Electrodeposition from Ionic Liquids
Edited by
Frank Endres, Douglas MacFarlane,
and Andrew Abbott
Related Titles
Wasserscheid, P., Welton, T. (eds.)
Ionic Liquids in Synthesis
2007
ISBN 978-3-527-31239-9
Staikov, G. T. (ed.)
Electrocrystallization in Nanotechnology
2007
ISBN 978-3-527-31515-4
Paunovic, M., Schlesinger, M.
Fundamentals of Electrochemical Deposition
2006
ISBN 978-0-471-71221-3
Ohno, H.
Electrochemical Aspects of Ionic Liquids
2005
ISBN 978-0-471-64851-2
Electrodeposition from Ionic Liquids
Edited by
Frank Endres, Douglas MacFarlane,
and Andrew Abbott
WILEY-VCH Verlag GmbH & Co. KGaA
The Editors
Prof. Dr. Frank Endres
Faculty of Natural and Material Sciences
Clausthal University of Technology
38678 Clausthal-Zellerfeld
Germany
Prof. Douglas MacFarlane
School of Chemistry
Monash University
Clayton, Victoria 3800
Australia
Prof. Dr. Andrew Abbott
Chemistry Department
University of Leicester
Leicester LE1 7RH
United Kingdom
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carefully produced. Nevertheless, authors,
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information contained in these books,
including this book, to be free of errors.
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details or other items may inadvertently be
inaccurate.
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c 2008 WILEY-VCH Verlag GmbH & Co.
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ISBN 978-3-527-31565-9
v
Contents
Preface
Foreword
IX
XIII
List of Contributors
List of Abbreviations
1
1.1
1.2
1.3
1.4
1.5
2
2.1
2.2
2.3
3
3.1
3.2
3.3
3.4
3.5
3.6
3.7
XV
XIX
Why use Ionic Liquids for Electrodeposition? 1
Andrew P. Abbott, Ian Dalrymple, Frank Endres, and Douglas R. MacFarlane
Non-aqueous Solutions 3
Ionic Fluids 3
What is an Ionic Liquid? 4
Technological Potential of Ionic Liquids 6
Concluding Remarks 12
References 12
Synthesis of Ionic Liquids 15
Tom Beyersdorff, Thomas J. S. Schubert, Urs Welz-Biermann, Will Pitner,
Andrew P. Abbott, Katy J. McKenzie, and Karl S. Ryder
Synthesis of Chloroaluminate Ionic Liquids 15
Air- and Water-stable Ionic Liquids 21
Eutectic-based Ionic Liquids 31
References 42
Physical Properties of Ionic Liquids for Electrochemical Applications 47
Hiroyuki Ohno
Introduction 47
Thermal Properties 47
Viscosity 54
Density 55
Refractive Index 56
Polarity 58
Solubility of Metal Salts 64
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
vi
Contents
3.8
3.9
Electrochemical Properties 66
Conclusion and Future Prospects
Acknowledgement 77
References 78
4
Electrodeposition of Metals 83
Thomas Schubert, Sherif Zein, El Abedin, Andrew P. Abbott,
Katy J. McKenzie, Karl S. Ryder, and Frank Endres
Electrodeposition in AlCl3 -based Ionic Liquids 84
Electrodeposition of Metals in Air- and Water-stable Ionic Liquids
Deposition of Metals from Non-chloroaluminate Eutectic
Mixtures 103
Troublesome Aspects 114
References 120
4.1
4.2
4.3
4.4
5
5.1
5.2
5.3
5.4
5.5
5.6
5.7
6
6.1
6.2
6.3
6.4
6.5
6.6
6.7
6.8
6.9
6.10
7
7.1
7.2
77
Electrodeposition of Alloys 125
I.-Wen Sun, and Po-Yu Chen
Introduction 125
Electrodeposition of Al-containing Alloys from Chloroaluminate
Ionic Liquids 126
Electrodeposition of Zn-containing Alloys from Chlorozincate
Ionic Liquids 132
Fabrication of a Porous Metal Surface by Electrochemical Alloying
and De-alloying 137
Nb–Sn 139
Air- and Water-stable Ionic Liquids 140
Summary 145
References 145
Electrodeposition of Semiconductors in Ionic Liquids 147
Natalia Borisenko, Sherif Zein El Abedin, and Frank Endres
Introduction 147
Gallium Arsenide 149
Indium Antimonide 149
Aluminum Antimonide 150
Zinc Telluride 150
Cadmium Telluride 151
Germanium 151
Silicon 155
Grey Selenium 160
Conclusions 164
References 164
Conducting Polymers 167
Jennifer M. Pringle, Maria Forsyth, and Douglas R. MacFarlane
Introduction 167
Electropolymerization – General Experimental Techniques 171
92
Contents
7.3
7.4
7.5
7.6
Synthesis of Conducting Polymers
Characterization 191
Future Directions 203
Conclusions 207
References 208
8
Nanostructured Metals and Alloys Deposited from Ionic Liquids 213
Rolf Hempelmann, and Harald Natter
Introduction 213
Pulsed Electrodeposition from Aqueous Electrolytes 215
Special Features of Ionic Liquids as Electrolytes 220
Nanocrystalline Metals and Alloys from Chlorometallate-based
Ionic Liquids 222
Nanocrystalline Metals from Air- and Water-stable Ionic Liquids 227
Conclusion and Outlook 234
Acknowledgement 235
References 235
8.1
8.2
8.3
8.4
8.5
8.6
9
9.1
9.2
9.3
9.4
9.5
9.6
10
10.1
10.2
10.3
10.4
10.5
10.6
11
11.1
11.2
11.3
177
Electrodeposition on the Nanometer Scale: In Situ Scanning
Tunneling Microscopy 239
Frank Endres, and Sherif Zein El Abedin
Introduction 239
In situ STM in [Py1,4 ] TFSA 241
Electrodeposition of Aluminum 245
Electrodeposition of Tantalum 250
Electrodeposition of Poly(p-phenylene) 252
Summary 256
References 256
Plasma Electrochemistry with Ionic Liquids 259
Jürgen Janek, Marcus Rohnke, Manuel Pölleth, and Sebastian A. Meiss
Introduction 259
Concepts and Principles 260
Early Studies 265
The Stability of Ionic Liquids in Plasma Experiments 269
Plasma Electrochemical Metal Deposition in Ionic Liquids 274
Conclusions and Outlook 282
Acknowledgement 283
References 283
Technical Aspects 287
Debbie S. Silvester, Emma I. Rogers, Richard G. Compton, Katy J. McKenzie,
Karl S. Ryder, Frank Endres, Douglas MacFarlane, and Andrew P. Abbott
Metal Dissolution Processes/Counter Electrode Reactions 287
Reference Electrodes for Use in Room-temperature Ionic Liquids 296
Process Scale Up 310
vii
viii
Contents
11.4
11.5
Towards Regeneration and Reuse of Ionic Liquids in Electroplating
Impurities 334
Appendix: Protocol for the Deposition of Zinc from a Type III
Ionic Liquid 344
References 345
12
Plating Protocols 353
Frank Endres, Sherif Zein El Abedin, Q. Liu, Douglas R. MacFarlane,
Karl S. Ryder, and Andrew P. Abbott
Electrodeposition of Al from 1-Ethyl-3-methylimidazolium
chloride/AlCl3 353
Electrodeposition of Al from 1-Butyl-3-methylimidazolium
chloride–AlCl3 – Toluene 356
Electrodeposition of Al from 1-Ethyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)amide/AlCl3 358
Electrodeposition of Al from 1-Butyl-1-methylpyrrolidinium
bis(trifluoromethylsulfonyl)amide/AlCl3 360
Electrodeposition of Li from 1-Butyl-1-methylpyrrolidinium
bis(trifluoromethylsulfonyl)amide/Lithium
bis(trifluoromethylsulfonyl)amide 362
Electrodeposition of Ta from 1-Butyl-1-methylpyrrolidinium
bis(trifluoromethylsulfonyl)amide 364
Electrodeposition of Zinc Coatings from a Choline Chloride:
Ethylene Glycol-based Deep Eutectic Solvent 365
References 367
12.1
12.2
12.3
12.4
12.5
12.6
12.7
13
13.1
13.2
13.3
13.4
13.5
13.6
13.7
13.8
13.9
13.10
13.11
Future Directions and Challenges 369
Frank Endres, Andrew P. Abbott, and Douglas MacFarlane
Impurities 369
Counter Electrodes/Compartments 370
Ionic Liquids for Reactive (Nano-)materials 371
Nanomaterials/Nanoparticles 372
Cation/Anion Effects 373
Polymers for Batteries and Solar Cells 373
Variable Temperature Studies 374
Intrinsic Process Safety 374
Economics (Price, Recycling) 375
Which Liquid to Start With? 375
Fundamental Knowledge Gaps 376
Subject Index
379
319
ix
Preface
Around ten years ago there were only about twenty papers per year dealing with
“ionic liquids” or “room-temperature molten salts”. Hence the term “ionic liquid”
was unknown to most of the scientific community at that time. Furthermore,
there was practically no knowledge of it in industry, and just a handful of groups
worldwide were investigating ionic liquids. Ionic liquids were perceived as an
academic curiosity. When one of us (F.E.) started his independent research in
1996 with the subject “room-temperature molten salts” many people cautioned
him about the eccentric topic. What was the reason for these opinions? From the
1950s to about 1995 most of the people in the community performed studies with
ionic liquids based on AlCl3 , often called “first generation” ionic liquids. These
are hygroscopic liquids, liberating HCl and a variety of oxo-chloroaluminates upon
exposure to moisture. Reproducible operation in these liquids requires either a
strictly controlled inert gas atmosphere with extremely low water concentration or
at least closed vessel conditions with limited contamination. Thus, these liquids
were considered to be difficult to work with and of little practical importance. On the
other hand as “room-temperature molten salts” they had attractive electrochemical
windows and allowed the electrodeposition of noble metals and of aluminum and
its alloys in micrometer thick layers. Aluminum is quite an interesting metal as
it is self-passivating, thus under air it forms spontaneously an oxide layer which
protects the metal underneath from further corrosion.
It was John Wilkes who realized that “room-temperature molten salts” would
only experience a widespread interest and uptake if they were stable under environmental conditions. Wilkes’ group published details of the first such liquid in 1992
using the BF4– and the PF–
6 anions, the latter showing a miscibility gap with water.
Thus these liquids could, in principle, be made water free. (Today we know that
–
ionic liquids containing BF–
4 and PF6 are subject to decomposition in the presence
of water.) Electrochemical studies showed that even these “early” ionic liquids had
wide electrochemical windows of about 4 V with cathodic limits of –2 to –2.5 V. vs.
NHE. This cathodic limit should, from the thermodynamic point of view, be wide
enough to electrodeposit many reactive elements.
Around 1995, Seddon realized that the expression “room-temperature molten
salts” was counter-productive. The expression “molten salt” was always associated
with “high temperature”, as also the editors (and many authors) of this book had
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
x
Preface
to experience. The introduction of the term “ionic liquids” for low melting molten
salts with melting temperatures below 100◦ C (the definition was actually coined by
Walden in 1914) created a clear distinction and the new term “ionic liquid” began
to appear in an unprecedented number of publications. In the late 1990s the first
papers on the electrodeposition of silver and copper in ionic liquids based on BF4–
and PF–
6 appeared in the literature. From these early papers it was immediately
clear that electrodeposition from ionic liquids is not trivial and is actually more
complicated than using ionic liquids based on AlCl3 . In addition to viscosity and
conductivity concerns, impurities such as water, halides and organic compounds
proved to be major difficulties.
Thus, nearly twenty years of accumulated knowledge on AlCl3 -based ionic liquids
could only be transferred with difficulty to this new class of ionic liquids, because
their Lewis acidity/basicity was totally different. Thus, the electrochemistry of these
second generation ionic liquids had to be re-invented, more or less. Nevertheless,
progress was not slowed and in 2002 alone there were already 600 papers dealing
with ionic liquids, about 10% concentrating on electrochemical aspects. In the
following years more stable ionic liquids with wider electrochemical windows were
developed and cathodic decomposition potentials as low as –3 V vs. NHE were
reported, opening the door to the electrodeposition of many reactive elements such
as Si, Ge, Ta, Al.
Recently a novel class of deep eutectic solvents based on choline chloride have
been developed. These can be handled easily under environmental conditions and
circumvent many problems that occur in aqueous solutions. They also offer the
first economically viable liquids that can be used on an industrial scale. As the
interest of electrochemists and classical electroplaters in ionic liquids has risen
strongly in the last few years we decided, in 2006, to collect the key aspects of the
electrodeposition from ionic liquids in the present book. The book has been written
by a panel of expert authors during late 2006 and the first half of 2007 and thus
describes the state of the art as of that point in time.
In Chapter 1 we explain the motivation and basic concepts of electrodeposition
from ionic liquids. In Chapter 2 an introduction to the principles of ionic liquids
synthesis is provided as background for those who may be using these materials for
the first time. While most of the ionic liquids discussed in this book are available
from commercial sources it is important that the reader is aware of the synthetic
methods so that impurity issues are clearly understood. Nonetheless, since a comprehensive summary is beyond the scope of this book the reader is referred for
more details to the second edition of Ionic Liquids in Synthesis, edited by Peter
Wasserscheid and Tom Welton. Chapter 3 summarizes the physical properties of
ionic liquids, and in Chapter 4 selected electrodeposition results are presented.
Chapter 4 also highlights some of the troublesome aspects of ionic liquid use. One
might expect that with a decomposition potential down to –3 V vs. NHE all available
elements could be deposited; unfortunately, the situation is not as simple as that
and the deposition of tantalum is discussed as an example of the issues. In Chapters
5 to 7 the electrodeposition of alloys is reviewed, together with the deposition of
semiconductors and conducting polymers. The deposition of conducting polymers
Preface
is still a little neglected in the literature, although the wide anodic decomposition
limit allows even benzene to be easily polymerized to poly( p-phenylene) in ionic
liquids.
Chapter 8 summarizes the principles of nanometal deposition as well as the few
examples of nanometal deposition in ionic liquids. Chapter 9 shows how scanning
probe microscopy can be used to study the electrodeposition of metals on the
submicro- and nano-scale. In situ STM is also used to probe impurities in the
ultralow concentration regime. Chapter 10 is devoted to a novel field in the scene,
i.e. plasma electrochemistry. By applying a glow-discharge plasma to the surface of
an ionic liquid which contains metal ions, suspensions of nanoparticles can be
made that might be of interest, for example, as catalysts. Chapter 11 is devoted to
technical aspects such as counter electrode reactions, reference electrodes (a very
complicated subject), upscaling, recycling and impurities. As industry increases the
scale of production the focus on cost and purity will be of increasing importance.
In Chapter 12 we provide some plating protocols, which will enable the reader to
begin electrodeposition experiments in ionic liquids. In Chapter 13 we have tried
to summarize the future directions of the field as we see them and challenging
aspects which, in our opinion, warrant further study. Of course, as the field is in a
permanent state of development, such a chapter can hardly be comprehensive, but
we hope that our thoughts, which are based on many years of experience, will help
to stimulate further the field of “electrodeposition from ionic liquids”.
Frank Endres, Andrew Abbott and Douglas R. MacFarlane
Yokohama, Japan, December 2007
xi
xiii
Foreword
It is always an honour to be asked to write a foreword for what is clearly an important book, but it is also a curse! What can you say that is original and interesting? –
Particularly when the editors themselves have written a Preface!! But this IS an
important book – electrodeposition is at the roots and heart of ionic liquid technology. It was one of the earliest applications of ionic liquids, and currently is one
of the exciting areas which are developing at an amazing rate. It is a wonderful
example of industrial processes developing hand-in-hand with academic research.
So, I accepted this cursed honour, and am very glad that I did: the opportunity to
see the chapters of this book in advance has been a privilege.
So let’s start with the obvious. This book on electrodeposition from ionic liquids
comes on the tail of another excellent Wiley book, edited in 2005 by Hiroyuki
Ohno, entitled “Electrochemical Aspects of Ionic Liquids”, an updated revision of
a 2003 Japanese volume with the title “Ionic Liquids: The Front and Future of
Material Development” (CMC Press, Tokyo). Is there any overlap? Well, in the
thirty-two chapters of this earlier edited book, which covers the whole spectrum of
electrochemistry in ionic liquids, there were only twenty pages devoted to the topic
of electrodeposition (an article by Yasushi Katayama). So, there is no significant
overlap to worry about.
Then, there is the whole question of the philosophy of the edited book? Has it
holistic value, or is it just a random collection of articles by disparate authors? Well,
the editors here have taken the same approach as Wasserscheid and Welton (“Ionic
Liquids in Synthesis”, 2nd Edit., Wiley-VCH, 2007). There is a well developed plan
for the book, and the chapters are integrated, and dovetail well. In addition, the
authors have been carefully selected – this is a book written by the leading lights of
the field. The editors have done an excellent job of producing a volume which deals
with the literature, conceptual framework, and practical aspects of the subject. It
was particularly pleasing to see chapters and sections dealing with the problems
associated with the area, including impurities, recycling and scale-up, reference
electrodes, and counter electrodes. Further, as one might expect with Andy Abbott
as one of the editors, there is a clear distinction drawn between ionic liquids and
deep eutectic solvents.
So, is this book perfect? Well, no! One thing drove me to distraction, and
it is a problem redolent of the wider literature – the choice of abbreviations
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
xiv
Foreword
for the ionic liquids – or, more precisely, the lack of choice! Different chapters used different systems, and each cation and anion was represented in at
least four ways within the book (meaning up to, or more than, sixteen possible abbreviations for some ionic liquids. For the simple, symmetrical and
common ionic liquid cation, 1,3-dimethylimidazolium, there were five different
abbreviations used: [MMIM], [mmim], [C1 mim], [C1 MIM], and [DMIM]; for 1butyl-1-methylpyrrolidinium, there were six different abbreviations used: [Py1,4 ],
P1,4 , [BMP], BuMePy, [c4 mpyr], and [c4 mpyrr]. And, even more bizarrely, for
the common anion bis(trifluoromethylsulfonyl)amide, six different abbreviations
were used: (CF3 SO2 )2 N, NTF, Tf2 N, NTf2 , TFSI, and TFSA. Thus, in principle
(I didn’t count!), 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide
could have had thirty-six possible abbreviations!! This is way past ridiculous. But
the problem doesn’t just lie with the editors; this is reflection of the problem in
the wider literature. The burgeoning of the ionic liquid literature during the past
decade has meant that there has been no period of stability during which a consensus could be reached. The question of a uniform system, and the wider question of
the fundamental definition of an ionic liquid, will have to be addressed elsewhere –
the problem is manifest here, however. Another minor issue is that some of the
English has a distinctly Germanic ring to it – but never to the point of obscuring
the meaning.
To summarise then, this book is timely and edited by three of the four main
experts in the field. It is planned with meticulous detail, and – of paramount
importance – it is authoritative. It is inconceivable that any researcher in the future
will not need access to this book, and it will be extensively cited. I congratulate
Frank, Doug and Andy on a wonderful volume. Editing books of this type is a
service to the community (no one does it for the royalties!), and we owe them a debt
of gratitude for the huge investment of time they have made.
Kenneth R. Seddon
The QUILL Research Centre
The Queen’s University of Belfast
Belfast, B9 5AG, U.K.
xv
List of Contributors
Andrew P. Abbott
Chemistry Department
University of Leicester
Leicester LE1 7RH
UK
Tom Beyersdorff
IOLITEG GmbH & Co. KG
Ferdinand-Porsche-Straße 5/1
79211 Denzlinger
Germany
Natalia Borisenko
Faculty of Natural and Materials
Sciences
Clausthal University of Technology
38678 Clausthal-Zellerfeld
Germany
Po-Yu Chen
Faculty of Medicinal and Applied
Chemistry
Kaohsiung Medical University
807 Kaohsiung
Taiwan
Richard G. Compton
Oxford University
Physical and Theoretical Chemistry
Laboratory
South Parks Road
Oxford OX1 3QZ
United Kingdom
Jan Dalrymple
C-Tech Innovation Ltd.
C-Tech
United Kingdom
Frank Endres
Faculty of Natural and Materials
Sciences
Clausthal University of Technology
38678 Clausthal-Zellerfeld
Germany
Email: frank.endres@tu-clausthal.de
Maria Forsyth
Australian Centre of Excellence for
Electromaterials Science
Department of Materials Engineering
Monash University
Wellington Road
Clayton
VIC 3800
Australia
Rolf Hempelmann
Physical Chemistry Department
Saarland University
66123 Saarbrücken
Germany
Jürgen Janek
Physikalisch-Chemisches Institut
Justus-Liebig-Universitaet Giessen
Heinrich-Buff-Ring 58
35392 Giessen
Germany
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
xvi
List of Contributors
Qunxian Liu
Faculty of Natural and Material
Sciences
Clausthal University of Technology
38678 Clausthal-Zellerfeld
Germany
Will Pitner
Merck KGaA
PLS R&D LSS Ionic Liquids 1
Frankfurter Str. 250
64271 Darmstadt
Germany
Douglas MacFarlane
School of Chemistry
Monash University
Wellington Road
Clayton
VIC 3800
Australia
Manuel Pölleth
Physikalisch-Chemisches Institut
Justus-Liebig-Universitaet Giessen
Heinrich-Buff-Ring 58
35392 Giessen
Germany
Katy J. McKenzie
Chemistry Department
University of Leicester
Leicester LE1 7RH
UK
Sebastian A. Meiss
Physikalisch-Chemisches Institut
Justus-Liebig-Universitaet Giessen
Heinrich-Buff-Ring 58
35392 Giessen
Germany
Harald Natter
Physical Chemistry Department
Saarland University
66123 Saarbrücken
Germany
Email: h.natter@mx.uni-saarland.de
Hiroyuki Ohno
Department of Biotechnology
Tokyo University of Agriculture and
Technology
2-24-16 Nakacho, Koganei
Tokyo 184-8588
Japan
Jennifer M. Pringle
School of Chemistry
Monash University
Wellington Road
Clayton
VIC 3800
Australia
Email: Jenny.Pringle@sci.monash.edu.
au
Emma I. Rogers
Oxford University
Physical and Theoretical Chemistry
Laboratory
South Parks Road
Oxford OX1 3QZ
United Kingdom
Marcus Rohnke
Physikalisch-Chemisches Institut
Justus-Liebig-Universitaet Giessen
Heinrich-Buff-Ring 58
35392 Giessen
Germany
Karl S. Ryder
Chemistry Department
University of Leicester
Leicester LE1 7RH
UK
List of Contributors
Thomas J. S. Schubert
Managing Director
IOLITEC GmbH & Co. KG
Ferdinand-Porsche-Straβe 5/1
79211 Denzlingen
Germany
Debbie S. Silvester
Oxford University
Physical and Theoretical Chemistry
Laboratory
South Parks Road
Oxford OX1 3QZ
United Kingdom
I.-Wen Sun
Department of Chemistry
National Cheng Kung University
Tainan 70101
Taiwan
Jorg Thöming
UFT
Section of Chemical Engineering –
Regeneration and Recycling
University of Bremen,
Leobener Str.
28359 Bremen
Germany
Daniel Waterkamp
UFT
Section of Chemical Engineering –
Regeneration and Recycling
University of Bremen,
Leobener Str.
28359 Bremen
Germany
Urs Welz-Biermann
New Business- Chemicals/Ionic
Liquids (NB-C)
Merck KGaA
NB-C, D1/311
Frankfurter Str. 250
64293 Darmstadt
Germany
Sherif Zein El Abedin
Faculty of Natural and Materials
Sciences
Clausthal University of Technology
38678 Clausthal-Zellerfeld
Germany
Permanent address:
Electrochemistry and Corrosion
Laboratory
National Research Centre
Dokki
Cairo
Egypt
xvii
xix
List of Abbreviations
Cations:
Pyrrolidinium cations:
1-Butyl-1-methylpyrrolidinium:
1-Propyl-1-methylpyrrolidinium:
[Py1,4 ], P1,4 , [BMP], BuMePy, [c4 mpyr],
[c4 mpyrr]
P1,3
Imidazolium Cations
1-Methyl-3-methylimidazolium:
1-Ethyl-3-methylimidazolium:
1-Propyl-3-methylimidazolium:
1-Butyl-3-methylimidazolium:
1-Butyl-3-butylimidazolium:
1-Butyl-3H-imidazolium:
1-Ethyl-3H-imidazolium:
1-Hexyl-3-methylimidazolium:
1-Octyl-3-methylimidazolium:
1-Propyl-2,3-dimethylimidazolium:
1-Butyl-2,3-dimethylimidazolium:
1-Etyl-2,3-dimethylimidazolium:
1-Hexyl-2,3-dimethylimidazolium:
1-Decyl-3-methylimidazolium:
1-Benzyl-3-methylimidazolium:
1-Hydroxyethyl-3-methylimidazolium:
1,2-Di-ethyl-3,4-dimethylimidazolium:
1-Alkyl-3-methylimidazolium:
1-(2-hydroxyethyl)-3methylimidazolium:
1-(2-methoxyethyl)-3methylimidazolium:
1-[2-(2-methoxyethoxy)ethyl]-3methylimidazolium:
[MMIM], [mmim], [C1 mim], [C1 MIM],
[DMIM]
[EMIM], [emim], [C2 mim], [C2 MIM]
[PMIM], [pmim], [C3 mim], [C3 MIM]
[BMIM], [bmim], [C4 mim]
[BBIM], [bbim]
[Hbim]
[Heim]
[HMIM], [hmim], [C6 mim], [HMPL]
[OMIM], [omim], [C8 mim]
[p-DiMIM], [DMPIM]
[b-DiMIM], [C4 -DMIM]
[e-DiMIM]
[C6 -DMIM]
[decyl-MIM], [C10 MIM], [C10 mim]
[BZMIM]
[HO(CH2 )2 MIM], [C2 OHMIM]
[DEDMIM]
[Cn MIM], [Cn mim]
[C2 OHmim]
[C3 Omim]
[C5 O2 mim]
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
xx
List of Abbreviations
Pyridinium Cations:
N-Methylpyridinium
N-Ethylpyridinium
N-Propylpyridinium
N-Butylpyridinium:
N-Hexylpyridinium:
[MP]
[EP], [C2 py], [EtPy]
[PP]
[BP], [bpyr], [bpyrr], [C4 py]
[HP], [HPYR], [C16 py]
Piperidinium Cations:
N-Ethyl-N-methylpiperidinium:
N-Propyl-N-methylpiperidinium:
N-Butyl-N-methylpiperidinium:
[C2 mPip]
[C3 mPip], [PP13 ]
[C4 mPip], [PP14 ]
Phosphonium Cations:
Tri-hexyl-tetradecylphosphonium:
[Ph3 t], [P14,6,6,6 ], [P6,6,6,14 ]
Pyrazolium Cations:
N,N-Diethyl-3-methylpyrazolium
[DEMPZ]
Ammonium-Cations:
Trimethylammonium:
Tetramethylammonium:
1,1,1-Trimethyl-1methoxyethylammonium:
Butyl-trimethylammonium:
Benzyl-trimethylammonium:
Propyl-trimethylammonium:
1-Cyanomethyl-1,1,1trimethylammonium:
1,1-Dimethyl-1-ethyl-1methoxyethylammonium:
1,1-Diethyl-1-methyl-1methoxyethylammonium:
Tributyl-methylammonium:
Trimethyl-n-hexylammonium:
Tetraethylammonium:
Triethyl-hexylammonium:
Tetrabutylammonium:
Triethyl-hexylammonium:
Hydroxyethyl-trimethylammonium:
Butyl-diethyl-methylammonium:
[TMHA]
[N1111 ], [TMA]
[N111,2O1 ]
[N1114 ], [N4111 ], [BTMA]
[BTMA]
[N1113 ], [N3111 ], [PTMA]
[N111,1-CN ]
[N112,2O1 ]
[N122,2O1 ]
[N4441 ], [TBMA]
[N1116 ], [TMHA]
[N2222 ], [TEA]
[N2226 ]
[N4444 ], [TBA], Bu4 N
[N6222 ]
[Me3 NC2 H4 OH], Ch also called choline
[N1224 ]
List of Abbreviations
Sulfonium Cations:
Trimethylsulfonium:
Triethylsulfonium:
Tributylsulfonium:
[S111 ]
TES, [S222 ]
TBS, [S444 ]
Anions:
Bis(trifluoromethylsulfonyl)
amide:
Trispentafluoroethyltrifluorophosphate:
Trifluoroacetate:
Trifluoromethylsulfonate:
Dicyanoamide:
Tricyanomethide:
Tetracyanoborate:
Tetraphenylborate:
Tris(trifluoromethylsulfonyl)
methide:
Thiocyanate:
(CF3 SO2 )2 N, NTF, Tf2 N, NTf2 , TFSI, TFSA
Sometimes this anion is also called
bis(trifluoromethylsulfonyl)imide or bistriflamide,
bistriflimide
FAP
ATF, TFA
OTF, OTf, TFO, Tf Also called
trifluoromethanesulfonate
DCA
TCM
TCB
[BPh4 ]
[CTf3 ]
SCN
Other chemicals:
[CHES]:
ChCl:
DCM:
EDOT:
EG:
Fc:
Fc+ :
GC:
ITO:
PC:
PEDOT:
TMPD:
TMS:
2-(Cyclohexylamino)ethylsulfonate
Choline chloride
Dichloromethane
Ethylenedioxythiophene
Ethyleneglycole
Ferrocene
Ferrocinium
Glassy carbon
Indium-tin-oxide
Propylenecarbonate
Polyethylenedioxythiophene
Tetramethylphenylenediamine
Tetramethylsilane
Abbreviations:
AAS:
ACD:
AFM:
ATR-FTIR:
Atomic Absorption Spectroscopy
Anomaleous Codeposition
Atomic Force Microscopy
Attenuated Total Reflection Fourier Transform
Infrared Spectroscopy
xxi
xxii
List of Abbreviations
BASIL:
CIS:
CV:
CVD:
DSSC:
ECALE:
EC-STM:
EDX, EDS, EDAX:
EIS:
EMF:
EQCM:
FAB MS:
FFG-NMR:
FWHM:
HBD:
HOPG:
HO-ESY:
H-REM, H-SEM:
H-TEM:
ICP:
LCA:
LED:
LSV:
MBE:
MNDO:
NHE:
NMR:
OCP:
OPD:
PECD:
PED:
PLED:
PPP:
PVD:
RTIL:
SAED:
SIGAL:
STM:
TSIL:
UPD:
UHV:
VFT, VTF:
XPS:
XRD:
Biphasic Acid Scavenging Utilizing Ionic Liquids
Copper-indium-selenide
Cyclic Voltammetry
Chemical Vapor Deposition
Dye Sensitized Solar Cell
Electrochemical Atomic Layer Epitaxy
Electrochemical in situ scanning tunnelling microscopy
Energy Dispersive X-ray analysis
Electrochemical Impedance Spectroscopy
Electromotive Force
Electrochemical Quarz Crystal Microbalance
Fast atom bombardment mass spectroscopy
Fixed Field Gradient Nuclear Magnetic Resonance
Spectroscopy
Full width at half maximum
Hydrogen Bond Donor
Highly Oriented Pyrolytic Graphite
Heteronuclear Overhauser Effect Spectroscopy
High Resolution Scanning Electron Microscopy
High Resolution Transmission Electron Microscopy
Inductively Coupled Plasma (Spectroscopy)
Life Cycle Analysis
Light Emitting Diode
Linear Sweep Voltammetry
Molecular Beam Epitaxy
Modified neglect of diatomic overlap
Normal Hydrogen Electrode
Nuclear Magnetic Resonance
Open Circuit Potential
Overpotential deposition
Plasmaelectrochemical deposition
Pulsed Electrodeposition
Polymer Light Emitting Diode
Poly-para-phenylene
Physical Vapour Deposition
Room Temperature Ionic Liquid
Selected Area Electron Diffraction
Siemens Galvano-Aluminium
Scanning Tunnelling Microscopy
Task Specific Ionic Liquid
Underpotential deposition
Ultrahigh Vacuum
Vogel-Tammann-Fulcher
X-ray photoelectron spectroscopy
X-ray diffraction
1
1
Why use Ionic Liquids for Electrodeposition?
Andrew P. Abbott, Ian Dalrymple, Frank Endres, and Douglas R. MacFarlane
With any great voyage of discovery the explorer should always be asked at the
outset “Why are you doing this?” To answer the question “Why use ionic liquids
for electrodeposition?” it is first necessary to look at current best practice and find
its limitations.
It is widely recognised that in 1805 Italian chemist, Luigi Brugnatelli made
the first experiments in what we now know as electroplating. Brugnatelli used
the newly discovered Voltaic Pile to deposit gold “I have lately gilt in a complete
manner two large silver medals, by bringing them into communication by means
of a steel wire, with a negative pole of a voltaic pile, and keeping them one after the
other immersed in ammoniuret of gold newly made and well saturated” [1]. The
process was later improved by John Wright who found that potassium cyanide was
a beneficial electrolyte to add for silver and gold plating as it allowed thick adherent
deposits to be obtained. Until the middle of the 19th century the production of
jewellery and the gilding of decorative items were the main uses of electrodeposition.
With an increased understanding of electrochemistry, the practice of metal
deposition spread to non-decorative metals such as nickel, brass, tin, and zinc
by the 1850s. Even though electroplated goods entered many aspects of manufacturing industry very little changed about the physical processes involved in
electrodeposition for about 100 years. It was only with the advent of the electronics industry in the middle of the 20th century that significant changes occurred in the hardware and chemistry of the plating solutions. The post-war
period saw an increase in gold plating for electronic components and the use
of less hazardous plating solutions. This trend has continued with increased
control of hazardous materials to the environment. Improved solution composition and power supply technology has allowed the development of fast and
continuous plating of wire, metal strips, semiconductors and complex substrate
geometries.
Many of the technological developments seen in the electronics industry depend
upon sophisticated electroplating including the use of exotic metals and this is
one of the drivers for new technology within the electroplating sector. The other
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
2 1 Why use Ionic Liquids for Electrodeposition?
main driver is the search for alternative technologies for metals such as chromium,
nickel and cadmium. Anti-corrosion and wear-resistant coatings are predominant
markets in the electroplating sector and environmental directives will evidently
limit their usage in the future.
The main metals of that commercially deposited are Cr, Ni, Cu, Au, Ag, Zn
and Cd together with a number of copper and zinc-based alloys [1]. The whole
electroplating sector is based on aqueous solutions. There are some niche markets
based on organic solvents such as aluminum but these are very much exceptions.
Metals outside this list are generally deposited using plasma or chemical vapor
deposition techniques (PVD and CVD). These methods allow the coating of most
substrates (metal, plastic, glass, ceramic etc.) not only with metal but also with alloys
or compounds (oxide, nitride, carbide, etc.), without damaging the environment.
Although these techniques are technically interesting, it is regrettable that they
always involve high capital investment and it is difficult to prepare thick coatings,
thus they are only applied to high value niche markets.
Clearly the key advantages of using aqueous solutions are:
Ĺ
Ĺ
Ĺ
Ĺ
Ĺ
Ĺ
Cost
Non-flammable
High solubility of electrolytes
High conductivities resulting in low ohmic losses and good throwing power
High solubility of metal salts
High rates of mass transfer.
For these reasons water will remain the mainstay of the metal plating industry,
however, there are also limitations of aqueous solutions including:
Ĺ Limited potential windows
Ĺ Gas evolution processes can be technically difficult to handle and result in hydrogen embrittlement
Ĺ Passivation of metals can cause difficulties with both anodic and cathodic materials
Ĺ Necessity for complexing agents such as cyanide
Ĺ All water must eventually be returned to the water course.
These prevent aqueous solutions being applied to the deposition of several technically important materials.
The key technological goals include replacement of environmentally toxic metal
coatings, deposition of new alloys and semiconductors and new coating methods
for reactive metals. The main driving force for non-aqueous electrolytes has been
the desire to deposit refractory metals such as Ti, Al and W. These metals are
abundant and excellent for corrosion resistance. It is, however, the stability of their
oxides that makes these metals difficult to extract from minerals and apply as
surface coatings.
1.2 Ionic Fluids 3
1.1
Non-aqueous Solutions
There is clearly a range of alternative non-aqueous solutions that could be used.
Ideally, to obtain the properties required for an electrolyte solution, polar molecules
have to be used and these should preferably be small to obtain the requisite high
fluidity. Unfortunately, all polar molecules result from having electronegative elements which by their nature makes them good electron donors. Accordingly, they
will strongly coordinate to metal ions making them difficult to reduce. While a
number of metals have been deposited from polar organic solvents these tend to
be the rather noble metals and the processes offer few advantages over aqueous
solutions. Some studies have been made using non-polar organic solvents, predominantly aromatic hydrocarbons but these suffer from the serious disadvantage that
the dissolved electrolytes are highly associated and the solutions suffer from poor
conductivity. The solutions do, however, have wide potential windows and it has
been demonstrated that metals such as aluminum and titanium (Ti at least in very
thin layers) can be deposited from them. One of the most successful non-aqueous
processes is the SIGAL process developed in the late 1980s for the deposition of
aluminum from toluene [2, 3]. The aluminum source is triethyl aluminum which is
pyrophoric and, despite the high flammability of the electrolyte solution, the process
has been commercialized and is currently the only electrochemical method for the
deposition of aluminum. A review of electrochemistry in non-aqueous solutions is
given by Izutsu [4].
1.2
Ionic Fluids
Clearly an alternative to molecular solvents is the use of ionic fluids. Ionic materials usually melt at high temperatures due to their large lattice energies. Hightemperature molten salts have been extensively used for the electrowinning of
metals such as Li, Na, Ti and Al at temperatures of up to 1000 ◦ C [5–7]. They have
wide potential windows, high conductivities and high solubilities for metal salts,
in fact they have most of the advantages of aqueous solutions and overcome most
of the limitations of aqueous solutions, but clearly they suffer from the major limitation that the operational conditions are difficult to achieve and limit the range of
substrates that can be used for deposition.
The alternative to high-temperature molten salts is to use an ionic substance that
melts at a low temperature. While this may sound like an oxymoron it is logical
to suppose that the melting point of an ionic substance is related to ionic size and
if the ions are made large enough the material will eventually melt at ambient
conditions. A significant amount of work was carried out in the middle of the
20th century with the aim of developing lower temperature molten salts. One of
the key aims was a lower temperature melt for aluminum deposition which led
4 1 Why use Ionic Liquids for Electrodeposition?
to the formation of Li+ / K+ / AlCl3 eutectics which have freezing points close to
100 ◦ C [8]. These low freezing points arise due to the large chloroaluminate anions
(AlCl4 − and Al2 Cl7 − ) that form in the eutectic mixtures and have low lattice energies. The use of quaternary ammonium salts, particularly pyridinium and imidazolium salts, has pushed the freezing point down to ambient conditions. The term
“ionic liquids” was coined to distinguish these lower temperature ionic fluids from
the high temperature analogues which are composed predominantly of inorganic
ions.
The synthesis and properties of a range of ionic liquids are briefly summarized in
the following chapter while the history and chemical properties of these liquids are
covered in several well known reviews [9–12]. Several applications of ionic liquids
are being tested and these are as diverse as fuel desulfurization [13] and precious
metal processing [14] but few have yet come to practical fruition.
R
BASF’s BASIL process [15] and the Dimersol process [16] have both been commercialized. The former uses the ionic liquid as a phase transfer catalyst to produce
alkoxyphenylphosphines which are precursors for the synthesis of photoinitiators
used in printing inks and wood coatings. The imidazole acts as a proton scavenger
in the reaction of phenyl-chlorophosphines with alcohols to produce phosphines.
R
The Dimersol process uses a Lewis acid catalyst for the dimerization of butenes
to produce C8 olefins which are usually further hydroformylated giving C9 alcohols
used in the manufacture of plasticizers. Several other processes are also at the pilot
plant scale and some ionic liquids are used commercially as additive e.g. binders
in paints.
1.3
What is an Ionic Liquid?
The recognised definition of an ionic liquid is “an ionic material that is liquid
below 100 ◦ C” but leaves the significant question as to what constitutes an ionic
material. Some authors limit the definition to cations with discrete anions e.g. BF4 − ,
NO3 − . This definition excludes the original work on chloroaluminate systems and
the considerable work on other eutectic systems and is therefore unsatisfactory.
Systems with anionic species formed by complex equilibria are difficult to categorise
as the relative amounts of ionic species depend strongly on the composition of the
different components.
Ionic liquids have also been separated into first and second generation liquids
[10]; where first generation liquids are those based on eutectics and second generation have discrete anions [17]. Others have sought to further divide the first
generation liquids into separate types depending on the nature of the Lewis or
Brønsted acid that complexes [18]. While there is some dispute whether eutectics
with Brønsted acids constitute ionic liquids at all there are others who seek to widen
the description of ionic liquids to include materials such as salt hydrates [19].
In general, ionic liquids form because the charge on the ions is delocalized and
this gives rise to a reduction in lattice energy. The majority of ionic liquids are
1.3 What is an Ionic Liquid?
described by the equilibrium:
cation + anion + complexing agent ↔ cation + complex anion
(1.1)
Potentially, complex cations could also be formed using species such as cryptands
or crown ethers:
cation + anion + complexing agent ↔ complex cation + anion
(1.2)
The confusion arises from the magnitude of the equilibrium constant. For discrete anions such as BF4 − and even ((CF3 SO2 )2 N)− the equilibrium lies clearly to
the right of Eq. (1.1). For some eutectic-based liquids the equilibrium constant is
also to the right e.g.
Cat+ Cl− + AlCl3 ↔ Cat+ + AlCl−
4
(1.3)
But the addition of more Lewis acid produces other anionic species.
Cat+ Cl− + 2AlCl3 ↔ Cat+ + Al2 Cl−
7
(1.4)
The use of less Lewis acidic metals e.g. ZnCl2 or SnCl2 will lead to a small amount
of Cl− . The species formed between the anion and the complexing agent becomes
weaker when a Brønsted acid e.g. urea is used [18].
Cat+ Cl− + urea ↔ Cat+ + Cl− · urea
(1.5)
Others have claimed that, in the extreme, water can act as a good Brønsted acid
and, in the extreme, hydrate salts can act as ionic liquids [19].
LiClO4 + 3.5 H2 O ↔ Li+ · xH2 O + ClO−
4 · yH2 O
(1.6)
Ionic liquids with discrete anions have a fixed anion structure but in the eutecticbased liquids at some composition point the Lewis or Brønsted acid will be in
considerable excess and the system becomes a solution of salt in the acid. A similar
scenario also exists with the incorporation of diluents or impurities and hence we
need to define at what composition an ionic liquid is formed. Many ionic liquids
with discrete anions are hydrophilic and the absorption of water is found sometimes
to have a significant effect upon the viscosity and conductivity of the liquid [20–22].
Two recent approaches to overcome this difficulty have been to classify ionic liquids
in terms of their charge mobility characteristics [23] and the correlation between the
molar conductivity and fluidity of the liquids [24]. This latter approach is thought
by some to be due to the validity of the Walden rule
η = constant
(1.7)
5
6 1 Why use Ionic Liquids for Electrodeposition?
in ionic liquids, where is the molar conductivity and η is the viscosity. This is,
however a misrepresentation of Eq. (1.7) which was found empirically and is only
strictly valid for a specific ion at infinite dilution and constant temperature. The
Walden rule is a useful tool for approximate classification of ionic liquids but it
actually follows from the Nernst–Stokes–Einstein equation (See Chapters 2.3 and
11.3) [23]. Most importantly, deviations from the Walden rule do not necessarily
show that a salt is not an ionic liquid but more usually occur where ionic species
deviate from the model of centro-symmetric spherical ions with similar ionic radii.
The Walden rule can, however, be used to give evidence of different charge transfer
mechanisms e.g. a Grotthus mechanism for protonic ionic liquids [24].
In this book a broad-church of ionic liquids will be assumed, encompassing all
of the above types because, in the discipline of electrodeposition, it is the resultant
deposit that is important rather than the means. As will be seen later there is also
a very fine line between a concentrated electrolyte solution and an ionic liquid
containing diluents.
1.4
Technological Potential of Ionic Liquids
A series of transition- and main group-metal-containing ionic liquids have been
formulated and the feasibility of achieving electrodeposition has been demonstrated
for the majority of these metals, Figure 1.1 shows the elements in the periodic table
that have been deposited using ionic liquids. Details of these systems are given
in the subsequent chapters and concise summaries exists in recently published
reviews [18,25,26].
It must be stressed that while the deposition of a wide range of metals has been
demonstrated from a number of ionic liquids the practical aspects of controlling
deposit morphology have not been significantly addressed due to the complex
nature of the process parameters that still need to be understood. Despite the lack
of reliable models to describe mass transport and material growth in ionic liquids,
there are tantalizing advantages that ionic liquid solvents have over aqueous baths
that make the understanding of their properties vitally important. Some of these
advantages include:
Ĺ Electroplating of a range of metals impossible to deposit in water due to hydrolysis
e.g. Al, Ti, Ta, Nb, Mo, W. As an example, the deposition of Al by electrolysis
in a low-temperature process has long been a highly desirable goal, with many
potential applications in aerospace for anti-friction properties, as well as replacing
Cr in decorative coatings. The deposition of Ti, Ta, Nb, Mo, W will open important
opportunities in various industries, because of their specific properties (heat,
corrosion, abrasion resistance, low or high density etc.).
Ĺ Direct electroplating of metals on water-sensitive substrate materials such as Al,
Mg and light alloys with good adherence should be possible using ionic liquids.
Ĺ There is potential for quality coatings to be obtained with ionic liquids rather
than with water. Currently available metallic coatings suffer from hydrogen
1.4 Technological Potential of Ionic Liquids 7
Fig. 1.1 Summary of the elements deposited as single metals or alloys.
Ĺ
Ĺ
Ĺ
Ĺ
embrittlement; a major problem caused by gaseous hydrogen produced during
water electrolysis. During electroplating with ionic liquids, negligible hydrogen
is produced, and coatings will have the better mechanical properties.
Metal ion electrodeposition potentials are much closer together in ionic liquids
compared with water, enabling easier preparation of alloys and the possibility
of a much wider range of possible electroplated alloys, which are difficult or
impossible in water.
Ionic liquids complex metals and therefore offer the possibility to develop novel
electroless plating baths for coating polymers (e.g. in electronics) without the
need for the toxic and problematic organic complexants used in water.
Although the cost of ionic liquids will be greater than aqueous electrolytes,
high conductivity and better efficiency will provide significant energy savings
compared with water, and capital costs will be much lower than the alternative
techniques PVD and CVD.
When used in electropolishing and electropickling processes, strongly acidic
aqueous electrolytes create large quantities of metal-laden, corrosive effluent
solution, whereas in ionic liquid electrolytes the metals will precipitate and be
readily separated and recycled.
8 1 Why use Ionic Liquids for Electrodeposition?
Ĺ The replacement of many hazardous and toxic materials currently used in water,
e.g. toxic form of chromium(VI), cyanide, highly corrosive and caustic electrolytes, would save about 10% of the current treatment costs.
Ĺ Nanocomposite coatings – nanoparticles giving improved properties compared to
microparticles e.g. thermal and electrical conductivity, transparency, uniformity,
low friction.
Ĺ An increased range of metal coatings on polymers is accessible by electroless
plating using ionic liquids containing reducing agents
In the longer term, specialist, ionic liquids will enable technically complex highadded-value products to be introduced, e.g. semiconductor coatings, special magnetic alloys, nanoparticle composite coatings with special erosion/corrosion properties, metal foams for energy storage activated surfaces for self-sterilization purposes
(e.g. through photo-catalysis), etc.
Also, metals have significantly different reduction potentials in ionic liquid solutions compared to water. For example the difference in reduction potential between
Cr and Pt in ionic liquids may be as little as 100 mV whereas in aqueous solutions
it is in excess of 2 V. One consequence of this characteristic is that alloy coatings
may be prepared more readily and that it should be possible to develop many novel
alloy coatings.
A fundamental advantage of using ionic liquid electrolytes in electroplating is
that, since these are non-aqueous solutions, there is negligible hydrogen evolution during electroplating and the coatings possess the much superior mechanical
properties of the pure metal. Hence essentially crack-free, more corrosion-resistant
deposits are possible. This may allow thinner deposits to be used, thus reducing
overall material and power consumption still further.
The electrodeposition of metals from ionic liquids is a novel method for the production of nanocrystalline metals and alloys, because the grain size can be adjusted
by varying the electrochemical parameters such as over-potential, current density,
pulse parameters, bath composition and temperature and the liquids themselves.
Recently, for the first time, nanocrystalline electrodeposition of Al, Fe and Al–Mn
alloy has been demonstrated.
The properties of the new electrolyte media could also provide much higher
health and safety standards for employees in the workplace, i.e. elimination of
hazardous vapors, elimination of highly corrosive acidic/alkaline solutions and
substantially reduced use of toxic chemicals. These issues are dealt with later in the
book. Current aqueous processing systems have a strongly negative impact on the
environment (risk of groundwater contamination, soil pollution), which obliges
the treatment of wastewater and the dumping of the ultimate waste in landfill.
The metal finishing industry in general estimates that at least 15% of turnover is
related to the cost of treatment for environmental protection. Legislation within
the framework of sustainable development is increasingly stringent (e.g. European
Commission directive 96/61/EC concerning “integrated pollution prevention and
control”). Thus, industries using metal finishing processes must search for new
techniques to achieve these environmental goals. In addition to the growing costs
1.4 Technological Potential of Ionic Liquids 9
and negative effects on competitiveness, it is a question of survival in the years to
come. The technology developed by this project is generic to most metal-plating
systems and, as such, should represent a significant advancement for the environmental sustainability of the metal finishing and electronics manufacturing sectors.
Ionic liquids are based on large non-centrosymmetric organic cations with complex anions, which are liquid at room temperature. The range of new ionic liquids
has insignificant vapor pressure (thus odorless), some are non-toxic (and even
completely biodegradable) and most are highly conductive compared with organic
electrolyte solutions. This statement has to be tempered, naturally, because compared to the current state of the art, i.e. concentrated inorganic acids, the conductivities are at best 10 to 100 times lower. An advantage might be that ionic liquids
can be operated at temperatures above 100 ◦ C where ionic conductivities of up to
0.2 −1 cm−1 are achievable. The ongoing development of ionic liquids might lead
to even better conducting liquids.
There are, however numerous risk elements in the development of ionic liquids:
Ĺ Coatings must achieve quality standards and a large amount of process development is required.
Ĺ A life cycle analysis (LCA) and an environmental impact study have not been
completed for any of this technology.
Ĺ Issues of scale-up and integration design of generic prototype systems have not
been addressed systematically.
Ĺ Some applications are at a fundamental research stage with associated higher
risk, i.e. electroless coating, semiconductors, anodising, nanocomposite coatings.
Ĺ Process economics are expected to be favorable for high-added-value products,
but there are likely to be applications where economics are less favorable.
Ĺ For improved existing products, customer acceptance is likely to be a significant
factor, i.e. reluctance to change product specifications.
The potential impact is extremely broad and fundamental in nature, because the
research will explore a totally innovative approach to metal finishing technology,
which has never been exploited previously. The use of this completely different
type of solvent/electrolyte system, entirely changes the normal behavior of metal
finishing processes seen in traditional aqueous electrolytes and an extensive range
of entirely new processes and products can be expected.
The following chapters discuss the history, development and physical properties
of low-temperature ionic substances but in this section it is useful to discuss the
differences that arise in changing from a molecular to an ionic environment and
the implications that this will have for electrodeposition processes occurring at an
electrode surfaces.
There are several physical plating parameters that are different in an ionic liquid
from those in an aqueous solution.
Temperature: Ionic liquids have wide liquid regions, typically in the range −50
to 250 ◦ C which allow more thermodynamic control than is possible in aqueous
solutions. This may have potential benefits for the development of new alloys.
10 1 Why use Ionic Liquids for Electrodeposition?
Diluents: Ionic liquids can be diluted with a range of organic and aqueous solvents
and these significantly affect conductivity, viscosity and metal speciation. The effects
have not as yet been characterized and a significant amount of fundamental data
still needs to be obtained. A significant amount of work has been carried out on the
chloroaluminate-based ionic liquids, although their use in other ionic liquids has
been generally ignored.
Cation: Cationic structure and size will affect the viscosity and conductivity of the
liquid and hence will control mass transport of metal ions to the electrode surface.
They will also be adsorbed at the electrode surface at the deposition potential and
hence the structure of the double layer is dominated by cations. Some studies
have shown that changing the cationic component of the ionic liquid changes the
structure of deposits from microcrystalline to nanocrystalline [27]. While these
changes are undeniable more studies need to be carried out to confirm that it is
a double layer effect. If this is in fact the case then the potential exists to use the
cationic component in the liquid as a built-in brightener.
Double Layer Structure: Surprisingly few studies have been carried out into the
double layer structure of ionic liquids. This is partially due to experimental difficulties but also to interpretation of the resulting impedance spectra. What is clearly
evident, however, is that the double layer in an ionic liquid cannot be described
by applying the models used for aqueous solutions [28–30]. A study using imidazolium bistriflamide, (F3 CSO2 )2 N− and BF4 − salts suggested that a model of
alternating anion and cation layers may be applicable to the data [29, 30]. Baldelli
[31, 32] concluded that the double layer is one ion layer thick using sum frequency
generation spectroscopy and electrochemistry to probe the electric field at the ionic
liquid/electrode interface. The double layer capacitance in an ionic liquid is considerably smaller than in an aqueous solution and less than that predicted by having
a perfect Helmholtz layer at the interface, which could result from the presence
of ion pairs at the electrode surface at all potentials. Most likely the double layer
structure is also influenced by cation/anion interaction.
While the structure at the electrode/ionic liquid interface is uncertain it is clear
that in the absence of neutral molecules the concentration of anions and cations
at the interface will be potential dependent. The main difference between aqueous solutions and ionic liquids is the size of the ions. The ionic radii of most
metal ions are in the range 1–2 Å whereas for most ions of an ionic liquid they
are more typically 3–5 Å. This means that in an ionic liquid the electrode will
be coated with a layer of ions at least 6–7 Å thick. To dissolve in an ionic liquid most metal species are anionic and hence the concentration of metal ions
close to the electrode surface will be potential dependent. The more negative the
applied potential the smaller the concentration of anions. This means that reactive metals such as Al, Ta, Ti and W will be difficult to deposit as the effective
concentration of metal might be too low to nucleate. It is proposed, as one explanation, that this is the reason that aluminum cannot be electrodeposited from
Lewis basic chloroaluminate ionic liquids. More reactive metals such as lithium can
however be deposited using ionic liquids because they are cationic and therefore
1.4 Technological Potential of Ionic Liquids 11
present at high concentrations close to the electrode surface at large negative overpotentials. The strategy to electrodeposit reactive metals must therefore be to either make cationic metal complexes or to work with metal salts at high concentrations.
Anode material: In aqueous solutions the anodic processes are either breakdown
of the electrolyte solution (with oxygen evolution at an inert anode being favored)
or the use of soluble anodes. The use of soluble anodes is limited by the passivation
of many metals in aqueous solutions. In ionic liquids, however, the first option
is not viable due to the cost and the nature of the anodic breakdown products.
New strategies will therefore have to be developed to use soluble anodes where
possible or add a sacrificial species that is oxidized to give a benign gaseous product.
Preliminary data have shown that for some metals the anodic dissolution process
is rate limiting and this affects the current distribution around the cathode and the
current density that can be applied.
Electrolytes: The above issue of double layer structure is important to the mechanism of nucleation and growth in ionic liquids, it may therefore be possible to
control the structure at the electrode/solution interface by addition of an inert
electrolyte. In this respect most Group 1 metals are soluble in most ionic liquids,
although it is only generally lithium salts that exhibit high solubility. In ionic liquids
with discrete anions the presence of Group 1 metal ions can be detrimental to the
deposition of reactive metals such as Al and Ta where they have been shown to be
co-deposited despite their presence in trace concentrations.
Brighteners: Brighteners are added to most aqueous electroplating solutions and
work by either complexing the metal ions and decreasing the rate of nucleation or
by acting as an interfacial adsorbate, blocking nucleation and hindering growth.
Aqueous brighteners have not been studied in depth in ionic liquids and it is
doubtful that they will function in the same way as they do in water because of
the difference in double layer structure and mass transport. In unpublished work
we have surveyed the use of aqueous brightener compounds and applied them to
the deposition of zinc and chromium from Type 2 or Type 3 eutectics (see also
Chapter 2). None of these were found to be effective in ionic liquids.
A small amount of work has been carried out into brighteners that complex
the metal ions in solution (see Chapter 11.3) but again no systematic studies
have been carried out. Brighteners which rely on electrostatic or hydrophobic interactions may function in ionic liquids but their efficacy is likely to be surface
and cation/anion specific. To date all systems that have produced bright metallic finishes have been found to have a nanocrystalline structure which may be
due to a progressive nucleation mechanism. This is currently under investigation
and if confirmed it will help significantly with the design of future brightener
systems.
As with other solutes in ionic liquids, the general rule of like dissolving like
is applicable i.e. ionic species will generally be soluble as will species capable of
interacting with the anion. Aromatic species tend to exhibit poor solubility in ionic
liquids consisting of aliphatic cations and vice versa.
12 1 Why use Ionic Liquids for Electrodeposition?
1.5
Concluding Remarks
What is clear from this introduction is that the journey into the area of metal
deposition from ionic liquids has tantalizing benefits. It is also clear that we have
only just begun to scratch the surface of this topic. Our models for the physical
properties of these novel fluids are only in an early state of development and
considerably more work is required to understand issues such as mass transport,
speciation and double layer structure. Nucleation and growth mechanisms in ionic
liquids will be considerably more complex than in their aqueous counterparts but
the potential to adjust mass transport, composition and speciation independently
for numerous metal ions opens the opportunity to deposit new metals, alloys
and composite materials which have hitherto been outside the grasp of electroplaters.
References
1 (2000) Modern Electroplating, 4th edn (eds
M. Schlesinger and M. Paunovic), John
Wiley & Sons, Inc., New York.
2 Peled, E. and Gileadi, E. (1976) J.
Electrochem. Soc., 123, 15–19.
3 Simanavicius, L. (1990) Chemija, 3, 3–30.
4 Izutsu, K. (2002) Electrochemistry in
Non-aqueous Solutions, Wiley-VCH,
Verlag GmbH.
5 Kruesi, W.H. and Fray, D.J. (1993) Metall.
Trans. B., 24, 605.
6 Fray, D.J. and Chen, G.Z. (2004) Mater.
Sci. Technol., 20, 295.
7 Grjotheim, K., Krohn, C., Malinovsky, M.,
Matiasovsky, K., and Thonstad, J. (1982)
Aluminum Electrolysis, 2nd edn,
Aluminium-Verlag, Dusseldorf.
8 Lantelme, F., Alexopoulos, H., Chemla,
M., and Haas, O. (1988) Electrochim. Acta,
33, 761.
9 Wasserscheid, P. and Welton, T. (2003)
Ionic Liquids in Synthesis, Wiley-VCH,
Verlag GmbH.
10 Welton, T. (1999) Chem. Rev., 99, 2071.
11 Wasserscheid, P. and Keim, W. (2000)
Angew. Chem. Int. Ed., 39, 3772.
12 Earle, M.J. and Seddon, K.R. (2000) Pure
Appl. Chem., 72, 1391.
13 Zhang, S. and Conrad Zhang, Z. (2002)
Green Chemistry, 4, 376.
14 Whitehead, J.A., Lawrence, G.A., and
McCluskey, A. (2004) Green Chem., 6, 313.
15 Maase, M. (2005) in Multiphase Homogeneous Catalysis (eds B. Cornils et al.),
Wiley-VCH, Weinheim, Germany,
p. 560.
16 Chauvin, Y., Olivier, H., Wyrvalski, C.N.,
Simon, L.C., de Souza, R., and Dupont, J.
(1997) J. Catal., 165, 275.
17 Chiappe, C. and Pieraccini, D. (2005) J.
Phys. Org. Chem., 18, 275–297.
18 Abbott, A.P. and McKenzie, K.J. (2006)
Phys. Chem. Chem. Phys., 8, 4265–4279.
19 Xu, W. and Angell, C.A. (2003) Science,
302, 422.
20 Billard, I., Mekki, S., Gaillard, C.,
Hesemann, P., Moutiers, G., Mariet, C.,
Labet, A., and Buenzli, J.G. (2004) Eur. J.
Inorg. Chem., 6, 1190–1197.
21 Jarosik, A., Krajewski, S.R., Lewandowski,
A., and Radzimski, P. (2006) J. Mol. Liq.,
123, 43–50.
22 Widegren, J.A., Saurer, E.M., Marsh,
K.N., and Magee, J.W. (2005) J. Chem.
Thermodyn., 37, 569–575.
23 Abbott, A.P., Harris, R.C., and Ryder,
K.S. (2007) J. Phys. Chem. B, 111,
4910–4914.
24 Yoshizawa, M., Xu, W., and Angell, C.A.
(2003) J. Am. Chem. Soc., 125, 15411.
25 Endres, F. and Zein El Abedin, S. (2006)
Phys. Chem. Chem. Phys., 8, 2101.
26 Zein El-Abedin, S. and Endres, F. (2002)
Phys. Chem. Chem. Phys., 4, 1640.
References
27 Moustafa, E.M., Zein El Abedin, S.,
Shkurankov, A., Zschippang, E., Saad,
A.Y., Bund, A., and Endres, F. (2007)
J. Phys. Chem. B, 111, 4693.
28 Gale, R.J. and Osteryoung, R.A. (1980)
Electrochim. Acta, 25, 1527.
29 Nanjundiah, C., McDevitt, S.F., and Koch,
V.R. (1997) J. Electrochem. Soc., 144, 3392.
30 Nanjundiah, C., Goldman, J.L., McDevitt,
S.F., and Koch, V.R. (1997) Proc.
Electrochem. Soc., 96-25, 301.
31 Baldelli, S. (2005) J. Phys. Chem. B, 27,
109.
32 Rivera-Rubero, S. and Baldelli, S. (2004)
J. Phys. Chem. B, 108, 15133.
13
15
2
Synthesis of Ionic Liquids
Tom Beyersdorff, Thomas J. S. Schubert, Urs Welz-Biermann, Will Pitner,
Andrew P. Abbott, Katy J. McKenzie, and Karl S. Ryder
As is well known in the Ionic Liquids Community 109 to 1018 ionic liquids, binary
and ternary mixtures have been predicted to be – theoretically – achievable. Of
course, this is an incredible number and it will hardly be possible to synthesize
all these liquids and investigate all of them in detail for electrochemical purposes.
This chapter presents an introduction to some ionic liquids that are interesting for
electrochemistry. As the field is still ongoing this chapter can only give an introduction to the principles of ionic liquids synthesis. Section 2.1 briefly summarizes
the major aspects of first generation ionic liquids based on AlCl3 , Section 2.2 gives
a short introduction to the synthesis of air- and water-stable ionic liquids of the
third generation, and Section 2.3 introduces a class of deep eutectic solvents/ionic
liquids based on comparatively well-priced educts such as choline chloride. For a
more detailed introduction to the chemistry of ionic liquids we would like to refer
readers to the 2nd edition of “Ionic Liquids in Synthesis”, ed. by Peter Wasserscheid
and Tom Welton (ISBN: 978–3-527–31239-9).
2.1
Synthesis of Chloroaluminate Ionic Liquids
2.1.1
Introduction
Ionic liquids (IL) are a new class of salt-like materials that are entirely composed
of ions and that are liquid at unusually low temperatures. For the most commonly
used definition of the term ionic liquid the boiling point of water was chosen as a
reference point, most likely for emotional reasons: “The term ionic liquids refers to
compounds consisting entirely of ions and existing in the liquid state below 100 ◦ C.” In
many cases the melting point is even below room temperature.
The history of ionic liquids began with the synthesis of ethylammonium nitrate
reported in 1914 by Walden [1]. This material is probably the first described in the
literature that fulfills the definition of ionic liquids used today. In this context it
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
16 2 Synthesis of Ionic Liquids
should be noted that at that time Walden had of course no idea of this definition
or the whole concept of ionic liquids. Consequently, it is not surprising that at the
time no attention was paid to the potential of this class of materials.
A major breakthrough was achieved in 1951 with the report of Hurley and Wier.
They noticed that a mixture of N-ethylpyridinium bromide (EtPyBr) and AlCl3 with
a eutectic composition of 1:2 {X(AlCl3 ) = 0.66}1) of EtPyBr to AlCl3 became liquid
at unusually low temperatures [2]. They investigated these melts with regard to
their potential use in the electrodeposition of aluminum at ambient temperature
[3]. Several studies were carried out on this system, however, its use was very limited
since it is only liquid at a mole fraction of X(AlCl3 ) = 0.66 and the ease of oxidation
of the bromide ion limits the electrochemical stability. In the following years the
main interest in ionic liquids was focused on electrochemical applications [4–6].
In 1978 Osteryoung and coworkers replaced EtPyBr with N-butylpyridinium
chloride (BuPyCl) and found that the properties of the resulting ionic liquids improved significantly [7, 8]. The new chloroaluminate melts were found to be liquid
at room temperature over a composition range from X(AlCl3 ) = 0.66 to 0.43. In
addition the anodic limit had improved by changing from bromide to chloride. The
main disadvantage of these systems was the relative ease of both chemical and
electrochemical reduction of the buytlpyridinium cation [9]. Wilkes and coworkers
performed MNDO (modified neglect of diatomic overlap) calculations on a variety
of organic cations in 1982 and found that N,N ′ -dialkylimidazolium cations are more
stable than the N-butylpyridinium cation due to the higher electron affinity of these
cations [10]. Many of the melts resulting from mixing N,N ′ -dialkylimidazolium
halides with AlCl3 even displayed lower melting points than the N-butylpyridiniumbased ionic liquids. In the case of the 1-ethyl-3-methyl-imidazolium chloride/AlCl3
mixtures the liquid range at room temperature extends from X(AlCl3 ) = 0.66 to
0.30 [11]. Further research on air- and water-stable anions and new cations has
been carried out during the past years resulting in more than 1500 materials being
described in the literature today [12].
The first part of this chapter focuses on the synthesis and properties of the socalled “first generation of ionic liquids”, the haloaluminate-based ionic liquids and
in particular on those of chloroaluminate melts.
2.1.2
Synthesis of Room-temperature Chloroaluminate-based Ionic Liquids
2.1.2.1 Introduction
The synthesis of haloaluminate-based ionic liquids from halide salts and aluminum
Lewis acids (most commonly AlX3 ; X=Cl, Br) can generally be split into two steps: (i)
fomation of the desired cation by the reaction of a trialkylamine, trialkylphosphine
or dialkylsulfide with a haloalkane, and (ii) formation of the haloaluminate anion
by addition of an appropriate aluminum halide to this salt (Scheme 2.1).
1) The composition of haloaluminate ionic liquids is often described by the mole fraction
of AlCl3 X(AlCl3 ) present in the mixture.
2.1 Synthesis of Chloroaluminate Ionic Liquids 17
Scheme 2.1 General synthesis route to haloaluminate-based ionic liquids.
Nowadays, as many halide salts are commercially available at reasonable prices,
often only the second step is required.
The most commonly used groups of cations are presented in Figure 2.1.
The following section focuses on the quaternization reaction of 1-alkylimidazoles
since these are the most commonly used starting materials for ionic liquids and have
dominated ionic liquids research over the last twenty years. However, the general
method for the quaternization reaction is similar for pyridines [13], isoquinolines
[14], 1-methylpyrrolidine [15], trialkylamines [16], phosphines [17] and sulfides [18].
2.1.2.2 The Quaternization Reaction
From a practical point of view ionic liquids have no significant vapor pressure. As a
consequence, their purification using conventional methods is extremely difficult.
Thus, it is recommended to remove as many impurities as possible from the
starting materials and to use synthetic procedures that produce as few side products
as possible, or allow their easy separation from the final product. In addition, all
starting materials should be dried prior to use considering the water-sensitive
nature of many of the products.
All reagents used for the synthesis of cations should be purified according to
literature procedures before use [19]. Amines such as 1-alkylimidazoles or pyridines
are typically distilled from sodium hydroxide or calcium hydride if dry amines are
required and stored under dry nitrogen or argon at 0 ◦ C. Haloalkanes are washed
with sulfuric acid until no further color is extracted into the acid layer and then
neutralized with NaHCO3 and deionized water prior to distillation from CaCl2 . All
solvents used in the syntheses should be dried and distilled prior to use. In order
to obtain colorless halide salts it is recommended to perform all reactions under a
protective atmosphere of a dry inert gas in order to exclude moisture and oxygen
from the reaction.
In order to obtain colorless chloroaluminate liquids it is recommended to sublime
the AlCl3 several times prior to use after the addition of sodium chloride and
aluminum wire [8].
The synthesis of the cation is typically performed by alkylation of an amine,
phosphine or sulfide, most commonly using an alkyl halide [ ]. In most cases the
reaction is carried out with chloro-, bromo- and iodoalkanes as readily available
alkylating reagents, with the reaction conditions becoming more gentle changing
from chloride to bromide to iodide, as can be expected for nucleophilic substitution
Fig. 2.1 Examples of cations commonly used for the synthesis of ionic liquids.
18 2 Synthesis of Ionic Liquids
Scheme 2.2 Quaternization reaction of 1-alkylimidazoles.
reactions. Onium fluorides cannot be synthesized in this manner due to the poor
leaving-group qualities of fluoride anions (Scheme 2.2).
A typical lab scale alkylating reaction is performed in a round-bottomed flask
equipped with a reflux condenser and a dropping funnel with a nitrogen or argon
inlet. The alkylating reagent is dissolved in the solvent and the amine is added
dropwise. After complete addition, the reaction mixture is heated until all of the
amine has been consumed. The reaction conditions for the quaternization are
strongly dependent on the haloalkanes used, with the chloroalkanes being the least
reactive and the iodoalkanes the most. In general chloroalkanes have to be heated
to 80 ◦ C for several days to ensure complete reaction, whereas reactions employing
bromoalkanes are usually complete after 24 hours at lower temperatures, between
50 and 60 ◦ C. Alkylation reactions with iodoalkanes can often be performed at
room temperature with exclusion of light, since iodoalkanes and the resulting
iodide salts are light-sensitive. Taking safety aspects into account, care has to be
taken with large-scale reactions employing bromoalkanes, as such reactions are
strongly exothermic with increased reaction rates. Besides safety considerations,
high thermal stress can also result in discoloration of the final product.
The reactivity of haloalkanes in alkylation reactions also decreases with increasing
chain length. In general, syntheses of salts with short alkyl substituents are more
complex due to the low boiling points of the haloalkanes. The most frequently
used halide salt in this field, 1-ethyl-3-methylimidazolium chloride ([EMIM] Cl),
is typically synthesized in an autoclave with the chloroethane cooled to below its
boiling point (12 ◦ C) before addition.
In general, the use of solvents is not inevitably necessary as the reagents are
liquid and mutually miscible, while the halide salts are usually immiscible with
the starting materials. Nevertheless, solvents are often used to keep the reaction
homogeneous and thus to ensure better heat transfer within the reaction mixture. Examples of solvents include the haloalkane itself [10], dichloromethane,
acetonitrile, 1,1,1-trichloroethane [20], ethyl acetate [21] and toluene [22]. These solvents can be divided into two classes: those that are miscible with the product salt
(dichloromethane, acetonitrile) and those that are immiscible with the halide salt
product (1,1,1-trichloroethane, toluene, ethyl acetate). Reactions performed in the
former solvents result in homogenous reaction mixtures from which the product
can be precipitated, in many cases, by addition of an immiscible co-solvent. For
reactions in the latter solvents, removal of the solvent and unreacted starting materials can be achieved by simple decantation and washing of the product with an
immiscible solvent, as the product is generally denser than the solvents and starting
materials. Purification of the halide salts is in all cases dependent on their state
of aggregation. In many cases the halide salts are solids at room temperature and
2.1 Synthesis of Chloroaluminate Ionic Liquids 19
can be recrystallized from mixtures of dry acetonitrile and ethyl acetate. However,
if the product does not crystallize, it is advisable to wash the oily product with an
immiscible solvent to remove excess starting materials. In all cases it is necessary to
remove all excess starting materials, solvents and moisture by heating the product
salt under vacuum. Care has to be taken at this stage, as overheating can result
in the decomposition of the product via retro-alkylation. It is recommended not to
heat the halide salt at temperatures higher than 80 ◦ C.
An alternative approach to the reaction conditions described above employs
microwave irradiation for the quaternization reaction of 1-methylimidazole with
various haloalkanes and 1,ω-dihaloalkanes [23]. High yields and acceptable purities
can be obtained in short reaction times (minutes instead of hours) and scaling up
this technology to an industrial scale can easily be achieved.
As a new class of materials, ionic liquids require special analytical methods.
In the case of imidazolium halides and similar compounds the most common
impurities are amines, alkyl halides and of course water. Seddon et al. described
a method for the detection of residual amines using the strong UV absorbance of
copper tetramine complexes. These complexes are readily formed by the addition of
Cu2+ ions [24]. The detection of both amines and alkyl halides is possible by NMR
spectroscopy but with limited resolution [25]. By far the most powerful analytical
method is liquid chromatography combined with UV detection. This sensitive
method allows the detection of traces of amines and halides [26]. Unreacted amines
can be also detected by ion chromatography combined with a suppressor module.
In this case detection is achieved using a continuous flow conductivity cell since
amines are protonated and thus detectable. For traces of other ionic impurities
ion chromatography is also the most powerful analytical tool [27]. Finally, residual
water can be quantified using Karl Fischer titration or coulometry [28].
2.1.2.3 Chloroaluminate Synthesis
Treatment of a quaternary halide salt Q+ X− with a Lewis acid MXn results in
the formation of a salt with the composition Q+ MXn+1 − . In general, more than
just one anion species is formed, depending on the relative proportions of MXn
and the halide salt Q+ X− . A representative example is the reaction of 1-ethyl-3methylimidazolium chloride [EMIM]Cl with AlCl3 (Scheme 2.3).
If the mole fraction X(AlCl3 ) is less than 0.5 in the final product, the ionic liquids
are basic, as chloride ions are present which are not bound to aluminum and which
act as Lewis bases. For mole fractions X(AlCl3 ) > 0.5 an excess of Lewis acid AlCl3
is present and the melts are acidic. If the mole fraction X(AlCl3 ) = 0.5 the salts
are neutral as all of the chloride ions are bound to aluminum and the only species
present is the [AlCl4 ]− ion. However, as a consequence of the autosolvolysis of
Scheme 2.3 Reaction between [EMIM]Cl and AlCl3 .
20 2 Synthesis of Ionic Liquids
Scheme 2.4 Autosolvolysis of AlCl4 melts.
AlCl4 − , Cl− and [Al2 Cl7 ]− species are always present in neutral liquids (Scheme
2.4) [ ].
A detailed description of the analysis of the chloroaluminate species in these
ionic liquids is given by Welton et al. [29].
It has to be mentioned that chloroaluminate melts are not the only Lewis acidbased ionic liquids produced in this manner. Other examples include for example
AlEtCl2 [30], AlBr3 [31], BCl3 [32], CuCl [33], SnCl2 [34], FeCl3 [35], ZnCl2 [36]. The
preparation of these salts is similar to that described for the [AlCl4 ]− salts. Even
the treatment of halide salts with metal halides or metal oxides that are not typical
Lewis acids has been used to synthesize ionic liquids. Examples of these salts are
[EMIM]2 [MCl4 ] (M=Co, Ni) [37], [EMIM]2 [VOCl4 ] [38], [BMIM][CrO3 Cl] [39].
The most common method for the synthesis of chloroaluminate-based ionic
liquids is a solid-phase synthesis by mixing AlCl3 and a quaternary halide Q+ X−
salt under vigorous stirring. This type of reaction should be carried out using
Schlenk techniques or, preferably, in a glove box. Since the ionic liquid is formed
directly in an exothermic reaction on contact of the two starting materials, care
should be taken upon mixing the reagents. Although the starting materials as well
as the products are relatively thermally stable in general, local overheating can result
in decomposition and darkening of the ionic liquid. To prevent this, the reaction
vessel should be cooled. An important point to avoid overheating is to add one
starting material to the other in small portions in order to allow the reaction heat
to dissipate. In addition the resulting ionic liquids should be stored under argon in
a Schlenk-type flask or in a glove box until use.
However, if a glove box is not available for the synthesis, the reaction can also
be performed in a dry, inert solvent which covers the reaction mixture and protects
it from hydrolysis. An advantage of this procedure is that the solvent, which is
typically an alkane, can also react as a heat carrier in the exothermic reaction. After
completion of the reaction the ionic liquid forms a second layer below the solvent.
The solvent can be removed by simple distillation before use of the ionic liquid.
However, the ionic liquid will be contaminated with the organic solvent, which has
to be removed under vacuum.
Another method involves microwave irradiation. It has been described for the
synthesis of 1,3-dialkylimidazolium tetrachloroaluminates [40]. This method precludes the use of volatile organic solvents and is faster, more efficient and also
ecofriendly, affording high yields of the desired products.
As mentioned above, purification of the resulting ionic liquids cannot be achieved
by distillation of the products since these materials show no significant vapor pressure. In most cases AlCl3 -based ionic liquids contain traces of oxo ion impurities
such as [AlOCl2 ]− as major impurities, especially if water and oxygen are not totally excluded during synthesis. As shown by 17 O NMR experiments a complex
set of equilibria is then present [41]. These impurities can easily be removed by
2.2 Air- and Water-stable Ionic Liquids 21
bubbling phosgene [42] or, considering the high toxicity of phosgene, triphosgene
[43] through the ionic liquid. The by-product formed in this reaction is CO2 , which
can be easily removed under vacuum. Further purification of room temperature
haloaluminate-based ionic liquids is not recommended, as these materials are extremely sensitive towards moisture and must be handled either in vacuum or under
an inert gas atmosphere. Although classic Schlenk techniques can be used to handle
these materials, working in a glove box is recommended.
Analysis of haloaluminate ionic liquids is much more limited than that of other
ionic liquids. The most important analytical technique is surely NMR spectroscopy.
The determination of residual water is difficult because of the instability of these
materials. Hence it is crucial to work accurately to achieve the best results.
To summarize the most important points for the synthesis of pure and colorless
ionic liquids it is recommended to:
Ĺ Purify all starting materials before use
Ĺ Exclude oxygen and moisture from the reactions by working in a dry inert atmosphere to prevent darkening of the ionic liquids
Ĺ Keep the reaction temperatures as low as possible, as overheating often results
in discoloration of the products
Ĺ Use Schlenk techniques or work in a glove box, as the chloride and bromide
salts are highly hygroscopic and the chloroaluminate melts are highly moisture
sensitive.
The importance of these first generation ionic liquids for metal deposition is
summarized in Chapter 4.1.
2.1.3
Physical Data of Haloaluminate-based Ionic Liquids
A selection of physical data of selected haloaluminate-based ionic liquids is given
in Table 2.1.
2.2
Air- and Water-stable Ionic Liquids
2.2.1
Introduction
For forty years following the introduction of haloaluminate-based ionic liquids by
Hurley and Wier, [44, 45] the majority of research in this field was carried out
on systems which were reactive with air and, more specifically, with water. The
difficulty of working with these materials, using elaborate Schlenk-line airless techniques or expensive and difficult-to-maintain controlled-atmosphere glove boxes,
had the effect of limiting the research to four American-based research groups,
mostly funded by the US Air Force [46]. Well aware of this limitation, John Wilkes
and coworkers made the decision to substitute the reactive haloaluminate anion
Composition of the IL System
Cation
Anion
Viscosity
(mPa s)
Conductivity
(mS cm−1 )
Molar
conductivity
Density
(g cm−3 )
Ref.
[Al2 Cl7 ]−
[Al2 Cl7 ]−
[AlCl4 ]−
Cl− ,
[AlCl4 ]−
[Al2 Br7 ]−
Br− ,
[AlBr4 ]−
[AlCl4 ]− ,
[Al2 Cl7 ]−
[AlCl4 ]−
Cl− ,
[AlCl4 ]−
[Al2 Cl7 ]−
[AlCl4 ]−
[Al2 Cl7 ]−
[AlCl4 ]−
[Al2 Cl7 ]−
[Al2 Cl7 ]−
[Al2 Cl7 ]−
[Al2 Cl7 ]−
17
14
18
47
15.0
15.0
23.0
6.5
4.26
4.46
4.98
1.22
1.404
1.389
1.294
1.256
[10]
[10]
[10]
[10]
32
67
5.8
5.7
1.89
1.15
2.219
1.828
[31]
[31]
18
11.0
2.94
1.351
[10]
27
12.0
3.3
2.79
1.262
[10]
[10]
19
27
24
38
21
18
18
21
9.2
10.0
6.0
54.0
8.1
10.0
8.0
6.7
3.04
2.49
2.32
1.50
2.23
2.91
2.47
2.18
1.334
1.238
1.252
1.164
1.441
1.408
1.375
1.346
[10]
[10]
[10]
[10]
[36]
[36]
[36]
[36]
34.0/66.0
34.0/66.0
50.0/50.0
60.0/40.0
(mol%)
(mol%)
(mol%)
(mol%)
[MMIM]Cl/AlCl3
[EMIM]Cl/AlCl3
[EMIM]Cl/AlCl3
[EMIM]Cl/AlCl3
[MMIM]+
[EMIM]+
[EMIM]+
[EMIM]+
34.0/66.0
60.0/40.0
(mol%)
(mol%)
[EMIM]Br/AlBr3
[EMIM]Br/AlBr3
[EMIM]+
[EMIM]+
40.0/60.0
(mol%)
[PMIM]Cl/AlCl3
[PMIM]+
50.0/50.0
60.0/40.0
(mol%)
(mol%)
[PMIM]Cl/AlCl3
[PMIM]Cl/AlCl3
[PMIM]+
[PMIM]+
34.0/66.0
50.0/50.0
34.0/66.0
50.0/50.0
33.3/66.7
33.3/66.7
33.3/66.7
33.3/66.7
(mol%)
(mol%)
(mol%)
(mol%)
(mol%)
(mol%)
(mol%)
(mol%)
[BMIM]Cl/AlCl3
[BMIM]Cl/AlCl3
[BBIM]Cl/AlCl3
[BBIM]Cl/AlCl3
[MP]Cl/AlCl3
[EP]Cl/AlCl3
[PP]Cl/AlCl3
[BP]Cl/AlCl3
[BMIM]+
[BMIM]+
[BBIM]+
[BBIM]+
[MP]+
[EP]+
[PP]+
[BP]+
22 2 Synthesis of Ionic Liquids
Table 2.1 Physical data of selected ionic liquids.
2.2 Air- and Water-stable Ionic Liquids 23
systems with less reactive anions, under the belief that other anions could also produce low-melting organic salts [47]. Perhaps they were aware of the singular work
by Walden to generate low melting, high conductivity salts [48], work which is often
used to mark him as the discoverer or inventor of ionic liquids. However, Walden
based his “molten salts” upon protonated primary amines such as ethylamine, and
such mixtures of organic bases with mineral and organic acids will always exist as
mixtures of the acid, the base and salt formed through their neutralization. Such
equilibrium mixtures are well known to be thermally unstable, due to the vapor
pressure of the two neutral components [49]. The ease with which the protonated
cation can be reduced to yield hydrogen gas also limits the usefulness of these
materials in electrochemical applications. The non-chloroaluminate ionic liquids
introduced by Wilkes have the advantages of high thermal and electrochemical
stability (with respect to Walden’s acid–base equilibrium mixtures) and ease of
handling under ambient, humid conditions (as compared to Hurley and Wier’s
haloaluminate ionic liquids).
The methathesis route employed by Wilkes in the production of these materials
has generally been followed for the majority of air- and moisture-stable ionic liquids.
In brief (Scheme 2.5), an organic base (such as N-methylimidazole, pyridine or Nmethylpyrrolidine) is alkylated using a haloalkane to generate an organic halide
salt. Anion exchange is carried out, generally in water, with the appropriate acid
or metal salt. The ionic liquid is extracted from the aqueous salt into an organic
phase, and the halide impurities removed through repeated washings with water.
The more hydrophilic the ionic liquid, the more difficult it is to purify, as extraction
of halides with water is complicated by loss of the ionic liquid to the aqueous phase.
Scheme 2.5 General synthetic route to producing air- and moisture-stable ionic liquids.
24 2 Synthesis of Ionic Liquids
Fig. 2.2 Structure, full name and abbreviations for the anions discussed in this section.
The discovery made by Wilkes in the early nineties was completely dependent
upon a change in the anionic systems: the cation components of the haloaluminate
systems previously in use were not the cause of their reactivity with water and the
pyridinium and imidazolium cations remain key components of air- and moisturestable ionic liquids under investigation today. Therefore the focus of this section will
be air- and moisture-stable anionic systems (Figure 2.2), with the cation relegated
to the rôle of junior partner in an ionic couple. Due to their ease of handling,
this report will also focus on those anion–cation combinations which yield roomtemperature ionic liquids (RTILs), even though operation at room-temperature is
not a prerequisite for a commercially-viable electroplating bath.
2.2.2
Tetrafluoroborate and Hexafluorophosphate-based Ionic Liquids
Wilkes launched the field of air- and moisture-stable ionic liquids by introducing five
new materials, each containing the 1-ethyl-3-methylimidazolium cation [EMIM]+
with one of five anions: nitrate [NO3 ]− , nitrite [NO2 ]− , sulfate [SO4 ]2− , methyl
carbonate [CH3 CO2 ]− and tetrafluoroborate [BF4 ]− [47]. Only the last two materials
had melting points lower than room temperature, and the reactive nature of the
methyl carbonate would make it unsuitable for many applications. This led to
the early adoption of [EMIM][BF4 ] as a favored ionic liquid, which has since been
the subject of over 350 scientific publications. One of the first appeared in 1997 [50],
reporting the investigation of [EMIM][BF4 ] as the electrolyte system for a number of
processes, including the electrodeposition of lithium (intended for use in lithium
ion batteries).
2.2 Air- and Water-stable Ionic Liquids 25
The [BF4 ]− anion is frequently used in battery electrolyte formulations, and it
was not long before other anions from this branch were investigated for their
capacity to make ionic liquids. The combination of hexafluorophosphate [PF6 ]−
with [EMIM]+ produced a salt with a reported melting point of 58–60 ◦ C [51].
This was too high a melting point for most researchers to bother working with,
but in 1995 Chauvin, Mussmann and Olivier reported the use of the analogous
[BMIM][PF6 ], which is liquid at room temperature, as well as [BMIM][BF4 ] [52].
These two RTILs would dominate ionic liquid publications for the next decade. The
preference for [BMIM][BF4 ] over [EMIM][BF4 ] can probably be explained by the fact
that the synthesis of [BMIM]Cl is an easier process than the synthesis of [EMIM]Cl,
which requires a pressurised reaction vessel, and by the high water-solubility of
[EMIM][BF4 ], which makes it much more difficult to purify than [BMIM][BF4 ].
Another reason was the experimental symmetry afforded by switching from an
ionic liquid which is completely miscible with water ([BMIM][BF4 ]) to one which
formed biphasic aqueous mixtures ([BMIM][PF6 ]).
Other cation combinations with [BF4 ]− and [PF6 ]− have proved uninteresting in
the study of electrochemical systems. Although N-butylpyridinium tetrafluoroborate [bpyr][BF4 ]− is known to be a RTIL [53], the lower electrochemical stability of
pyridinium-based cations relative to imidazolium limits their electrochemical applicability. On the other hand, pyrrolidinium-based cations are known to be more
electrochemically stable than imidazolium salts, N-alkyl-N-methylpyrrolidinium
salts of [BF4 ]− and [PF6 ]− are made less attractive to researchers by the fact that
they are solids at room temperature [54, 55]. Therefore, most of the electrochemical investigations of ionic liquids containing [BF4 ]− and [PF6 ]− have focused on
[BMIM][PF6 ], [BMIM][BF4 ] and, to a lesser extent, [EMIM][BF4 ].
Concerns about the stability of [BF4 ]− and [PF6 ]− and ionic liquids which contain
these anions have led many researchers to turn their backs on these materials.
Anecdotal evidence of glassware shattering during heating and vacuum drying is
commonplace, but more rigorous investigations confirm the rumours that [BF4 ]−
and [PF6 ]− -based ionic liquids hydrolyse to generate HF, a corrosive and toxic
material [56]. Experiments performed by Merck KGaA demonstrate this instability
with respect to ionic liquids based on other fluorinated anions (Figure 2.3). Despite
such warnings, however, research continues on these materials for a number of
reasons: the large amount of baseline data on [BMIM][PF6 ], [BMIM][BF4 ] and
[EMIM][BF4 ] which is available from prior experimentation and publications; the
ease with which these materials can be produced; and their low cost relative to other
more complex and stable anion systems. In addition, [BF4 ]− and [PF6 ]− -based ionic
liquids can possess properties which, for a given application, provide a superior
performance to other ionic liquids [57].
2.2.3
Triflate- and Trifluoroacetate-based Ionic Liquids
Small, fluorinated organic anions, such as trifluoromethanesulfonate (or triflate)
and trifluoroacetate, were quickly considered as alternatives to inorganic fluorinated phosphates and borates. Carlin and coworkers were the first to report
26 2 Synthesis of Ionic Liquids
Fig. 2.3 A comparison of the hydrolytic stability of four 1-hexyl-3-methylimidazolium
ionic liquids [HMIM]X ionic liquids. To
25 g of each ionic liquid, 7.1 mol% water
was added. These solutions were heated to
60 ◦ C and the fluoride content measured
once an hour for eight hours. All measurements performed at Merck KGaA, Darmstadt, Germany.
an investigation using 1-ethyl-3-methylimidazolium triflate [EMIM][OTF], looking
at different RTILs for use in an ongoing battery project [58]. The 1-butyl-3methylimidazolium triflate [BMIM][OTF] was suggested as an electrolyte component for dye-sensitised solar cells (DSSCs) [59], as was 1-ethyl-3-methylimidazolium
trifluoroacetate [EMIM][ATF] [60]. These materials in general form ionic liquids
with relatively low viscosities, and are characterized by reasonably large electrochemical windows (though not comparable with the inorganic fluorinated anions,
see Table 2.2) [61]. The sulfate and carboxylate functional groups make them
strongly coordinating anions, although the electron-withdrawing trifluoromethane
Table 2.2 Dependence of selected physicochemical properties (at 20 ◦ C)
of ionic liquids [EMIM]X on the anion X− .
Ionic liquid
[EMIM][BF4 ]
[EMIM][FAP]
[EMIM][ATF]
[EMIM][OTF]
[EMIM][NTF]
[EMIM][SCN]
[EMIM][DCA]
[EMIM][TCM]
[EMIM][TCB]
a
Density/g cm−3
Dynamic
viscosity/mPa s
Specific
conductivity/mS cm−1
ERed–Ox /V
1.30
1.72
1.30
1.39
1.52
1.15
1.08a,b
1.11b
1.04
60
75
41
52
40
44
16a (17)b
18b
20
11
4
5
7
8
14
28a (27)b
18b
13
5.2
6.5
3.4
4.1
6.3
3.2
3.5a,b
3.5b
4.5
At 25 ◦ C [J. Phys. Chem. B, 111(18), 2007].
At 22 ◦ C [Inorganic Chemistry, 43(4), 2004].
All other measurements performed at Merck KGaA, Darmstadt, Germany.
b
2.2 Air- and Water-stable Ionic Liquids 27
component makes them much less basic than their methanesulfonate and acetate
analogs. While the decreased electrochemical stability may be viewed as a negative
when considering such ionic liquids as potential electroplating baths, their lower
viscosities are an attractive feature while their ability to coordinate may also work
in their favor by increasing the solubility of metal salts.
2.2.4
Bistriflamide-based Ionic Liquids
Lithium bis(trifluoromethylsulfonyl)amide Li[NTF] has been widely recognised as
a possible component in battery electrolyte compositions since 1990 [62]. A readily
available material, Li[NTF] can be converted into [NTF]− -based ionic liquids through
a very simple ion exchange step in an aqueous mixture because most [NTF]− based ionic liquids form biphasic aqueous mixtures, as reported by Bonhôte and
coworkers in 1996 [59, 60]. The ionic-liquid-rich phase is easily separated and can be
purified to a high level through simple washing with water. In the same year, Watanabe and Mizumura reported ionic liquids based upon Li[NTF] in combination with
lithium acetate and triethyl methyl ammonium benzoate. In addition to their ease
of preparation, [NTF]− -based ionic liquids are generally characterised by higher
electrochemical and thermal stability, lower viscosity and higher conductivity
than ionic liquids based on [BF4 ]− and [PF6 ]− (Table 2.2). This collection of favorable properties is one reason why there is currently great interest in this
class of ionic liquids. While initial interest in [NTF]− -based ionic liquids focused on imidazolium salts, it soon became clear that a broader range of cations
could be paired with [NTF]− to generate RTILs [63–65]. For electrochemists, the
higher electrochemical stability of tetraalkylphosphonium, tetraalkylammonium,
N,N-dialkylpyrrolidinium and N,N-dialkylpiperidinium [NTF]− -based RTILs make
them attractive alternatives to 1,3-dialkylimidazolium and N-alkylpyridinium salts.
This is especially true for applications involving the electrodeposition of active metals, where reactions between the electrodeposited metal and the ionic liquid plating
bath should be avoided.
More recently, concerns have arisen with respect to the assumed stability of
[NTF]− -based ionic liquids. In an investigation of the electrochemical behavior of
lithium in such ionic liquids, MacFarlane and coworkers reported the decomposition of [NTF]− , due either to unwanted reactions between the active metal surface
of the electrode or the electrochemical reduction of the anion at the negative potentials required for lithium reduction [66]. This potential instability of [NTF]− to
reduction had been predicted by Makato and coworkers two years earlier, using
ab initio molecular orbital calculations [67]. In addition, their possession of a negligible vapor pressure, which was previously assumed to be a general property of
all ionic liquids, has been called into question with reports of the distillation of
[NTF]− -based ionic liquids [68]. Moreover, while [NTF]− has been demonstrated to
be much more stable than [BF4 ]− and [PF6 ]− to hydrolysis (Figure 2.3), ab initio
calculations suggest that [NTF]− -based ionic liquids may be much more volatile
than those based on [BF4 ]− and [PF6 ]− [69]. In addition, [NTF]− -based ionic liquids, though often classified as “hydrophobic” due to their formation of biphasic
28 2 Synthesis of Ionic Liquids
Table 2.3 A comparison of the relative hydrophobicity of eight ionic
liquids.
Ionic liquid
[HMIM][BF4 ]
[HMIM][TCB]
[HMIM][PF6 ]
[Ph3 t][NTF]
[HPYR][NTF]
[HMIM][NTF]
[HMPL][NTF]
[HMIM][FAP]
[Ph3 t][FAP]
[HMPL][FAP]
Water content
of IL/wt%
IL content of
water/ppm
18.1
5.39
1.84
1.58
1.13
1.12
0.900
0.195
0.180
0.114
7.27
0.44
—
—
0.25
0.13
—
0.02
—
—
aqueous mixtures, are demonstrably soluble in water (Table 2.3); even a a small loss
of [NTF]− to the water-rich phase during synthesis could result in circumstances
where the economics of commercial applications are called into question by the
loss of this high-cost component.
2.2.5
Trispentafluoroethyltrifluorophosphate-based Ionic Liquids
The anion trispentafluoroethyltrifluorophosphate [FAP]− belongs to the broader
class of perfluoroalkylphosphate-based anions first reported in the 1960s [70].
Merck KGaA began investigating the use of Li[FAP] as a component in battery
electrolytes in 2001, as a replacement for Li[PF6 ] [71]. The hydrolytic instability
of the hexafluorophosphate anions is due to the facial protonation of the fluorine
atom, followed by HF elimination and further reaction with water. This problem
is addressed by replacement of some of the fluorine atoms by hydrophobic
perfluoroalkyl groups, reducing the rate of hydrolysis through steric hindrance of
attacks on the phosphate center.
Recently, Merck KGaA developed a convenient method for the synthesis of
[FAP]− -based ionic liquids as replacement for [PF6 ]− -based ionic liquids [72]. Like
their [PF6 ]− analogs, [FAP]− -based ionic liquids form biphasic aqueous mixtures
and can be separated and recovered easily from aqueous reaction mixtures. They
can be easily obtained with very low water and chloride content by washing with
water followed by heating under reduced pressure. The hydrolytic stability (Figure
2.3) and electrochemical stability (Table 2.2) of [FAP]− and its ionic liquids are
superior to [PF6 ]− and [BF4 ]− and comparable with [NTF]− .
2.2.6
Cyano-based Ionic Liquids
A family of ionic liquids has developed around anions containing a central element
coordinated by one or more cyano groups. The stability of the carbon–nitrogen triple
2.2 Air- and Water-stable Ionic Liquids 29
bond of the cyano group, its high electronegativity and ability to increase charge
delocalization combine to give this family a unique set of chemical properties. While
generally electrochemically stable, they are also capable of strong coordination and
solvation of polar hydrogen-bond donating materials such as cellulose and sugars
[73, 74]. Cyano ligands are prone to polymerization during decomposition, and
preliminary investigations have indicated that the main concern with this type of
anion, the formation of HCN during their thermal decomposition, is negligible
[75, 76]. They are generally much more hydrophilic than fluorinated anions.
The simplest example is the thiocyanate anion [SCN]− . Thiocyanate-based ionic
liquids, such as [EMIM][SCN] [77] show good thermal stability, low melting points
and electrochemical stability sufficient for a wide range of electrochemical applications, most especially for dye-sensitised solar cells [78–80]. Their ability to dissolve
metal thiocyanates in high quantities is significant [81]. Nitrogen can coordinate
two cyano ligands, to form the dicyanamide anion [DCA]− , which can form ionic
liquids [82]. Certainly the popularity of [EMIM][DCA] is its extremely low viscosity
of 17 cP (extremely low for an ionic liquid, that is). This anion can form RTILs with
a broad range of electrochemically stable cations, including imidazoliums, ammoniums and pyrrolidiniums [83]. Because of these properties, dicyanamide-based
ionic liquids have been considered for a wide variety of electrochemical applications [84], most especially for dye-sensitized solar cells (DSSC) [85, 86]. While the
tricyanomethide anion [TCM]− can also be used to make RTILs [87] and has also
been investigated for photovoltaic applications [88], this anion system has been
much less investigated than thiocyanates and dicyanamides.
The synthesis of the tetracyanoborate anion [TCB]− was first described by Bessler
[89, 90], but only the improvement of the sinter process of the key intermediate
potassium [TCB] [91] has made this material available in reasonable amounts and
thus allowed the synthesis of [TCB]− -based ionic liquids [92]. The low viscosity
(20 cP at 20 ◦ C) and thermal and chemical robustness led to the use of [EMIM][TCB]
as an electrolyte for DSSC [93].
2.2.7
Effect of Anion on Ionic Liquid Physicochemical Properties
The choice of anion will have a known effect on the physicochemical properties of
the ionic liquid. To demonstrate the anion effect, selected data on properties of general interest to electrochemists (density, viscosity, conductivity and electrochemical
window) have been gathered in Table 2.2. In each case, the anion is paired with
the same cation: 1-ethyl-3-methylimidazolium. Certain trends from this data can
be generalized, as well as in other collections of such data (for example, see Ref.
[61] and references therein), that hold true regardless of the identity of the cation.
For example, the effect of the anion on density follows the trend:
[TCB]− < [DCA]− < [TCM]− < [SCN]− ≪ [BF4 ]− < [ATF]− < [OTF]−
≪ [PF6 ]− < [NTF]− <<< [FAP]−
Anions which are particularly large or which strongly coordinate tend to have
increased densities.
30 2 Synthesis of Ionic Liquids
The trend for viscosity is:
[DCA]− ∼ [TCB]− < [TCM]− ≪ [NTF]− < [ATF]− < [SCN]− < [OTF]−
< [BF4 ]− < [FAP]− < [PF6 ]−
and for specific conductivity:
[DCA]− > [TCM]− > [TCB]− ∼ [SCN]− ≫ [BF4 ]− > [NTF]−
∼ [ATF]− ∼ [OTF]− > [FAP]− > [PF6 ]−
likewise appearing to be related to anion size and coordination strength, as well as
the amount of charge delocalization.
Ionic liquids which form biphasic aqueous mixtures are often classified as hydrophobic, regardless of the fact that this phase behavior is temperature dependent
and that ionic liquids are, in general, hygroscopic. The hydrophobicity of an ionic
liquid used as an electroplating bath is an important factor if exclusion of water from
the bath is important: the more hydrophobic the ionic liquids, the lower the water
content will be upon saturation. The relative hydrophobicity of an ionic liquid is a
factor of both the anion and cation, as has been demonstrated by research carried out
by Merck KGaA (Table 2.3). To make this comparison, equal volumes of an ionic liquid and water were mixed for 2 h at room temperature, then allowed to separate into
two phases. The water content of the ionic-liquid-rich phase was then determined
with Karl–Fischer titration, and the ionic liquid content of the water-rich phase was
determined using high performance liquid chromatography (HPLC) (1-hexyl-3methylimidazolium [HMIM]+ and N-hexylpyridinium [HPYR]+ ) or ion chromatography (IC) (1-hexyl-1-methylpyrrolidinium [HMPL]+ ). No satisfactory method was
found for quantifying the trihexyl-tetradecylphosphonium bistriflamide [Ph3 t][NTF]
content of the water-rich phase. The clear trend of increasing hydrophobicity for
the four cations evaluated is
[Ph3 t]+ < [HMIM]+ ∼ [HPYR]+ < [HMPL]+
while the trend for the five anions is
[BF4 ]− < [TCB]− < [PF6 ]− < [NTF]− < [FAP]−
In summary, there are many anion types which offer useful properties for the
creation of an electroplating medium. Choices must be made regarding electrochemical stability, relative hydrophobicity, the ability to coordinate metal salts and
the mass transport properties of viscosity and conductivity.
2.2.8
Concluding Remarks
In the years following Wilkes introduction of air- and moisture-stable ionic liquids,
these materials have been transformed from laboratory curiosities which each
2.3 Eutectic-based Ionic Liquids
Table 2.4 Standard purity grades of ionic liquids available from Merck
KGaA.
Merck purity grade
“for synthesis”
“high purity”
“ultra pure”
> 98
< 1000
<1%
> 99
< 100
< 100 ppm
> 99
< 10
< 10 ppm
overall purity (%)
halide content (ppm)
water content
researcher had to prepare in-house, to commercially available materials. Over 200
individual ionic liquids in a variety of purity grades can be ordered from a range
of manufacturing companies including Merck KGaA (EMD Chemicals), BASF
and Cytec and chemical supply companies such as Fluka, Simga-Aldrich, VWR
and Kanto Chemical Co. Although the technical grades of ionic liquids which are
available from most companies are unsuitable for electrochemical applications,
a number of suppliers do offer higher purity grades. For example, Merck KGaA
offers its products in three purity grades (Table 2.4): “for synthesis” grade, which
is suitable for most non-electrochemical applications; “high purity” grade, which
is suitable for many catalytic and electrochemical applications; and “ultra pure”
grade, which was specified with electrochemical applications in mind.
As ionic liquids are adopted for industrial applications, questions are arising
concerning their toxicity, their impact on the environment and the registration of
these new chemicals with regulatory bodies. Several academic groups have led the
way in exploring the relationship between ionic liquids structures and their (eco)
toxicological effects [94–96] and their biodegradability [96, 97]. Information from
these studies will be useful in the design of more benign ionic liquids. Although
thorough studies of ionic liquids are rare, the fact that 1-ethyl-3-methylimidazolium
ethylsulfate has been classified as a non-toxic material gives a good indication that
other benign ionic liquids are highly likely. Although this ionic liquid is unlikely to
find use in electrochemical processes, due to the instability of the anion, there is
so far no reason to doubt that one or more environmentally friendly ionic liquids
exist with the physicochemical properties suitable to make a key component in an
electroplating system.
2.3
Eutectic-based Ionic Liquids
The melting point of two component mixtures is dependent upon the interaction
between the components. For non-interacting components the freezing point can
vary linearly with mole fraction whereas large negative deviations can occur when
the components interact strongly with each other. This is shown schematically
in Figure 2.4. The composition at which the minimum freezing point occurs is
31
32 2 Synthesis of Ionic Liquids
Fig. 2.4 Schematic representation of a eutectic point on a two-component phase diagram.
known as the eutectic point and this is also the temperature where the phases
simultaneously crystallise from molten solution. The word eutectic comes from
eutektos which is Greek for easily melted.
Eutectic mixtures have been used extensively for applications of molten salts to
reduce the operating temperature and this is where the significant area of ionic liquids developed from i.e. the quest to find aluminum-based salt mixtures. While the
development of aluminum-containing ionic liquids is technologically very important for the field of metal deposition it is clear that there are many other issues that
also need to be addressed and hence methods need to be developed to incorporate
a wide range of other metals into ionic liquid formulations.
While the first aluminum-based ionic liquids were reported in the 1950s [94],
it was not until the late 1990s that other metal salts were used to form ionic
liquids. Work by Abbott et al. [95, 96] and Sun et al. [97, 98] showed that eutectic
mixtures of zinc halides and quaternary ammonium halides also have melting
points close to ambient conditions. This has been further extended to a wide range
of other salts and organic compounds that form eutectic mixtures with quaternary
ammonium salts. This area has received comparatively little attention compared
with the chloroaluminate and discrete anions but the principle is simple in that the
complexing agent just needs to be able to complex the simple anion to effectively
delocalize the charge and decrease the interaction with the cation. This is shown
schematically in Figure 2.5.
The systems so far described can be expressed in terms of the general formula
Cat+ X− · z Y, where Cat+ is in principle any ammonium, phosphonium or sulfonium cation, X is generally a halide anion (usually Cl− ). They are based on equilibria
set up between X− and a Lewis or Brønsted acid Y, z refers to the number of Y
molecules which complex X− . The ionic liquids described can be subdivided into
three types depending on the nature of the complexing agent used.
2.3 Eutectic-based Ionic Liquids
Fig. 2.5 Schematic representation of the complexation occurring when
a Lewis acid or a Brønsted acid interacts with a quaternary ammonium salt.
Eutectic Type 1 Y = MClx, M = Zn [95–98], Sn [96], Fe [96], Al [94], Ga [99], In [100]
Eutectic Type 2 Y = MClx·yH2 O, M = Cr [101], Co, Cu, Ni, Fe
Eutectic Type 3 Y = RZ, Z = CONH2 [102], COOH [103], OH [104]
To date the only Cat+ species studied have been based on pyridinium, imidazolium
and quaternary ammonium moieties. In general, as with the chloroaluminate and
discrete anion systems, the imidazolium-based liquids have the lowest freezing
points and viscosities and higher conductivities. The depression of freezing point
is related to the strength of interaction between the anion and complexing agent
although this has not really been quantified as yet due primarily to a lack of thermodynamic data for the individual components.
One of the key advantages of these types of ionic liquids is the ease of manufacture. The liquid formation is generally mildly endothermic and requires simply
mixing the two components with gentle heating. Another key advantage is that
they are water insensitive which is very important for practical electroplating systems. As will be shown in Chapter 6.3, the electrochemistry of metals is relatively
unaffected by relatively large concentrations of water either naturally absorbed or
deliberately added to the ionic liquids. The final key advantage of eutectic-based
systems is that because they are simple mixtures of known chemicals they do not
have to be registered as new entities as they revert to their constituent components
upon excessive dilution in water.
2.3.1
Type I Eutectics
An extensive range of metal salts [96] have been studied but the only ones which
produce ionic liquids (i.e. liquid below 100 ◦ C) with pyridinium, imidazolium and
quaternary ammonium halides are FeCl3 , ZnCl2 , SnCl2 , CuCl [105], InCl3 [100]
and AuCl3 [106, 107].
It is thought that the ability of a metal salt to form a low melting point ionic liquid
will be related to its own melting point. The reason for this is apparent from Figure
2.4. Hence aluminum chloride (mp 190 ◦ C) has been shown to be useful with a wide
range of quaternary ammonium salts. Table 2.5 shows that relatively low melting
points are also possessed by ZnCl2 , SnCl2 and FeCl3 . Metal salts that do not form
ionic liquids with ammonium salts tend to have high melting points resulting from
large lattice energies. It is generally true that the metals have linear or tetrahedral
33
34 2 Synthesis of Ionic Liquids
Table 2.5 Freezing temperature data for a variety of metal salts and
amides when mixed with choline chloride in 2:1 ratio.
Type I
Type II
Type III
a
ZnCl2
SnCl2
FeCl3
CrCl3 .6H2 O
MgCl2 .6H2 O
CoCl2 .6H2 O
LaCl3 .6H2 O
CuCl2 .2H2 O
Urea
1, Methyl urea,
1,3 Dimethyl urea
1,1 Dimethyl urea
Thiourea
Acetamide
Benzamide
T f /◦ C
T f * /◦ C
T f /◦ C
24
37
65
4
10
16
6
48
12
29
70
149
69
51
92
283
246
306
83
117a
86
91
100a
134
93
102
180
175
80
129
259
209
241
79
107
70
85
52
122
64
32
31
106
29
37
Denotes decomposition temperatures.
geometries and tend to form predominantly univalent anionic complexes. In the
cases of FeCl3 , ZnCl2 and SnCl2 a variety of complex anions are known to form
whereas for CuCl, InCl3 , AuCl3 and TeCl4 only monometallate anions are known
to form i.e. CuCl2 − , InCl4 − , AuCl4 − and TeCl6 2− .
For the zinc chloride: choline chloride mixtures the eutectic is observed at a 2:1
composition, whereas for the tin chloride: choline chloride mixtures it is observed
at 2.5:1. This is presumably because SnCl2 is less Lewis acidic than ZnCl2 and
hence more SnCl2 is require to push the equilibrium for the reaction SnCl2 +
SnCl3 − ⇆ Sn2 Cl5 − to the optimum Sn2 Cl5 − composition.
The ZnCl2 system has probably been studied in the most detail. Fast atom
bombardment mass spectrometry (FAB MS) has been used to identify the species
present. It was found that ZnCl3 − , Zn2 Cl5 − and Zn3 Cl7 − species are all present
in the liquids. The relative proportions of anionic species depend on the ionic
liquid composition. Lecocq et al. [108] used electrospray ionization to look at the
various species present and found that in Lewis basic liquids x(ZnCl2 ) < 0.5 ZnCl3 −
whereas the di- and tri-metallate species were more prevalent in Lewis acidic liquids.
Presumably, small changes in concentration of each of the complex anions
change the ion–ion interactions markedly and this in turn changes the freezing
point. For example, ZnCl3 − ions are smaller and have a higher charge density than
Zn2 Cl5 − anions so are likely to have stronger electrostatic interactions with the
cation thus increasing the freezing point. Hence, as the mole fraction of ZnCl2
increases from 50% the amount of Zn2 Cl5 − relative to ZnCl3 − should increase
and the freezing point decreases. Above a mole fraction of 66% ZnCl2 the freezing
point increases again.
2.3 Eutectic-based Ionic Liquids
The composition of the various chlorozincate anions in the Lewis acid ionic
liquids was determined using potentiometry in an analogous manner to that used
by Heerman and D’Olislager [109] who measured the potential of the cell
Al|BuPyCl, AlCl3 (ref )||AlCl3 (x), BuPyCl(1 − x)|Al
They found that the equilibrium constant for the process 2Al2 Cl7 − DAlCl4 − +
Al3 Cl10 − was 2.93 × 10−3 at 60 ◦ C i.e. Al2 Cl7 − is the most abundant species in
solution.
The cell
Zn|ZnCl2 (0.667) ChCl (0.333)||ZnCl2 (x) ChCl (1 − x)|Zn
was used to determine an equilibrium constant of 2.0 × 10−5 for the reaction
→ ZnCl3 − + Zn3 Cl7 −
2Zn2 Cl5 −←
The value is lower than that for the analogous aluminum case, which would be
expected because of the difference in Lewis acidity. Hence the main species at the
eutectic composition was found to be Zn2 Cl5 − [96].
Liu et al. [110] studied the crystal structures of chlorozincate–choline chloride
complexes and identified that in an equimolar ratio the liquid is made up of two
species. It was shown that two types of crystal could be grown from the super-cooled
liquid, one rod-like and the other sheet-like, and these were thought to be due to
the Zn2 Cl5 − and ZnCl3 − salts, respectively.
13
C and 35 Cl NMR spectra of [BMIM]Cl and [BMIM]ZnCl3 showed that at 25 ◦ C
there is a significant difference between the two systems whereas at 110 ◦ C the
systems are similar; this showed that the zinc-containing liquid is highly associated
at lower temperatures. A more dissociated structure is favored at high temperatures.
This is significant for metal deposition studies as the coordination geometry will
affect the way in which the metal is reduced.
The phase behavior of ionic liquids will depend upon the potential energy between the ions but this is difficult to model for a eutectic-based ionic liquid because
of the complex nature of the anion and the non-centrosymmetric charge distribution on the cation. However, if the difference in freezing point between that of the
quaternary ammonium salt and that of the complex with the metal salt is considered
then the issue becomes significantly easier. The change in interionic potential energy, E p , and the resulting change in freezing temperature, T f will be related to the
expansion of the ionic lattice resulting from the formation of a complex anion. Since
Ep =
q1q2
4πεo r
(2.1)
where q is the charge on the ions, ε o is the vacuum permittivity and r is the
separation between the two charges. Therefore
E p ∝
rc − rs
rs
(2.2)
35
36 2 Synthesis of Ionic Liquids
where r c and r s are the charge separation in the complex mixture and simple
quaternary ammonium halide salt, respectively (assuming that at the eutectic
composition the complex anion is predominantly of the form M2 X5 − ). The
depression of freezing point, T f , is taken as the difference between the measured
freezing point at the eutectic composition, T f and the freezing temperature of the
pure quaternary ammonium halide. It was shown [96] that by plotting the freezing
point depression as a function of the normalized change in charge separation
(Eq. (2.2)) produces a good correlation for ZnCl2 -based ionic liquids. This is a
significant result as it allows phase behavior to be predicted from simple ionic size
considerations and it shows that symmetry has a negligible effect on the depression
of freezing point, but it does change the absolute freezing point. Angell [111]
recently used a similar approach to show that the glass transition temperature of
a range of ionic liquids is related to the molar volume of the ions.
The effect of the quaternary ammonium cations is quite complex because the
smaller cations depress the freezing point more because the halide salts of the
smaller cations also have a higher freezing point; the net result is that all of the
eutectic mixtures will have reasonably similar freezing points. Hence the cation is
observed to have little effect on the absolute freezing point of the eutectic-based
ionic liquids.
Lecocq et al. [108] studied ionic liquids formed between zinc chloride and
1-butyl-2,3-dimethylimidazolium chloride [BMMIM]Cl with the amount of ZnCl2
between 0 and 0.75 mol%. Analysis using NMR, and mass spectrometry showed
Cl− and [ZnCl3 ]− in Lewis basic liquids and [ZnCl3 ]− and [Zn3 Cl7 ]− in Lewis acidic
liquids. Infrared spectra with pyridine were used to quantify the Lewis acidity and
high temperature (110 ◦ C) NMR experiments showed that the structure varies with
time from [BMMIM][ZnCl3 ] to [BMMIM. . .Cl. . .ZnCl2 ].
The iron-based systems have two eutectic points in an analogous manner to
the chloroaluminate systems. The eutectic points occur at 33 and 67 mol% FeCl3
[96]. We were only able to identify the species FeCl4 − by FAB MS in the choline
chloride–FeCl3 system but this could be because other species are too weak to
be observed by this technique. Other groups have prepared iron-containing liquids with FeCl2 and FeCl3 . Sitze et al. [112] found that [BMIM]Cl formed liquids
with FeCl2 in the molar ratio 0.3 FeCl2 : 1 [BMIM]Cl whereas the ferric chloride
formed in the molar ratio 0.53 to 1.7. Raman scattering and ab initio calculations
showed that FeCl4 2− was the prevalent anion present with ferrous chloride, whereas
FeCl4 − and Fe2 Cl7 − were present in the ferric chloride system. The relative concentrations were dependent upon the Lewis acidity in an analogous manner to the
zinc and aluminum systems. Zhang et al. [113] also studied FeCl3 and 1-methyl-3butylimidazolium chloride ([BMIM]Cl) with a molar ratio of 1:1 and characterized
the physical properties of the liquid. Hayashi et al. [114] also studied the [BMIM]
FeCl4 system and found that the liquid is ferromagnetic.
The ability to vary the composition of Lewis or Brønsted acid adds an additional
dimension to the tuneability of the eutectic-based ionic liquids. It has been shown
that the Lewis acidity of the liquid affects not only the physical properties of the
liquids but also the electrochemical behavior. Type I ionic liquids are also clearly
2.3 Eutectic-based Ionic Liquids
Scheme 2.6
useful for electroplating if the metal of interest falls in the category defined above
as the metal ion concentration can be as high as 10 mol dm−3 .
Eutectic mixtures of imidazolium chloride with GaCl3 and InCl3 have also been
reported [99, 100]. These metals will be of limited interest for electrodeposition
although some studies have been made on the deposition of semiconductors [115,
116]. Other metal halides that have been used include AuCl3 , NiCl2 and CoCl2
[106, 107]. These tend to have higher melting points than other metal salts for the
reasons explained above. They have been used for synthetic applications and while,
in principle, they could be used for electrodeposition there are better alternatives
that would be more suitable.
Seddon et al. [117] have produced ionic liquids of the type [EMIM]2 [UCl6 ] from
the reaction of UCl4 with [EMIM]Cl. The uranium was isolated using electrochemical reduction and it was proposed that this was a potential method for recycling
spent nuclear fuel. Hagiwara [118] produced ionic liquids containing niobium and
tantalum from the reaction of [EMIM]F.2.3HF with TaF5 and NbF5 to produce
[EMIM] TaF6 and [EMIM] NbF6 .
While the majority of studies in this area have concentrated on halide salts some
intriguing work has been carried out using metal oxides. Noguera et al. [119] showed
that CrO3 and Na2 MoO4 could be incorporated into ionic liquids. Scheme 2.6 shows
the synthesis of two ionic liquids and although the electrodeposition of the metal
was not reported it could, in principle, be used for such applications. These liquids
have been shown to be good oxidants for organic reactions. A number of other
strategies have been published for the production of metal-containing ionic liquids
and while most of these are very exotic and have been used for catalysis some of the
generic methodologies may eventually find application in electrodeposition. This
area has recently been reviewed by Lin and Vasam [120].
The conductivities of Type I ionic liquids based on anhydrous zinc and iron
salts tend to be lower than those of the corresponding aluminum ionic liquids.
This is due largely to the higher viscosity of the former, primarily because of the
large size of the ions and the availability of suitably sized holes in the ionic liquids
for the ions to move into. This has been quantified by the application of hole
theory as is explained in Section 2.3.4. In general imidazolium-based liquids have
lower viscosities and higher conductivities than the corresponding pyridinium or
quaternary ammonium eutectics formed under the same conditions.
37
38 2 Synthesis of Ionic Liquids
2.3.2
Type II Eutectics
These were developed in an endeavor to expand the range of metals that could be
incorporated into an ionic liquid. The presence of waters of hydration decreases
the melting point of metal salts because it decreases the lattice energy. Hence,
as Figure 2.4 shows, hydrated salts should be more likely to form mixtures with
quaternary ammonium salts that are liquid at ambient temperature than anhydrous
salts. Table 2.5 shows a list of some of the metal salts that have been made into
ionic liquids with choline chloride and the freezing point of a 1ChCl:2metal salt
mixture.
Electrospray MS of the eutectic mixture showed two primary signals M+ 104
[Choline]+ and M− 192/194* /196 [CrCl4 ]− (the waters of hydration are bound too
weakly to be observed and Cr(H2 O)3 Cl3 is neutral and therefore not detected) [101].
The UV–vis spectrum of the eutectic mixture showed the presence of predominantly
Cr(H2 O)3 Cl3 with some evidence of [CrCl4 ·2H2 O]− . It was concluded that the main
charge carrying species were [Choline]+ and [Cl·3H2 O]− . This would account for
the high conductivities of these liquids compared to the anhydrous salt mixtures.
The addition of LiCl to the ionic liquid was found to have only a small effect upon
the conductivity of the liquid, but it did affect the speciation [121], producing more
of the [CrCl4 ·2H2 O]− . It was anticipated that the small Li+ ion would have a high
mobility in the liquid but the conductivity is less than expected, suggesting that the
ion must be strongly solvated or highly associated with the anion.
Unlike the anhydrous metal salts, these mixtures are very sensitive to temperature
fluctuations. At ambient temperatures they are extremely hygroscopic and rapidly
absorb up to 10 wt% water from the atmosphere. Above 70 ◦ C the liquids lose water
and this is characterized by a change in color of the chromium-based liquid from
dark green to purple. At about 50 to 60 ◦ C the water concentration in the liquid
remains constant and can be used in an open atmosphere without significant
alteration in the liquid composition. Thermogravimetry shows that the waters of
hydration are released in two steps; the first starts at about 85 ◦ C, which equates to
approximately 3 waters, and the second at about 180 ◦ C, corresponding to the other
3 water molecules [101].
To date the only concerted study has been carried out using chromium chloride,
but it has been reported that a number of other metals form this type of eutectic
mixture and Table 2.5 lists just some of the metal salts that have been studied,
together with their freezing points in eutectic mixtures with choline chloride. Potentially there are some very interesting systems but to date only Cr and Co have
been deposited from these liquids. The deposition of metals such as Al and Ca is
not possible due to the limited potential window of these liquids.
2.3.3
Type III Eutectics
It has recently been shown that the principle of creating an ionic fluid by complexing
a halide salt can be applied to mixtures of quaternary ammonium salts with a
2.3 Eutectic-based Ionic Liquids
range of amides [102, 103]. The charge delocalization is achieved through hydrogen
bonding between the halide anion with an amide, carboxylic acid or alcohol moiety.
Table 2.5 lists the freezing points of a number of hydrogen bond donor (HBD)
mixtures with choline chloride. These liquids have interesting solvent properties
that are similar to other eutectic ionic liquids and a wide variety of solutes were
found to exhibit high solubilities [102, 103]. The depression of freezing point (with
respect to an ideal mixture of the two components) for a number of these eutectic
systems is extremely large e.g. the oxalic acid–choline chloride system, is 212 ◦ C and
the choline chloride–urea system is 178 ◦ C [103]. The freezing point depressions
are not as large as the choline chloride–zinc chloride system (272 ◦ C) [97] due to
the covalent bonds formed in the metal chloride case. To differentiate these liquids
from ionic liquids the term Deep Eutectic Solvents (DES) has been adopted. Unlike
the room-temperature ionic liquids, these eutectic mixtures are easy to prepare in
a pure state. They are non-reactive with water, many are biodegradable and the
toxicological properties of the components are well characterized.
It is thought that the chloride complexes with 2 HBDs and this accounts for the
varying eutectic composition. For monofunctional HBDs e.g. urea, phenylpropionic acid, the eutectic point occurs at 67 mol% HBD, for difunctional HBDs, e.g.
oxalic acid and malonic acid, the eutectic point occurs at 50 mol% HBD and for
citric acid the eutectic occurs at 33 mol% HBD. The tris-carboxylic acids exhibit
the rheology of gels and presumably have extensive bridging of the acids between
neighboring chloride ions. The existence of hydrogen bonding in ChCl/urea eutectic mixtures can be observed using NMR spectroscopy [102]. Heteronuclear Overhauser effect spectroscopy (HOESY) of HOCH2 CH2 N + (CH3 )3 F · 2(NH2 )2 CO
shows intense cross-correlation between the fluoride ion and the NH2 protons on
the urea molecule. Some anion complexes have been identified using FAB MS
and it is evident that the HBD is sufficiently strongly coordinated to the chloride
anion to be detected by this technique. In a 1 choline chloride: 2 urea mixture the
presence of Cl− with two ureas (M− = 155) and Cl− with one urea (M− = 95) was
observed.
As with the chlorometallate eutectics a model for the effect of HBD on the freezing
point depression of the mixture would be beneficial for the design of new liquids.
No correlations were observed between the freezing point of the mixtures and the
enthalpy of formation or fusion of the pure acids but Table 2.5 shows qualitatively
that the larger depressions of freezing point occur with the lower molecular weight
HBDs.
The freezing point of the HBD–salt mixtures will be dependent upon the lattice
energies of the salt and HBD and how these are counteracted by the anion–HBD
interaction and the entropy changes arising from forming a liquid. For a given
quaternary ammonium salt, the lattice energy of the HBD will be related to the
anion–HBD interaction and hence, to a first approximation, the depression of
freezing point will be a measure of the entropy change. It has been shown [103]
that the depression of freezing point correlates well with the mass fraction of HBD
in the mixture.
The lowest viscosities and highest conductivities are obtained with diol-based
HBDs. It is thought that the comparatively weak interactions between the alcohol
39
40 2 Synthesis of Ionic Liquids
Table 2.6 Viscosity and conductivity of a variety of ionic liquids at
298 K.
Cation
Anion
EMIM
EMIM
BMIM
BuMePy
choline
choline
choline
choline
choline
acetylcholine
choline
choline
BF4 −
N(CF3 SO2 )2
BF4 −
N(CF3 SO2 )2
Zn2 Cl5 −
CrCl4 ·6H2 O
CoCl3 ·6H2 O
Cl·2urea
Cl·2propanediol
Cl·2propanediol
Cl·malonic acid
Cl·2ethylene glycol
j/mS cm−1
14
8.4
3.5
2.2
0.02
0.37
1.7
0.75
2.2
0.51
0.36
7.6
g/cP
32
28
180
85
76 000
2346
392
632
89
117
3340
36
and the chloride mean that some ‘free’ glycol is able to move, decreasing the
viscosity of the liquid. The glycol-based liquids tend also to have comparatively large
potential windows. Hence the Abbott group has carried out a number of studies
using ethylene glycol with choline chloride. This mixture has been shown to be
useful for the deposition of zinc and zinc alloys [122] as well as the electropolishing
of stainless steel [104] (see Chapter 11.1). The liquid is inexpensive, non-toxic,
non-viscous and highly conducting compared to other ionic liquids.
2.3.4
Modelling Viscosity and Conductivity
One of the main differences between ionic liquids and aqueous solutions is the comparatively high viscosity of the former. Table 2.6 shows that viscosities are typically
in the range 10–500 cP (0.01–0.50 Pa s) and this affects the diffusion coefficients of
species in solution.
Most new liquids have viscosity that varies as a function of temperature and the
majority vary in an Arrhenius manner with temperature [123]:
ln η = ln η0 +
Eη
RT
(2.3)
where Eη is the activation energy viscous flow and η0 is a constant. Other researchers
have found that the viscosity obeys a Vogel–Tamman–Fulcher relationship [124]. A
comprehensive study of viscosity is that of VanderNoot [124] and there are several
collections of viscosity data in recent reviews [125–127].
We have fitted the viscosity of ionic liquids using hole theory [123]. The theory
was developed for molten salts but has been shown to be very useful for ionic
liquids. It was shown that the value of E η is related to the size of the ions and
the size of the voids present in the liquid [103]. The viscosity of ionic liquids is
2.3 Eutectic-based Ionic Liquids
several orders of magnitude higher than that of high-temperature molten salts due
partially to the difference in size of the ions, but also to the increased void volume
in the latter. It has been shown [128] that hole theory can be applied to both ionic
and molecular fluids to account for viscosity. The viscosity of a fluid, η, can be
modeled by assuming it behaves like an ideal gas, but its motion is restricted by the
availability of sites for the ions/molecules to move into. Hence it was shown that
η=
mc̄/2.12σ
P(r > R)
(2.4)
where m is the molecular mass (for ionic fluids this was taken as the geometric
mean), c̄ is the average speed of the molecule (=(8kT/πm)1/2 ) and σ is the collision
diameter of the molecule (4πR2 ). The probability of finding a hole of radius, r,
greater than the radius of the solvent molecule, R, in a given liquid, (P(r > R)) is
given by integration of the following expression [123]:
Pdr =
16 7/2 6 −ar 2
dr
√ a r e
15 π
(2.5)
where a = 4πγ /kT and γ is the surface tension. The good correlation obtained
between the calculated and measured viscosities shows that it is valid to think of
the viscosity of fluids as being limited by the availability of holes. It is evident
from Eqs. (2.4) and (2.5) that decreased viscosity can be obtained by decreasing the
surface tension of the liquid, i.e. increasing the free volume, or by decreasing the
ionic radius.
Hence the ionic liquids with the lowest viscosity tend to have highly fluorinated
anions as these shield the charge density and result in low surface tensions. The
cation also affects the viscosity of ionic liquids. For imidazolium cations, the viscosity initially decreases as the length of the R group increases, as the ion–ion
interactions decrease and hence the surface tension decreases. However, as the
alkyl group increases in size its mobility will decrease due to a lack of suitably sized
voids for the cations to move into. This can be seen in the data presented by Tokuda
et al. who showed a minimum in viscosity for ethyl methyl imidazolium salts [129].
The conductivity of ionic liquids can be modeled in the same manner as the viscosity, i.e. despite the high ionic strength of the liquid, ionic migration is limited by
the availability of suitably sized voids [130]. Since the fraction of suitably sized holes
in ambient temperature ionic liquids is effectively at infinite dilution, migration
should be described by a combination of the Stokes–Einstein and Nernst–Einstein
equations. This is explained in greater detail in Chapter 11.3 on process scale-up
but it is sufficient to say that an expression can be derived for the conductivity, κ
κ=
z2 F e
6πη
1
1
+
R+
R
ρ
Mw
(2.6)
where ρ is the density and Mw is the molar mass of the ionic fluid. Hence the molar
conductivity ( = κ/c) is, in effect, independent of the number of charge carriers
and this is the reason why the empirical Walden rule [123, 126] ( η = constant)
41
42 2 Synthesis of Ionic Liquids
is applicable to ionic liquids. The Walden rule is normally only valid for ions at
infinite dilution where ion–ion interactions can be ignored, which is clearly not the
case in ionic liquids.
It is apparent from the above discussion that ionic mobility is controlled by the
free volume of a liquid and the size of the ions. The size of the voids in the liquid and
their effect on liquid density can be changed by decreasing the ion–ion interactions.
This will manifest itself by a decrease in surface tension and, in general, the liquids
with lower surface tensions are more fluid and have higher conductivities. This is
the reason why ionic liquids with discrete, highly fluorinated anions such as PF6
and (F3 CSO2 )2 N have become popular.
It has recently been shown that the same principle can be applied to deep eutectic
solvents by using small quaternary ammonium cations such as ethylammonium
and fluorinated hydrogen bond donors such as trifluoroacetamide. However, there
is only a limited benefit that can be achieved using this approach as the physical
parameters cannot be varied totally independently of one another. For example
there will be an optimum ion size; too small and the lattice energy will increase the
surface tension, too large and the ionic mobility will be impeded.
2.3.5
Conclusions
This chapter shows that eutectic-based ionic liquids can be made in a variety of
ways. The above description of liquids falling into three types is by no means exclusive and will certainly expand over the coming years. While there are disadvantages
in terms of viscosity and conductivity these are outweighed for many metal deposition processes by issues such as cost, ease of manufacture, decreased toxicity
and insensitivity to moisture. The high viscosity of some of these liquids could be
ameliorated in many circumstances by the addition of inert diluents.
The physical principles underlying eutectic-based ionic liquids are now relatively
well understood, however, the liquids described above have tended to be less academically fashionable and have received comparatively little attention. Concerted
effort with these types of liquids could lead to optimization of their properties such
that they would be suitable for commercial deposition processes.
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47
3
Physical Properties of Ionic Liquids for
Electrochemical Applications
Hiroyuki Ohno
3.1
Introduction
In spite of the explosion in studies on ionic liquids (ILs), there is only a small
number of studies of their basic characteristics. There are limitless possibilities for
the design of ILs by changing their component ion structures. However, the chance
of success is not very great without accurate information on the structure–properties
relationship. Physico-chemical property data for ILs are therefore very important for
the present and future of the field of ILs. In this chapter, some basic properties of airstable ILs have been summarized. Some are not directly related to electrochemistry
but are very important and useful for a wide range of science and technology related
to ILs.
3.2
Thermal Properties
ILs are defined as organic salts having a melting point (T m ) below 100 ◦ C [1–5].
In order to use these ILs as non-volatile electrolyte solutions, it is necessary to
maintain the liquid phase over a wide temperature range. Consequently, T m and
the thermal degradation temperature (T d ) of ILs are important properties for ILs
as electrochemical media. In this section, the thermal properties of ILs, especially
of imidazolium salts, are summarized. The difference between ILs and general
electrolyte solutions based on molecular solvents is clarified. Recent results on the
correlation between the structure and properties of ILs will also be mentioned.
3.2.1
Melting Point
ILs are differentiated from typical inorganic salts by their low T m . Typical inorganic
salts have a high T m , around 1000 ◦ C reflecting high lattice energies, i.e., the high
T m is attributable to a strong electrostatic attractive force between the ions. Since
ILs are organic compounds, van der Waals interaction, hydrogen bonding, and π–π
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
48 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Table 3.1 Melting point (◦ C) of several salts
Cation
Na+ (1.02 Å)
K+ (1.38 Å)
Cs+ (1.67 Å)
[N1111 ]+
[N2222 ]+ (3.35 Å)
[emim]+ (3.04 Å)
Anion
Cl− (1.81 Å)
Br− (1.96 Å)
I− (2.20 Å)
BF4 − (2.29 Å)
808
772
645
420
>300
87
747
734
636
230a
285
77
662
685
621
230[a]
300
78
384
530
72
11
(CF3 SO2 )2 N−
(3.25 Å)
104
−15
Ion radius is given in parentheses. a Decomposition temperature, [N1111 ]+ , tetramethylammonium
cation; [N2222 ]+ , tetraethylammonium cation; [EMIM]+ , 1-ethyl-3-methylimidazolium cation.
interaction are additionally present among the component ions. These interactions
affect the T m of ILs. Accordingly, structural design of component ions to weaken
the electrostatic interaction and other interactions is directly effective in lowering
the T m of the salts, as discussed in detail below. However, it is still difficult to
predict the T m of any given salt from its structure.
3.2.1.1 Effect of Ion Radius
When ions have equivalent charges, the electrostatic interaction decreases with
increasing ion radius because the surface charge density decreases with increasing
ion radius and the separation between the ions also increases. The electrostatic
interaction of larger ions is therefore weaker, and accordingly the salts show lower
T m . Table 3.1 shows the T m of typical salts and their ion radii [6–9]. In general,
organic salts have lower T m than inorganic salts because of their larger ion size, as
shown in Table 3.1. However, the tetraethylammonium cation ([N2222 ]+ ) is larger
than 1-ethyl-3-methylimidazolium cation ([EMIM]+ ), yet the [N2222 ]+ salts show
higher T m than [EMIM]+ salts. This is a typical example that illustrates the case
where T m depends not only on the electrostatic interaction. Since the delocalization
of charge also contributes to lowering the electrostatic interaction, the existence of
π electron orbitals is important in lowering the T m .
3.2.1.2 Effect of Cation Structure on the Melting Point
Onium Cations. Major families of ILs are composed of quaternary onium cations
such as imidazolium, pyridinium, ammonium, phosphonium, sulfonium cations
and so on. As described above, the fact that most ILs are composed of organic
cations is attributed to weaker electrostatic interaction among component ions.
There have been several reports on the effect of cation structure on the T m of ILs.
The relationship between T m and the basic structure of onium cations is important
in developing a protocol to prepare low melting ILs.
The cation structure and T m of bis(trifluoromethanesulfonyl)imide type ILs are
shown in Table 3.2. These cations, having similar length of alkyl chain, are chosen
3.2 Thermal Properties
Table 3.2 Melting point (T m ) of a series of
bis(trifluoromethanesulfonyl)imide type ionic liquids
Cation
T m /◦ C
Ref.
−15
7
−4
10
20
11
86
12
−18
12
44
13
29.2
14
28.7
14
104
15
90
15
for comparison to exclude the effect of the alkyl chain on their T m s. Among the salts
in Table 3.2, the imidazolium and pyridinium cations are aromatic and the others
are aliphatic. These aromatic salts show relatively low T m because of delocalized
positive charge, as mentioned previously.
Effect of Side Chain Length. The side chain bound to the cation also affects the T m
due to flexibility and excluded volume effects. To discuss the relation between side
chain structure and T m , the onium cation structure was fixed. The relation between
the side chain structure and the thermal properties of imidazolium salts has already
been reported by Seddon et al. [16, 17]. The effect of alkyl chain length of 1-alkyl3-methylimidazolium tetrafluoroborate on the phase transition temperatures is
49
50 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Fig. 3.1 Phase diagram for 1-alkyl-3-methylimidazolium tetrafluoroborate showing the melting (•), glass (◦) and clearing () transitions
measured by differential scanning calorimetry. Data from Ref. [16].
shown in Figure 3.1 [16]. With a carbon number from 2 to 9 on the imidazolium
ring, the salts are liquid at room temperature. On the other hand, when the carbon
number of the alkyl chain of the imidazolium ring was 0, 1, or larger than 9, the
salts showed a clear T m . A liquid crystalline phase appeared for imidazolium salts
with a carbon chain longer than the dodecyl group. This liquid crystalline phase
arises due to the orientational effect of the long alkyl chains. A similar tendency
has also been observed in the case of PF6 imidazolium salts [17].
Symmetry is another factor to affect T m . The salts with symmetric ions generally show higher T m than those with asymmetric ones. For example, 1,3dimethylimidazolium tetrafluoroborate showed higher T m than 1-methylimidazolium or 1-ethyl-3-methylimidazolium salts, as shown in Figure 3.1. In the
case of tetraalkylammonium salts, their T m also increased with increasing symmetry of the cation structure [18]. This tendency is understood to relate to the
structural effect on crystallinity [19], i.e., highly symmetric ions are more efficiently
packed into the crystalline structure than unsymmetric ones. Other kinds of chain
structures such as polyether [20], perfluorocarbon [21], etc. [22] are obviously also
effective in influencing thermal properties.
3.2.1.3 Anion Species
There are many choices of anion species for IL synthesis. In particular, halogencontaining anions such as BF4 − , PF6 − , and TFSI− are often used to prepare ILs.
Room-temperature ILs are obtained with ions having weaker electrostatic interaction originating from negative charge delocalization and stabilization by the
electron-withdrawing effect of halogen atoms. Non-halogenated anion-containing
ILs with low T m have also been prepared after suitable structural design to
lower the anionic charge density [23]. Thermal properties of such 1-ethyl-3methylimidazolium and 1-butyl-3-methylimidazolium salts are summarized in Table 3.3. Larger anions generally form ILs with lower T m . Lowering the surface
3.2 Thermal Properties
Table 3.3 Thermal properties of imidazolium-type ionic liquids com-
posed of various anions
Salt
Cation
Anion
[EMIM]+
[BMIM]+
Tm / ◦C
Tg / ◦C
Td / ◦C
[Cl]−
[Br]−
[I]−
[BF4 ]−
[PF6 ]−
[NO3 ]−
[CH3 COO]−
[CF3 COO]−
[CH3 SO3 ]−
[CF3 SO3 ]−
[(CF3 SO2 )2 N]−
[(C2 F5 SO2 )2 N]−
[(CN)2 N]−
[(CF3 SO2 )3 C]−
[(CN)3 C]−
8915
7915
7915
1115 , 1524
6215 , 5825
117 , 3826
4515
–146
3927
–96
–157
–115
–2123a)
3915
–1128
––
––
––
–8624
––
––
––
28515
––
–987
–9528
44010
4557
42315
––
45015
––
[Cl]−
[Br]−
[I]−
[BF4 ]−
[PF6 ]−
[NO3 ]−
[CH3 COO]−
[CF3 COO]−
[CH3 SO3 ]−
[CF3 SO3 ]−
[(CF3 SO2 )2 N]−
[(C2 F5 SO2 )2 N]−
[(CN)2 N]−
[(CF3 SO2 )3 C]−
[(CN)3 C]−
6510
––
––
–8129
–810 ,1027
––
–5031
––
–9731
–8031
25031
27332
26515
40331
34931
––
–7832
22035
17632
1610
–410
––
–8733
–628
––
–9031
–6531
–10423a)
30315
42034
––
––
––
15010
40932
43931
40232
30031
41331
[EMIM]+ , 1-ethyl-3-methylimidazolium cation; [BMIM]+ , 1-butyl-3-methylimiadzolium cation.
Reference numbers are shown as superscripts to the data.
charge densities of the anion also lowers the T m of the ILs. In spite of the fact that
most anions used are symmetric, there are a few approaches where T m has been
lowered by using asymmetric anions [36].
3.2.2
Glass Transition Temperature
The glass transition temperature (T g ) is generally understood to be the temperature
where segmental motion begins on heating from the quenched amorphous solid.
51
52 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Accordingly, the ionic conductivity and viscosity of many ILs are a function of T g .
The T g is thus not so important for electrodeposition from ILs, but is important
for ion conduction of and in the ILs (see ionic conductivity). In the case of ILs,
there are many examples that show both T g and T m . Detailed phase studies have
been reported elsewhere [16]. The relation between T g and T m has already been
discussed by Angell et al. [37] who observed that T g is almost equal to two-thirds of
T m in Kelvin.
3.2.3
Thermal Decomposition Temperature
ILs are thermally stable but certainly decompose at high temperature. The decomposition temperature (T d ) of general imidazolium type ILs is summarized in Table
3.3. The T d of ILs depends on the component ion structure, similarly to other thermal properties [30]. ILs having excellent thermal stability up to 400 ◦ C have been
reported [15,31,34]. However, this does not mean that these ILs can be used at any
temperature below T d , because most T d values are determined by using temperature sweeping thermogravimetric measurements. ILs gradually decompose even
below T d . It is therefore important to analyze the thermal behavior at constant
temperature.
Previously reported T d values of several imidazolium-type ILs are plotted against
side chain length (n) in Figure 3.2. T d is affected by water content, impurities, type
of flow gases, and vessel material for thermal gravimetric measurements [15]; hence
it should be noted here that the T d shown in Figure 3.2 is not the absolute value for
each IL. As shown in Figure 3.2, the ILs composed of BF4 − , PF6 − and TFSI− have T d
100 ◦ C higher than ILs composed of halogen anions such as Cl− or I− . Chan et al.
reported that an alkyl chain at the N-position of the imidazolium cation suffers
nucleophilic attack by the halide anion in the manner shown in Scheme 3.1 [38].
While remarkable differences in T d are observed by changing the anion species,
T d only depends slightly on the alkyl chain length on the imidazolium cation.
Fig. 3.2 Relation between thermal decomposition temperature (T d ) of
1-alkyl-3-methylimidazolium-type ILs with alkyl chain length (n). Anion
species are Cl− : (), I− : (•), BF4 − : (), PF6 − : () and TFSI− : (♦).
3.2 Thermal Properties
Scheme 3.1 Pyrolysis mechanism of ionic liquids containinig halide anions.
There is a report that the thermal stability of imidazolium salts is largely the same,
despite a great difference in alkyl chain length between the butyl and octadecyl
groups, suggesting simple extension of alkyl chain hardly affects T d [30].
3.2.4
Liquid Crystallinity and Solid–Solid Transitions
Some compounds show meso-phases between the solid and liquid phases. These
phases are classified into two kinds, namely liquid crystals in which the molecules
have orientational order and disorganized position in one or more dimensions, and
plastic crystal in which the molecules have organized positions and orientational
disorder. Although the component ions in ILs are largely disordered, the appearance
of liquid crystalline or plastic crystal phases could be the function of ion structures,
when component ions have a tendency towards orientational or positional ordering
by alignment of the ions and/or interaction among ions. Onium salt-type plastic
crystals have been reported by MacFarlane [12,39].
A series of 1-alkyl-3-methylimidazolium salts show liquid crystalline phases by
elongation of the alkyl side chain as shown in Figure 3.1. Seddon et al. reported
the liquid crystalline phase for pyridinium salts with longer alkyl chains, as well as
imidazolium salts [16, 17,40]. Introduction of a hydrophobic moiety into the anion is
also effective in yielding liquid crystalline properties in imidazolium salts. However,
it should be noted here that not all hydrophobic anions show liquid crystal phases.
The salts composed of an imidazolium cation having multiple methyl groups and
long chain alkylsulfonate anions showed liquid crystalline properties, depending on
the position and quantity of substituent groups on the imidazolium ring [41]. These
materials have been discussed as anisotropic ion conductors [42] and anomalous
reaction media [43]. It might be interesting to examine electrodeoposition in these
materials for super-fine surface designs such as parallel nanowires.
3.2.5
Thermal Conductivity
The thermal conductivity of ILs is an important property when using ILs for electrochemical synthesis or thermal storage. The thermal conductivity of ILs was
reported, together with heat capacity, by Wilkes et al., as summarized in Table 3.4
[44]. The heat capacities of ILs are 3 or 4 times larger than that of copper, but smaller
than that of water. The thermal conductivity of general ILs is lower than that of cop
R
per or water. Therminol VP-1, diphenyl oxide/biphenyl type thermal conductor,
is commercially available as a heat transport fluid. The thermal conductivity and
heat capacity of ILs are, in general, similar to those of VP-1.
53
54 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Table 3.4 Heat capacity and thermal conductivity of ionic liquids
R
Therminol VP-1
[EMIM][BF4 ]
[BMIM][BF4 ]
[p-DMIM][(CF3 SO2 )2 N]
H2 O at 30 ◦ C
H2 O at 100 ◦ C
copper
Heat capacity at 100 ◦ C
J g−1 K−1
Thermal conductivity at 25 ◦ C
W m−1 K−1
Ref.
1.78
1.28
1.66±0.08
1.20±0.05
4.18
4.22
0.385
0.127
0.200±0.003
0.186±0.001
0.131±0.001
0.615
0.679
398
44
44
44
44
45
45
46
[p-DIMIM][(CF3 SO2 )2 N]: 1-propyl-2,3-dimethylimidazolium bis(trifluoromethanesulfonyl)imide.
3.2.6
Vapor Pressure
The vapor pressure of ILs is substantially zero under ambient conditions. Therefore
ILs have been recognized as non-volatile liquids at normal pressures. However, it
is known experimentally that some ILs, synthesized by the neutralization of protic
acid with organic base, easily evaporate on heating. Angell et al. pointed out that
the acid–base equilibrium of those ILs becomes imbalanced on heating and then
generates the volatile acid and base [47]. Based on this, MacFarlane et al. recently
reported that the ILs prepared by neutralization, N-methylpyrrolidinium formate,
could be distilled 100% at 70 ◦ C under 0.9 mmHg [48]. Distillable ILs can, therefore,
be prepared by neutralization of volatile bases with volatile acids. On the other hand,
in general, ILs composed of quaternized onium cations and anions do not show
such an equilibrium. These ILs are generally decomposed on heating without
evaporation.
Recently, Seddon et al. reported that many known ILs including [EMIM][TFSI]
can be evaporated at 300 ◦ C under high vacuum (less than 0.1 mbar) [49]. Details
of the evaporation mechanism are not yet clear; a cluster ion model is proposed
because it is hardly conceivable that individual anions and cations are vaporized,
even under high vacuum.
3.3
Viscosity
Viscosity is an important property of ILs used as electrolyte solutions. There are
some basic studies on the viscosity of ILs in the literature [50, 51]. The reported
viscosities of imidazolium type ILs composed of commercially available anions are
relatively low, as summarized in Table 3.5. The reported viscosity values are not
always the same for any given IL owing to water content, impurities, synthetic
route, starting materials, and measurement method.
3.4 Density 55
Table 3.5 Viscosity of several liquids at room temperature (25 ◦ C ± 1)
[EMIM]+
[BMIM]+
[HMIM]+
[OMIM]+
[BF4 ]−
[PF6 ]−
[(CF3 SO2 )2 N]−
[(CF3 CF2 SO2 )2 N]−
[CF3 CO2 ]−
[CF3 SO3 ]−
[BF4 ]−
[PF6 ]−
[(CF3 SO2 )2 N]−
[(CF3 CF2 SO2 )2 N]−
[CF3 CO2 ]−
[CF3 SO3 ]−
[BF4 ]−
[PF6 ]−
[(CF3 SO2 )2 N]−
[BF4 ]−
[PF6 ]−
[(CF3 SO2 )2 N]−
[(CF3 CF2 SO2 )2 N]−
g / cP
Ref.
43
15 (80 ◦ C)
28
61
35
45 (30 ◦ C)
219
450
69
77
70
93
314 (20 ◦ C)
585
68
439
682
93
492
7
7
7
7
6
54
30
30
30
54
35
55
56
30
35
57
30
58
59
g / cP
water
methanol
acetic acid
acetone
acetonitrile
N,N-dimethylformamide
ethylene glycol
propylene glycol
glycerol
0.89
0.54
1.13
0.30
0.34
0.80
16.1
40.4
934
[HMIM]+ : 1-hexyl-3-methylimidazolium, [OMIM]+ : 1-octyl-3-methylimidazolium.
The viscosity of ILs is typically 10 to 100 times higher than that of water or organic
solvents [50–52] as a result of the strong electrostatic and other interaction forces.
The fluorohydrogenate type ILs reported by Hagiwara et al. have some of the lowest
viscosities known [53]. Low viscosity ILs are obviously preferred in electrolyte or
other reaction solvent applications, but it is quite difficult to design low viscosity ILs.
The imidazolium ILs tend to show decreasing viscosity in the following order of
anion species; PF6 − , BF4 − , and TFSI− , depending on the alkyl side chain length.
In addition, CF3 CO2 − and CF3 SO3 − anions tend to form relatively low viscosity
ILs. There are only a few studies of ILs containing Cl− or Br− anion [22], because
these ILs are not in the liquid state at room temperature. Since viscosity is directly
affected by electrostatic interaction, it is expected that ILs composed of larger ions
or charge delocalized ions should show lower viscosity. The degree of dissociation
of salts is another important factor.
3.4
Density
Table 3.6 summarizes the densities of various ILs. Since ILs are composed only
of ions, almost all ILs are denser than water, from 1.0 to 1.6 g cm−3 depending
on their ion structure. The densities of some complex salts are even higher than
ordinary ILs. Details will be given elsewhere.
56 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Table 3.6 Density of several ionic liquids
Cation
Anion
[EMIM]+
[NO3 ]−
[BF4 ]−
[PF6 ]−
[CF3 COO]−
[C3 F7 COO]−
[CH3 SO3 ]−
[CF3 SO3 ]−
[(CF3 SO2 )2 N]−
[(C2 F5 SO2 )2 N]−
[(CN)2 N]−
[(CN)3 C]−
[PF6 ]−
[BF4 ]−
[CF3 SO3 ]−
[(CF3 SO2 )2 N]−
[Cl]−
[PF6 ]−
[(CF3 SO2 )2 N]−
[BMIM]+
[HMIM]+
q /g cm−3 Ref.
1.28
1.28
1.56
1.29
1.45
1.25
1.38
1.46
1.52
1.08
1.11
1.37
1.21
1.30
1.43
1.03
1.31
1.37
Cation
60
[OMIM]+
24a
24a
6
6 [b-diMIM]+
60
6
[bpyr]+
6
32
[N3111 ]+
60
[N4111 ]+
28
[N6222 ]+
60
[N8222 ]+
60
[P14 ]+
6
[P13 ]+
6
[S111 ]+
60
[S222 ]+
60
[S444 ]+
60
Anion
q /g cm−3 Ref.
[BF4 ]−
[PF6 ]−
[(CF3 SO2 )2 N]−
1.12
1.23
1.31
60
60
60
[BF4 ]−
[PF6 ]−
[BF4 ]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
1.20
1.35
1.22
1.45
1.44
1.39
1.27
1.25
1.41
1.44
1.58
1.46
1.29
60
60
24a
24a
61
62
18
18
12
12
63
63
63
[b-diMIM]+ : 1-butyl-2,3-dimethylimidazolium, [bpyr]+ : butylpyridinium.
The density has been found to decrease with increasing alkyl chain length on the
imidazolium cation [33]. Similarly, in the ammonium and sulfonium salts, the density decreases with increasing alkyl chain length. This clearly shows that the charged
ion unit is heavier than the hydrocarbon chain. Accordingly, the density of ILs is
tunable to some extent. The density of aromatic onium salts is higher than that of
aliphatic ammonium salts. Generally, density decreases in the order of pyridinium
salts > imidazolium salts > aliphatic ammonium salts and piperidinium salts.
The densities of ILs are also affected by the anion species. Similarly to the
trends for cations, the density of ILs decreases with increasing alkyl chain length
of the anion. The density of ILs is increased on the introduction of a “heavy” chain
such as fluoroalkyl chains. For example, 1-ethyl-3-methylimidazolium (EMIM) salts
became heavier with the following anion species; CH3 SO3 − < BF4 − and CF3 COO−
< CF3 SO3 − < (CF3 SO2 )2 N− < (C2 F5 SO2 )2 N− . It is easy to understand this order as
an effect of formula weight of the ions. However, these tendencies are still empirical,
and a perfect correlation between ion structure and density is not yet available.
3.5
Refractive Index
In the field of processing or engineering, there is a potential requirement for
materials with high refractive index. However, these materials are typically all
solid and those liquids that are known are poisonous. Accordingly liquids having
3.5 Refractive Index 57
Table 3.7 Refractive index of ionic liquids
Cation
Anion
[EMIM]+
[BMIM]+
[HMIM]+
[OMIM]+
[decyl-MIM]+
[EMIM]+
[BMIM]+
[HMIM]+
n
T/◦ C
Ref.
[CH3 CO2 ]−
[CF3 SO3 ]−
[(CF3 SO2 )2 N]−
[Br]−
[I]−
[CH3 CO2 ]−
[BF4 ]−
[CF3 SO3 ]−
[(CF3 SO2 )2 N]−
[Cl]−
[Cl]−
[BF4 ]−
[BF4 ]−
1.4405
1.4332
1.4231
1.54
1.572
1.4887
1.42
1.4380
1.4271
1.515
1.505
1.4322
1.4367
20
20
20
25
25
20
25
20
20
25
25
25
25
6
6
6
30
64
6
64
6
6
30
30
65
65
[IBr2 ]−
[BrI2 ]−
[I5 ]−
[Br3 ]−
[IBr2 ]−
[BrI2 ]−
[I5 ]−
[I7 ]−
[I9 ]−
[IBr2 ]−
[I3 ]−
1.715
1.833
2.23
1.699
1.701
1.805
2.16
2.3
2.4
1.685
1.88
––
––
––
––
––
––
––
––
––
––
––
67
67
67
67
67
67
67
67
67
67
67
[decyl-MIM]+ : 1-decyl-3-methylimidazolium cation.
high refractive index with low toxicity are highly desirable. The refractive index of
water or methanol is ca. 1.33 at room temperature, whereas that of ILs is 1.4 or
more. Higher values (ca. 1.5) have been found in the ILs composed of halogenated
anions. Table 3.7 summarizes the refractive index for imidazolium salts composed
of various anions.
From this table, it is clear that refractive index of ILs increases with increase in
the alkyl chain length of the imidazolium cation. Moreover, the refractive index is
strongly affected by the anion; the refractive index of ILs with larger anions, such
as [(CF3 SO2 )2 N]− (TFSI anion) is lower than that of ILs having smaller anions such
as acetate or halide anions.
As mentioned before, the densities of ILs depend on the ion structure. There
is also a tendency for ILs with lower density to show higher refractive index. However, it was reported that the ILs containing complex metal anions
([BMIM][Ln(NCS)x (H2 O)y ] show relatively high refractive index (ca. 1.57) [66]. In
the case of these ILs, a clear correlation between refractive index and density is
observed; higher refractive index is found for the ILs with higher density (Figure
3.3). This tendency is the opposite of what occurs with common ILs.
58 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Fig. 3.3 Plot of refractive indices (n) vs. density (D) for a series of ionic liquids.
Furthermore, Seddon et al. reported that the poly-halide salts, such as
[EMIM][IBr2 ] or [EMIM][I5 ], have a high refractive index of 1.6 or more, as shown
in Table 3.7 [67]. The high refractive indices of the lanthanide salts and the heavy
halogens and their trihalide salts are well predictable from their polarizabilities,
which in turn are well understood on the basis of periodic table trends: atoms/ions
with partly filled 4f, 5d etc. shells tend to be quite polarizable and hence have high
refractive indices.
3.6
Polarity
Polarity is one of the most important parameters of ILs for its effect on electrochemical reactions. It is important when we characterize ILs to measure not only
thermal properties such as melting point but also solvent properties such as polarity
[68–70]. The most common method of polarity measurement is a dielectric constant measurement. Weingartner et al. and Hefter et al. have shown, by applying
appropriately high frequency methods, that the dielectric constants are uniformly
around 10–15. Accordingly, the polarity of ILs should be estimated by other methods. Solvatochromism is heavily applied for this purpose due to its simplicity.
3.6.1
Solvatochromism
Solvatochromism is the shift of the maximum absorption wavelength of dye
molecules depending on the polarity of the solvents. The advantage of this method
3.6 Polarity 59
is the small amount of sample required for spectroscopic measurement. A review
by Reichardt introduced over 80 probe molecules [71].
3.6.2
Reichardt’s Betaine Dye
There are many kinds of probe molecules for estimation of polarity. Among them,
the most widely used dye is Reichardt’s dye (2,4,6-triphenylpyridinium-N-4-(2,6diphenylphenoxide) betaine) [72]. Both empirical scales for polarity; E T (30) and
E T N are frequently used for polarity studies in ILs. The solvatochromism of the
Reichardt’s dye is based on the interaction of the solvent with the ground state of the
dye. Kamlet and Taft proposed that about two-thirds of the shift of the maximum
absorption wavelength of Reichardt’s Dye could be assigned to the interactions
involving the phenoxide oxygen with the solvent [73]. E T (30) is estimated by Eq.
(3.1). Reichardt and Harbusch-Gornert have defined an E T N value according to
Eq. (3.2) as a dimensionless figure, using water and tetramethylsilane (TMS) as
references of extreme polar and non-polar solvents, respectively. Hence, the E T N
scale ranges from 0 to 1 [74].
E T (30) = 28591.5/λmax
E TN
= [E T (solvent) − 30.7]/32.4
(3.1)
(3.2)
The E T (30) (in kcal mol−1 ) and E T N scales of ILs are summarized in Table 3.8.
The E T (30) for 1-methyl-3-alkylimidazolium type ILs is about 51–53, similar to
that of methanol and ethanol. The alkyl chain length and the nature of the anion have no influence on the E T (30). For the E T (30) values of alkylimidazoliumtype ILs, substitution of C-2 proton with a methyl group lowered the E T (30) to
48–49, which is similar to that of octanol or isopropanol. On the other hand,
[HO(CH2 )2 MIM][(CF3 SO2 )2 N], which contains a hydroxy group, has a high E T (30)
value (61.4). This suggests that the C-2 proton on the imidazolium ring shows
high acidity/hydrogen bond donor capability; furthermore the hydroxy group on
the side chain shows even higher acidity/ hydrogen bond donor capability than
the C-2 proton. For the aliphatic cations, the E T (30) decreased in the following
order: primary > tertiary > quaternary. Additionally, for quaternary ammonium
cations, the values decrease with increasing cation size. Generally, E T (30) values
are dominated by the nature of the cations.
3.6.3
Kamlet–Taft Parameter
Kamlet–Taft parameters are known to express three distinct measures of the solvent polarity such as dipolarity/polarizability (π * ), hydrogen-bond acidity (α) and
hydrogen-bond basicity (β). These parameters have been determined by absorption
measurements for individual or pairs of the following dye molecules; N,N-diethyl4-nitroaniline, 4-netroaniline and Reichardt’s dye, as seen in Figure 3.4 [81–83].
60 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Table 3.8 The value of ET (30) and ET N for some ionic liquids
Cation
[EMIM]+
[PMIM]+
[BMIM]+
[HMIM]+
[OMIM]+
[DMIM]+
[b-diMIM]+
[BZMIM]+
[OH(CH2 )2 -MIM]+
[P14 ]+
n-butylammonium
sec-butylammonium
dipropylammonium
ethylammonium
n-propylammonium
tributylammonium
tetrabutylammonium
tetrapropylammonium
tetrapentylammonium
water
methanol
ethanol
acetonitrile
acetone
dichloromethane
toluene
hexane
dimethyl sulfoxide
Anion
ET (30)
ET N
Ref.
[BF4 ]−
[(CF3 SO2 )2 N]−
[BF4 ]−
[(CF3 SO2 )2 N]−
[BF4 ]−
[PF6 ]−
[TfO]−
[(CF3 SO2 )2 N]−
[SbF6 ]−
[CF3 SO3 ]−
[(CF3 SO2 )2 N]−
[BF4 ]−
[PF6 ]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[BF4 ]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[SCN]−
[SCN]−
[SCN]−
[NO3 ]−
[NO3 ]−
[NO3 ]−
[CHES]−
[CHES]−
[(CF3 SO2 )2 N]−
53.7
52.9
53.1
52.0
52.5
52.3
51.2
51.5
52.4
52.3
51.9
48.3
51.2
51.1
51
49.4
48.6
52.5
61.4
48.3
61.4
61.6
63.3
61.6
60.6
56.7
48
51
44
63.1
55.5
51.9
46
42.2
40.7
33.9
0.71
0.69
0.69
0.65
0.67
0.67
0.66
0.64
0.67
0.67
0.65
0.54
0.63
0.63
0.63
0.58
0.54
0.67
0.95
0.54
0.95
0.95
1.01
0.95
0.92
0.8
0.62
0.53
0.41
1.00
0.77
0.65
0.47
0.36
0.309
0.1
(0.009
0.44
75
76
75
58
77
77
77
77
77
77
77
79
77
77
58
78
77
58
58
78
79
79
79
79
58
79
80
80
69
45
[OMIM]+ : 1-octyl-3-methylimidazolium, [DMIM]+ : 1-decyl-3-methylimidazolium, [b-diMIM]+ :
1-butyl-2,3-dimethylimidazolium, [BZMIM]+ : 1-benzyl-3-methylimidazolium, [OH-EMIM]+ :
1-hydroxyethyl-3-methylimidazolium, [CHES]− : 2-(cyclohexylamino)ethanesulfate
3.6 Polarity 61
Fig. 3.4 Solvatochromic probe molecules.
N,N-Diethyl-4-nitroaniline, has an aromatic ring but no hydrogen bond donor
substituent, shows a π–π * transition based on a non-specific interaction between
ions. Dipolarity/polarizability, π * , is estimated by the solvatochromic shift of N,Ndiethyl-4-nitroaniline using Eq. (3.3), where λmax is the absorption maximum for
N,N-diethyl-4-nitroaniline.
π ∗ = 8.649 − 0.314ν1 (ν1 = 1/(λmax × 10−4 ))
(3.3)
4-Nitroaniline can interact with solvent molecules with an amino group at the C-1
position as a proton donor. The λmax shows a red shift when it interacts with a solvent
having a hydrogen bond acceptor group. The β value (hydrogen bond basicity) is
estimated with Eq. (3.4).using spectral data of both N,N-diethyl-4-nitroaniline and
4-nitroaniline.
B = (1.035ν2 − ν1 + 2.64)/2.80
(3.4)
Reichardt’s dye has thus been used to estimate hydrogen-bond acidity of solvents.
The absorption maximum of Reichardt’s dye shows a blue shift when the solvent
molecule interacts with the dye through a hydrogen bond. The α value (hydrogenbond acidity) is estimated using E T (30) and π * with Eq. (3.5).
α = 0.0649E T (30) − 0.72π ∗ − 2.03
(3.5)
Welton reported the effect of cations and anions on the Kamlet–Taft Parameters
[78]. The Kamlet–Taft parameters for some ILs are summarized in Table 3.9. As
seen, π * values for these ILs are high, 0.9–1.3, in comparison with those for protic
molecular solvents as shown in the same table. Both cation and anion affect the π *
value. For anions, the π * value for ILs having TFSI anion is low due to weakened
coulombic interaction caused by delocalized anionic charge.
The β values of ILs are mainly governed by the nature of the anions. They decrease
in the order [Cl]− > [RSO3 ]− > [BF4 ]− > [PF6 ]− . On the other hand, α values of ILs
are largely affected by the nature of the component cations, especially the presence
of hydrogen-bond donor groups. The nature of the anion seldom affects the α value
[78].
62 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Table 3.9 Kamlet–Taft parameters for typical ionic liquids
Cation
[EMIM]+
[BMIM]+
[HMIM]+
[b-diMIM]+
[P14 ]+
n-butylammonium
sec-butylammonium
dipropylammonium
ethylammonium
n-propylammonium
tributylammonium
water
methanol
ethanol
acetonitrile
acetone
dichloromethane
toluene
hexane
dimethyl sulfoxide
Anion
[(CF3 SO2 )2 N]−
[BF4 ]−
[Cl]−
[PF6 ]−
[CF3 SO3 ]−
[(CF3 SO2 )2 N]−
[SbF6 ]−
[(CF3 SO2 )2 N]−
[BF4 ]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[SCN]−
[SCN]−
[SCN]−
[NO3 ]−
[NO3 ]−
[NO3 ]−
Kamlet-Taft parameters
Ref.
p*
α
β
0.980
1.047
1.17
1.032
1.006
0.984
1.04
0.971
1.083
1.01
0.954
1.23
1.28
1.16
1.24
1.17
0.97
0.705
0.627
0.41
0.634
0.625
0.617
0.64
0.259
0.402
0.381
0.427
0.92
0.91
0.97
0.85
0.88
0.84
0.233
0.376
0.95
0.207
0.464
0.243
0.15
0.650
0.363
0.239
0.252
1.09
0.6
0.54
0.75
0.71
0.791
0.532
(–0.12)
1
1.17
0.93
0.83
0.19
0.08
0.042
–0.213
(0.07)
0
0.18
0.62
0.77
0.31
0.48
–0.014
0.077
(0.04)
0.76
0.39
0.46
0.52
76
58
68
77
77
77
78
76
78
77
77
79
79
79
79
58
79
3.6.4
Acetylacetonatotetramethylethyldiaminecopper (II)
[Cu(acac)(tmen)]BPh4 , is known to provide a good correlation between the donor
number (DN) of the solvent and the λmax corresponding to the lowest energy of
d–d transition [84]. In spite of the small number of experiments, there is a certain
relation between anion species and λmax , as shown in Table 3.10.
3.6.5
Pyrene
Pyrene is one of the most widely studied neutral fluorescence probes, and accordingly, this was sometimes used to determine the polarity of some ILs. The
polarity scale of the IL analyzed with pyrene is defined as the emission intensity
3.6 Polarity 63
Table 3.10 Polarity of ionic liquids determined by Nile Red and
[Cu(acac)(tmen)]+ BPh4 −
Cation
Anion
[Cu(acac)(tmen)]
[BPh4 ]
Nile Red
kmax
[EMIM]+
[PMIM]+
[BMIM]+
[HMIM]+
[OMIM]+
[DMIM]+
[BZMIM]+
[HMIM]+
[HEIM]+
[HBIM]+
water
methanol
ethanol
acetonitrile
acetone
hexane
DMSO
[BF4 ]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[BF4 ]−
[PF6 ]−
[CF3 SO3 ]−
[(CF3 SO2 )2 N]−
[NO3 ]−
[BF4 ]−
[PF6 ]−
[NO3 ]−
[BF4 ]−
[PF6 ]−
[(CF3 SO2 )2 N]−
[NO3 ]−
[(CF3 SO2 )2 N]−
[BF4 ]−
[(CF3 SO2 )2 N]−
[BF4 ]−
[BF4 ]−
[BF4 ]−
ENR
Ref.
550.3
550.8
547.5
52.3
51.9
52.2
88
88
88
548.7
555.7
551.9
551.7
552.9
549.5
549.8
52.1
51.5
51.8
51.8
51.7
52
52
88
88
88
88
88
88
88
550.1
560.5
545.7
546
562.3
562.9
562.8
584.5
542.9
539.8
520.7
217.4
52.2
52.4
51.8
50.9
50.8
50.8
48.2
52
52.2
53.8
88
88
88
88
88
89
89
544.8
59
52
kmax
Ref.
541
547
69
69
517
516.5
546
602
77
517
549
77
77
77
573
569
ratio “II /IIII ”, where band I corresponds to an S1 (ν = 0)→S0 (ν = 0) transition
(at 373 nm), and band III is an S1 (ν = 0)→S0 (ν = 1) transition (at 384 nm). The
“II /IIII ” emission intensity ratio is known to increase with increasing solvent polarity [85–87]. The II /IIII ratio for monoalkylammonium thiocyanates is 1.01–1.23. In
the case of [EMIM][(CF3 SO2 )2 N], it is 0.85, and [BMIM][PF6 ] shows a particularly
high ratio: 2.08 (cf. water = 1.87, acetonitrile = 1.79, methanol = 1.35) [68].
Additionally, estimation of the dielectric constant for some ILs has been carried
out from these measurements. Since [EMIM][(CF3 SO2 )2 N] shows a λmax at 431 nm,
the dielectric constant is estimated to be lower than 10 since λmax is shorter than
that in hexanol (dielectric constant: 13.5) [69]. This correlates well with the direct
measurements of dielectric constant by Weingartner.
64 3 Physical Properties of Ionic Liquids for Electrochemical Applications
3.6.6
Nile Red
Nile Red shows positive solvatochromism. The degree of λmax shift is known to
depend on the dipolarity/polarizability of the medium. The λmax and the calculated
E NR values of Nile Red are summarized in Table 3.10. For ILs of the [BMIM]+ cation,
the E NR value decreased with anions in the following order: [NO3 ]− > [BF4 ]− >
[(CF3 SO2 )2 N]− > [PF6 ]− [88, 89]. Additionally, Nile Red was applied to estimate
the polarity of protic ILs, while most other dyes were bleached in the presence of
protons [89].
It should be noted here that comparison of ILs via only one polarity parameter is
dangerous in discussion of the polarity of ILs.
3.7
Solubility of Metal Salts
Solubility of metal salts in ILs is extremely important in electrodeposition. In
this section, the solubility of metal salts in air stable ILs is summarized. The
solubility of metal salts in halometalate type ILs has been summarized in previous
reports [90, 91]. In addition, many IL systems have been reported as electrolytes for
lithium-ion secondary batteries. Some metal salts were reported to be soluble above
50 mol%. However, these systems were obtained by mixing ILs with metal salts in
organic solvent or water followed by removal of the solvent; this may produce
supersaturated solutions. In this section, these systems are omitted due to space
limitations.
In ILs, anions and/or cations have weakly coordinating properties, and this solvation energy is not large enough to break the electrostatic interactions between ions
or metal atoms in metal salts. Consequently, it is generally expected that common
ILs have very low solubilizing ability for metals or metal salts. Rogers et al. reported
the evaluation of distribution ratios of Cs+ , Na+ , Sr2+ , Cl− in [Cn MIM][PF6 ] (n = 4,
6, 8)/water mixtures [40b]. The distribution ratio is defined as the concentration ratio of solute in the IL phase to that in the aqueous phase. As shown in Table 3.11, all
distribution ratios are very low, such as 10−3 –10−2 . Although the solubility of these
ions in [Cn MIM][PF6 ] is unknown, it is expected from these results to be very low.
Alfonso et al. evaluated the solubility of LiCl, HgCl2 , and LaCl3 in [Cn MIM][BF4 ]
(n = 4, 8, 10) and [Cm MIM][PF6 ] (m = 4, 8) (Table 3.12 entries 1–5) [92]. The ILs
containing the BF4 anion solubilized these salts more than those containing the
PF6 anion. However, the highest solubility was around 10−4 wt%, still very low. In
addition, they prepared ILs from ions having ether or hydroxy groups, expecting
further interaction with ions. In these ILs the solubility of HgCl2 and LaCl3 was
certainly improved (Table 3.12 entries 6–12).
MacFarlane et al. evaluated the solubility of CoCl2 ·6H2 O and CuCl2 ·2H2 O in
[C2 MIM] dicyanamide [93]. Compared to traditional ILs containing the Tf2 N anion,
the IL containing DCA anion dissolved CoCl2 ·6H2 O and CuCl2 ·2H2 O to a greater
3.7 Solubility of Metal Salts 65
Table 3.11 Distribution ratios between RTIL/aqueous phases. Data
from Ref. [40d]
ion
[C4 MIM][PF6 ]
[C6 MIM][PF6 ]
[C8 MIM][PF6 ]
Na+
Cs+
Sr2+
Cl−
0.023
0.067
0.048
0.0017
0.011
0.068
0.029
0.0014
0.011
0.072
0.026
0.00041
Aqueous phase, pH 7
extent. In order to elevate the solubility, it is important to strengthen the interaction
between the ILs and the metal ion. As an approach to enhancing the interaction,
functional groups were incorporated in the cation or anion to prepare so-called taskspecific ILs (TSIL). Davis et al. synthesized TSIL with thioether or thiourea groups
introduced into a side chain of the imidazolium cation. These were effective in
extracting Hg2+ or Cd2+ ions from an aqueous phase [94]. Table 3.13 shows the
distribution ratios of Hg2+ and Cd2+ in TSIL/water mixtures. Although the ion
species are different, these distribution ratios were significantly improved. This is
attributed to the interaction between the sulfur atom in TSIL and Hg2+ or Cd2+ ions.
As another approach to enhancing the interaction between ILs and metal salts,
an extractant highly compatible with both ILs and the metal salt was added. There
are many reports on the extraction of metal ions from the aqueous phase into an
IL phase with such extractants. Typical examples include crown ethers [40b,95],
Table 3.12 Observed solubility constants (K s ) of inorganic salts in sev-
eral ionic liquids. Data from Ref. [92]
Entry
1
2
3
4
5
6
7
8
9
10
11
12
a
K s [a]
Ionic liquids
Cation
Anion
LiCl
HgCl2
LaCl3
[C4 MIM]+
[C4 MIM]+
[C8 MIM]+
[C8 MIM]+
[C10 MIM]+
[C2 OHMIM]+
[C2 OHMIM]+
[C3 OMIM]+
[C3 OMIM]+
[C5 O2 MIM]+
[C5 O2 MIM]+
[C5 O2 MIM]+
[PF6 ]−
[BF4 ]−
[PF6 ]−
[BF4 ]−
[BF4 ]−
[PF6 ]−
[BF4 ]−
[PF6 ]−
[BF4 ]−
[Cl]−
[PF6 ]−
[BF4 ]−
12.08
15.54
35.32
56.02
12.64
144.47
18.46
12.44
14.43
9.98
35.52
21.36
4.06
41.41
32.98
35.92
2.12
44.64
84.73
50.13
220.86
295.34
147.48
174.17
6.58
10.92
8.49
53.25
47.12
32.47
54.01
37.61
180.27
379.23
97.22
292.46
Observed K s (10−6 g of salt g−1 of ILs)
66 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Table 3.13 Distribution ratio for Hg2+ and Cd2+ in the mixed systems
of water and TSIL 1 or TSIL 2. Data from Ref. [94]
TSIL
Cation
pH
Distribution ratio
Hg2+
1
198
Hg2+
7
1
7
1
7
1
7
208
330
376
346
343
20
23
Cd2+
Cd2+
Hg2+
Hg2+
Cd2+
Cd2+
molecules containing phosphine oxide groups [96], and calixarenes [97]. The development of ILs having improved coordinating properties for metal salts will be an
important area of study in the future.
3.8
Electrochemical Properties
3.8.1
Potential Window
For electrochemical applications, the potential window of the electrolyte solution
is one of the important properties. The potential window is governed not only by
the chemical structure of the materials used but also by the electrode materials,
sweep rate of the potential, temperature, atmosphere, solvent, impurity and so
on. Since values of potential windows in the literature have been evaluated under
various conditions, it is not easy to compare the values. Use of various reference
electrodes (RE) for the determination of cathodic and anodic limits of ILs makes
the situation even more complicated. At least, the potential of REs should be
confirmed with common redox potentials for non-aqueous systems. For example,
the ferrocene(Fc)/ferrocenium(Fc+ ) redox couple is helpful as a standard for many
ILs. Although Ag/AgCl(aq), Ag/Ag+ (organic solvents) and pseudo-metal electrodes
such as Ag wire and Pt wire are often used as REs, these are not stable enough due
to the generation of unstable membrane potentials, chemical reactions on the metal
surface, and so on. As a stable RE, Katayama et al. reported an Ag/Ag+ (IL) reference
consisting of a silver wire inserted in a silver salt/IL solution as the inner solution
[98]. The Ag/Ag+ (IL) reference is stable also in the measurement under specific
conditions (under reduced pressure, high temperature, dry atmosphere, etc.). The
potential windows are usually evaluated by cyclic voltammetry (CV) or linear sweep
voltammetry (LSV). In the CV method, it must be noted that the electrochemically
oxidized (or reduced) products of the first sweep must affect the voltammograms
of the reverse sweeps. Such effects do not appear in the LSV method, since fresh
3.8 Electrochemical Properties 67
test solution and electrodes are employed for anodic sweep and cathodic sweep,
respectively. Instead, reproducibility must be checked for LSV measurements. In
both cases, the anodic and cathodic limits are defined as the voltage where the
current density reaches a certain value. The cut-off current density is generally
1.0 mA cm−1 with a sweeping rate of 50 mV s−1 . Generally, cathodic and anodic
limits of pure ILs are attributed to the oxidative decomposition of the anion and
the reductive decomposition of the cation, respectively. Impurities, especially water
and halide anions, must be removed carefully, otherwise these drastically narrow
the potential window.
Table 3.14 shows a series of potential windows and the conditions of measurement for a series of ILs. The potential windows of imidazolium salts are around
4 V. Imidazolium salts having active protons on the 2-carbon sometimes decompose easily. In fact, the potential windows of imidazolium salts are wider when
Table 3.14 Electrochemical windows for a variety of ionic liquids
Cation
[EMIM]+
[e-diMIM]+
[N1113 ]+
[N1114 ]+
[N1114 ]+
[N111,2O1 ]+
[N1224 ]+
[N122,2O1 ]+
[N2226 ]+
[P14 ]+
[P14 ]+
[PP13 ]+
[C2 -dabco]+
[S222 ]+
[S444 ]+
Anion
Working
electrode
Reference
electrode
Potential
window[a]
Ref.
[BF4 ]−
[BF4 ]−
[BF4 ]−
[BF4 ]−
[(CN)2 N]−
[CF3 CO2 ]−
[CF3 SO3 ]−
[CF3 BF3 ]−
[(CF3 SO2 )(CF3 CO)N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 CF2 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[CF3 BF3 ]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[C2 F5 BF3 ]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
Pt
GC
Pt
Pt
Pt
Pt
Pt
GC
GC
Pt
GC
GC
Pt
GC
GC
GC
GC
GC
Pt
GC
GC
GC
GC
Pt
GC
GC
Al/Al3+
Al/Al3+
I− /I3 −
Ag wire
Ag wire
I− /I3 −
I− /I3 −
Fc/Fc+
Fc/Fc+
I− /I3 −
Ag wire
Ag wire
I− /I3 −
I− /I3 −
Fc/Fc+
Ag wire
I− /I3 −
Fc/Fc+
Ag/AgCl aq
Ag wire
Ag wire
Fc/Fc+
I− /I3 −
Fc/Fc+
I− /I3 −
I− /I3 −
4.4
> 2.1[b]
4.4
4.4
3.0
3.2
3.8
4.6
3.1
4.2
4.1
4.1
4.4
5.8
5.9
5.6
5.4
5.8
5.8
5.6
5.5
5.4
5.8
5.0
5.2
5.2
105
105
106
107
108
6
6
109
101
6
7
7
6
61
110
57
61
110
102
57
57
111
103
112
63
63
a
Many of the potential windows were estimated from the voltammograms shown in the reference
papers; cut off current density ∼ 1 mA cm−2 . b Anodic limit was not given. [e-diMIM]+ :
1-ethyl-2,3-dimethylimidazolium, [PP13 ]+ : N-methyl-N-(n-propyl)piperidinium [C2 -dabco]+ :
N-ethyl-1,4-diazabicyclo[2.2.2]octane.
68 3 Physical Properties of Ionic Liquids for Electrochemical Applications
the 2-position is substituted by an alkyl chain. However, 2-substituted imidazolium
salts generally have higher melting temperatures or higher viscosity than unsubstituted ones. Aliphatic cations such as ammonium cations and piperidinium cations
are relatively strong against both oxidation and reduction. Therefore, their potential
windows are usually around or wider than 5 V. Thus far, ionic liquids having a potential window over 7 V have also been reported [99]. Generally, TFSI anion-based
ILs have relatively wide electrochemical windows on a wide variety of electrodes.
Also, BF4 − -based ILs have good properties, but it must be noted here that this anion
is not stable against carbon electrodes [100].
Since some ILs have excellent electrochemical stability, as shown in Table 3.14,
they are favorable for application as electrolyte materials. Recently, ionic liquids
have been investigated as conductive and redox media for lithium ions. Stable
electrochemical deposition and dissolution of Li metal (Li/Li+ ) was observed for the
lithium salt solution of [N1113 ][TFSI], [N122,2O1 ][TFSI], and [PP14 ][TFSI] [101–103].
In order to observe the redox couple of lithium metal, Ni should be used as working
electrode because it does not form alloys with lithium metal. In addition to this,
the atmosphere must be pure Ar, because Li metal reacts rapidly with N2 to form
conductive LiN.
The potential window is one of the most important physical properties for the
selection of a solvent for electrolysis. However, it should also be noted that the
surface layer on the electrode, which is formed by chemical or electrochemical
deposition, often stabilizes the system. For example, Katayama reported that a
lithium ion conductive passivated layer, which is formed on the tungsten electrode
as a result of reductive decomposition of the cation of the IL during the first sweeps,
enables the reversible deposition and dissociation of Li metal [104]. Howlett et al.
have also discussed this extensively [112a–c].
This surface film formation is being proposed to protect corrosion of some
reactive metals [112d].
3.8.2
Ionic Conductivity
The ionic conductivity (σ i ) can be described by the following equation.
ni e i µi
i =
(3.6)
where ni is the number of ith ions, e is the charge of an electron and µi is the
mobility of the ith ion. The net ionic conductivity is the sum of the product for each
effective carrier ion species in the system. In order to compare the ionic conductivity
of some ILs, one has to note that every IL has a different ion concentration (n).
Therefore, the molar conductivity () is usually helpful to know the contribution
of the ion mobility (µ) for the ionic conductivity.
= σi /d
where d is the salt concentration in mol L−1 .
(3.7)
3.8 Electrochemical Properties 69
For classical dilute aqueous electrolyte solutions, where the salts are perfectly
dissociated, the molar conductivity is governed by the viscosity of the system.
(3.8)
σ η = constant
This equation is known as Walden’s rule. The constant is called the Walden
product. Although the salt contents of bulk ILs are very high (about 3–7 mol L−1 ),
the Walden plots for a variety of ILs are similar to that of a conventional diluted
system [113]. This observation indicates that ILs are ionized effectively, even in the
bulk. However, ILs also contain ion aggregates which do not contribute to the ionic
conductivity. Recent research shows more specifically how much ILs are ionized
[114].
The Arrhenius plot of the viscosity of the ILs is not a straight line but a
Vogel–Fulcher–Tamman (VFT) type curve. Since ionic conductivity is the inverse
of the viscosity (Eq. (3.8)), it also obeys the VFT equation.
σi = σ0 exp
−B
T − T0
(3.9)
where σ 0 and B are constants and T 0 is the ideal T g .
Equation (3.9) clearly indicates that ionic conductivity could be improved by
lowering the T g of the system. The difference in the temperature dependences of
ionic conductivity (and viscosity) for ion-conductive glass-forming materials has
been discussed by Angell et al. using “fragility” parameters [115].
So far, the ionic conductivity of most ILs has been measured by the complex
impedance method [116]. In this method, charge transfer between carrier ions and
electrode is not necessary. Therefore platinum and stainless steel are frequently
used as “blocking” electrodes. However, it is often difficult to distinguish the resistance and dielectric properties from Nyquist plots obtained by the impedance
measurement. In order to clarify this, additional measurements using non-blocking
electrodes or DC polarization measurement are needed.
The ionic conductivities of ILs are lower than those of conventional aqueous
electrolytes, since the viscosity of ILs is generally high (> 30 cP, except for some
systems). However, comparing with salt solutions having similar viscosity such as
oligo(ethylene oxide)/lithium salt solutions [117], ILs show even higher ionic conductivity because of the much larger number of carrier ions. The ionic conductivity
and related properties of [EMIM] salts are summarized in Table 3.15. Among them,
[EMIM][TFSI] and [EMIM][BF4 ] show both relatively high ionic conductivity and
low viscosity. Imidazolium salts are known to show higher ionic conductivity than
those of ammonium ones having similar formula weight. The effect of the alkyl
chain length on the ions is also obvious. [EMIM][TFSI] shows the maximum ionic
conductivity among the [TFSI]-based imidazolium salts, but further elongation of
the alkyl chain causes a decrease in conductivity.
We have also investigated the ion conductive properties of a series of “neutralized”
ILs, prepared by neutralization of amines with equimolar amounts of Brønsted
70 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Table 3.15 Specific ionic conductivity and related properties of imidazolium salts at 25 ◦ C
Cation
[EMIM]+
Anion
Conductivity
r i /mS cm−1
Molar conductivity
/ S cm2 mol−1
Ref.
[BF]−
[BF4 ]−
[BF4 ]−
[PF6 ]−
[CF3 SO3 ]−
[CF3 CO2 ]−
[C3 F7 CO2 ]−
[CH3 COO]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
[(CF3 CF2 SO2 )2 N]−
[C(CN)3 ]−
[(CN)2 N]−
[NbF6 ]−
[TaF6 ]−
[CH3 BF3 ]−
[C2 H5 BF3 ]−
[n-C3 H7 BF3 ]−
[n-C4 H9 BF3 ]−
[n-C5 H11 BF3 ]−
[CH2 CHBF3 ]−
[CF3 BF3 ]−
[C2 F5 BF3 ]−
[n-C3 F7 BF3 ]−
[n-C4 F9 BF3 ]−
13.6
13.6
13 (26)
5.2 (26)
8.6 (20)
9.6 (20)
2.7 (20)
2.8 (20)
8.8 (20)
5.7
8.4 (26)
9.0
3.4 (26)
180 (20)
270 (20)
8.5
7.1
9.0
6.3
5.7
3.2
2.7
10.5
14.8
12.0
8.6
5.2
2.1
2.1
2.0 (26)
24a)
34
7
7
6
6
6
6
6
24a)
7
121
7
28
28
120
120
34
34
34
34
34
34
109
109
109
109
1.5
2.1 (26)
2.3
1.1 (26)
32.8 (20)
44.3 (20)
1.6
1.3
1.5
1.2
1.1
0.7
0.6
1.9
2.7
2.5
2.0
1.3
The number in parentheses is the measurement temperature.
acids [118]. Some conductivity values are shown in Table 3.16. Physical properties
of the neutralized ILs showed similar trends to those of the quaternary ones. Since
neutralized ILs are easy to prepare, these are useful models to find candidate ions
for new ILs. They are also expected to be proton conductors.
Some ion conductive properties of lithium salt/ IL solutions are summarized
in Table 3.17. Generally, the ionic conductivity of ILs containing lithium salts are
lower than those of pure ILs, even though the addition of lithium salt increases
the net number of ions in the system, due to smaller formula weight of lithium.
According to the literature, the major reason for these phenomena may be the
increase in viscosity and T g (some specific values are shown in Table 3.17). The
aggregation of lithium ions in ILs, which causes the decrease in the effective carrier
ion number, might be another reason for the decrease in ionic conductivity.
3.8 Electrochemical Properties 71
Table 3.16 Ionic conductivity (25◦ C) of amines neutralized by HBF4 .
Data from Ref. [118]
Amine
Structure
Conductivity r i /mS cm−1
Tg/ ◦C
T m /◦ C
1-methylpyrazole
19
−109.3
−5.9
2-methyl-1-pyrroline
16
−94.3
17.1
1-methylpyrrolidine
16
––
−31.9
1-ethylcarbazole
2.2
−68.0
––
2,3-lutidine
5.9 × 10−3
––
59.4
2,6-lutidine
< 10−4
−10.9
104.6
pyrrole
< 10−5
0.1
––
1-methylpyrrole
< 10−5
−15.9
––
Although the lithium ion transference numbers in lithium salt/ IL solutions
are important, especially for battery applications, few literature reports refer to the
specific values. Since the component ions of the ILs themselves have high mobility,
the lithium ion transference number should be low for most cases. In order to
suppress the mobility of the component ions, we proposed zwitterion compounds
having an imidazolium cation structure [119]. Normally, such zwitterions are solid
72 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Table 3.17 Ionic conductivity of ionic liquids containing lithium salts at
25 ◦ C
Cation
Anion
Added salt, amount
Conductivity
r i /mS cm−1
Viscosity/ cP
Ref.
[EMIM]+
[(CF3 SO2 )2 N]−
––
LiTf2 N, 0.32 mol kg−1
10.6 (30)
6.6 (30)
122
[DMPIM]+
[(CF3 SO2 )2 N]−
––
LiTf2 N, 0.32 mol kg−1
3.41 (30)
2.1 (30)
122
[EMIM]+
[(CF3 SO2 )2 N]−
––
LiTf2 N, 1m
10
2
30
∼200
123
[EMIM]+
[BF4 ]−
––
LiTf2 N,1 m
15
7
36
∼100
123
[N111,1-CN ]+
[(CF3 SO2 )2 N]−
––
LiTf2 N, 0.2 mol L−1
∼10−4
slight increase
124
[N1114 ]+
[(CF3 SO2 )2 N]−
––
LiTf2 N, 0.2 mol L−1
∼10−3
slight decrease
124
[N112, 2O1 ]+
[(CF3 SO2 )2 N]−
––
LiTf2 N, 0.9 mol L−1
4.0 (30)
0.37 (30)
∼100 (20)
300 (20)
102
[N112, 2O1 ]+
[BF4 ]−
––
LiTf2 N, 0.9 mol L−1
∼10−2 (30)
0.34 (30)
∼1000 (20)
3450 (20)
102
[DEDMIM]+
[(CF3 SO2 )2 N]−
[(CF3 SO2 )2 N]−
2.7 (20)
1.4 (20)
0.8 (20)
< 0.0001
> 0.001
125
[C2dabco]+
––
LiTf2 N, 0.4 mol L−1
LiTf2 N, 0.8 mol L−1
––
LiTf2 N, 33 mol %
[DEMPZ]+
[(CF3 SO2 )2 N]−
––
LiTf2 N, 10 mol %
2.6 (20)
1.7 (20)
126
112
The numbers in parentheses are the temperature of measurement, [DMPIM]+ :
1,2-dimethyl-3-(n-propyl)imidazolium, [DEDMIM]+ : 1,2-diethyl-3,4-dimethylimidazolium,
[DEMPZ]+ : N,N-diethyl-3-methylpyrazolium.
at room temperature and obviously show almost no ionic conductivity. However,
the zwitterions were readily changed to liquid by mixing with suitable lithium salts
having soft anions. The ionic conductivity drastically increased after adding lithium
salts, as shown in Table 3.18. The lithium ion transference number of such lithium
salt/zwitterion mixtures was estimated to be higher than 0.5. The usefulness of
zwitterions will be mentioned in Section 3.8.4.2.
Recently, proton conductive ILs and iodide ion conductive ILs have also been
investigated separately. These ILs for specific ion transport are quite important for
the development of energy devices such as lithium batteries, fuel cells, and solar
cells. This will be discussed further in the next section.
3.8 Electrochemical Properties 73
Table 3.18 Ionic conductivity of zwitterions containing equimolar
lithium salts
+ LiTFSI
+ LiBETI
+ LiCF3 SO3
+ LiBF4
+ LiClO4
r i /mS cm−1 at 100◦ C
Tm/ ◦C
< 10 −5
0.89
6.1 × 10−2
7.5 × 10−3
< 10−3
< 10−3
175
––
––
––
––
––
Tg/ ◦C
18
−37
−5
19
4
24
3.8.3
Diffusion Coefficients of Component Ions
In spite of the high ionic conductivity, there is no guarantee that the IL can transport
the desired ions such as metal ions or protons. It is therefore important to analyze
the ion transport properties in ILs. The ion conduction mechanism in ILs is different from that in molecular solvents. The ionic conductivity is generally coupled to
carrier ion migration and ionic conductivity (σ ) correlates to diffusion coefficient
(D) according to the Nernst–Einstein equation (see Eq. (3.10)) where n and q imply
the number of carrier ions and electric charge, respectively. R, T, and F stand for
the gas constant, the temperature in K, and the Faraday constant, respectively.
σ =
Dnq 2 F 2
RT
(3.10)
Compared with electrochemical measurements of ion mobility and diffusion
in ion conductive materials [127], pulse-field-gradient NMR (PFG-NMR) is useful
to directly measure the diffusion coefficient of ions containing measurable nuclei [128]. In general, diffusion coefficients of 1 H, 13 C, 19 F, and 7 Li are frequently
measured as target nuclei and the values obtained are used to calculate the diffusivity parameters of the corresponding component ions. Therefore, ions having
no measurable nuclei for NMR cannot be analyzed. Many low-viscosity ILs are
composed of fluorinated anions such as BF4 − and TFSI− , and hence it is rather
easy to distinguish the diffusion behavior of anions from that of the onium cations.
From previous studies [24a,129], diffusion coefficients of relatively low viscosity
ILs such as [EMIM][BF4 ] and [EMIM][TFSI] are reported as around 10−11 m2 s−1
at room temperature. The diffusivity of [EMIM][AlCl4 ] measured by electrochemical methods [127] is similar in value to that of [EMIM][BF4 ] and [EMIM][TFSI].
Compared with water, in which diffusion coefficients are around 10−9 m2 s−1 at
room temperature [130], it is understandable that component ions in ILs find it
hard to diffuse due to strong electrostatic interaction forces. It is readily seen that
the difference in these diffusion coefficients arises from the higher viscosity of ILs
74 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Table 3.19 Physical properties and diffusion coefficients of ionic liquids
at 30 ◦ C. Data from Ref. [114]
Cation
[BMIM]+
[MMIM]+
[EMIM]+
[BMIM]+
[HMIM]+
[OMIM]+
[N4111 ]+
[bpy]+
[P14 ]+
Anion
[BF4 ]−
[PF6 ]−
[CF3 CO2 ]−
[CF3 SO3 ]−
[(CF3 SO2 ]2 N]−
[(C2 F5 SO2 ]2 N]−
[(CF3 SO2 )2 N]−
g/cP
75
182
58
64
40
87
31
27
40
56
71
77
49
60
D/10−11 m2 s−1
r /mS cm−1
4.5
1.9
3.8
3.6
4.6
1.9
11
11
4.6
2.7
1.6
2.6
4
3.4
Cation
Anion
t+
1.8
0.89
2.2
2.2
3.4
1.6
5.8
6.2
3.4
2.2
1.5
1.7
2.8
2.2
1.8
0.71
1.9
1.6
2.6
1.1
3.3
3.7
2.6
1.9
1.5
1.4
2.2
1.8
0.50
0.56
0.54
0.58
0.57
0.59
0.64
0.63
0.57
0.54
0.50
0.55
0.56
0.55
t+ = Dcation / (Dcation + Danion )
compared with that of water. Diffusion coefficients of ILs containing fluorinated
(or fluorine-containing) anions are generally large due to weaker electrostatic forces
[131]. Reported diffusion coefficients of relatively low viscosity ILs are summarized
in Table 3.19.
MacFarlane et al. [129] and Watanabe et al. [24a,114] discussed the difference
in diffusivity of component ions. Reported diffusion coefficients of ILs are shown
in Table 3.19 together with viscosity and ionic conductivity. From that table, it is
easy to see that lower viscosity ILs show larger diffusion coefficients and higher
ionic conductivity. Cations generally have larger diffusion coefficient values than
do anions in ILs. This means that the cation diffuses more easily than the anion.
However, the transference numbers of onium cation (t+ ) in ILs calculated from the
results of PFG-NMR is in the range 0.5 to 0.6 and their contribution to the ionic conductivity is mostly the same, irrespective of the ion species. In the case of [bpy][BF4 ],
the BF4 − shows a larger diffusion coefficient than that of bpy+ , and therefore t+
is below 0.5 [24a]. Thus, as well as thermal and electrochemical properties, the
diffusion behavior of component ions is dependent on their structure.
Diffusion coefficients (Dimp ) obtained from measurement are calculated via the
Nernst–Einstein equation. Furthermore, electrochemical diffusion coefficient measurements are possible which directly measure the diffusion coefficient. The degree
of dissociation of a component ion in the IL can be estimated from the relation
(DNMR /Dimp ) between Dimp and the diffusion coefficient measure by PFG-NMR
(DNMR ) [132]. This parameter is called the “Haven ratio” and should be unity
3.8 Electrochemical Properties 75
if all components completely dissociate into ions. In most cases, the diffusion
coefficient of ILs measured by PFG-NMR is larger than that calculated from
impedance measurements [62,129]. These results imply that part of the component ions of ILs do not contribute to ion conduction. The gap between the diffusion
coefficient and ionic conductivity is attributed to the fact that PFG-NMR could not
distinguish the dissociation state of ions, i.e., either ion or ion pair. As a result,
measured diffusion coefficients are an average value obtained from the summation
of diffusion coefficients of ions and ion pairs. The difference between DNMR and
Dimp shows that a fraction of the component ions form ion pairs or an aggregated
state. Therefore the carrier ion number is smaller than the calculated value based
on the molar concentration of the ILs.
Lithium cation transportation in ILs can be analyzed with PFG-NMR. The mixtures [P13 ][TFSI]/LiTFSI and [EMIM][BF4 ]/LiBF4 have been analyzed [133]. When
inorganic salts were added to ILs, their viscosity increased and accordingly ionic
conductivity decreased. In both reported mixture systems, the diffusion coefficient
of the component ions became smaller with increasing inorganic salt concentration. The diffusion coefficient of the lithium cation is the smallest among the ions
in the mixture. The lithium cation, which has a smaller ion radius than any of the
other component ions, has the strongest electrostatic interactions. This low lithium
cation transport number is one of the reasons why ILs are not currently applied as
substituents for electrolyte solutions in secondary batteries. Design of specific ILs
for target ion transport will be mentioned in the next section.
3.8.4
Ionic Liquids for Specific Ion Conduction
Physicochemical properties of ILs can be changed by variation of the component
ions. There are important studies to achieve ILs having excellent properties such as
low T m , low viscosity, high ionic conductivity and wide electrochemical potential
windows. It is generally understood that ILs are difficult to apply as electrolyte
solution substituents because they contain a large number of ions which cannot work as carrier ions for electrochemical devices such as secondary batteries.
Therefore, structural design of ions for particular applications is important for
ILs. Selective ion conduction is one of the attractive and challenging tasks for IL
science.
3.8.4.1 Ionic Liquids Containing Specific Ions
Significant differences between molecular solvents and ILs are based on the component species. This is not as a serious problem for the electrochemical devices that
require non-specific ions. The unique properties of ILs, especially their high ionic
conductivity, thermal stability and non-volatility are great advantages for electrolyte
solutions in electrochemical devices
These devices need long-range transportation of particular carrier cations such
as lithium cations or protons between electrodes. These small cations interact with
76 3 Physical Properties of Ionic Liquids for Electrochemical Applications
Fig. 3.5. Ionic liquids with a multivalent anion.
a counter anion with a strong electrostatic interaction and, accordingly, they are
difficult to migrate. Then, suitable carrier ions such as lithium cation and proton
should be added to ILs for these purposes.
It is easy to generate a target carrier ion for ILs if lithium salts or acids are added
to the ILs [134]. On the other hand, preparation of ILs composed of the required
cation is a better way to provide a higher concentration of the required cations in
ILs [135]. A room-temperature molten lithium salt (lithium IL) has been designed
by introducing the lithium salt structure into the tail of polar and flexible polymers,
or by using anions having highly delocalized negative charge [69]. ILs inherently
containing target ions have been prepared by the combination of an onium cation
and a multivalent anion. It is believed that the salts composed of multivalent ions
hardly melt at room temperature owing to their strong electrostatic interaction.
However, some multivalent anions give room temperature ILs by coupling with
specific onium cations like alkylimidazolium cations which have been known to
form good ILs [131,136]. Multivalent anions can interact with multiple cations
and form ILs containing the target cation as shown in Fig. 3.5. For this strategy,
sulfate, phosphate, phosphite, and pyrophosphate have been used to couple with
both imidazolium cation and protons [136b]. They showed excellent properties
with moderate ionic conductivity of 10−5 to 3 × 10−3 S cm−1 at room temperature;
with this strategy lithium ion-containing ILs can be prepared.
3.8.4.2 Selective Ion Conduction
Since component ions of ILs are highly mobile, these ions potentially move together with target small ions such as the lithium cation and proton. Component
ion migration should be inhibited in order to use the ILs for target ion transport.
Zwitterionic salts have been proposed as IL derivatives to inhibit component ion
migration along the potential gradient [119,137]. Zwitterionic salts in which both
the onium cation and the counter anion are tethered covalently cannot migrate
along with the potential gradient. Therefore, pure zwitterionic salts show no ionic
conductivity. They become conductive when salts are added. The mixtures can be
regarded as target ion transport materials. The cation and anion structures of zwitterionic salts are shown in Figure 3.6. Most of the prepared zwitterionic salts are
solid at room temperature. However, the salt containing an equimolar mixture of
LiTFSI and zwitterionic salt (having imidazolium cation and sulfonic acid anion) is
liquid at room temperature and shows ionic conductivity of 1.0 × 10−5 S cm−1
at 50 ◦ C and a lithium ion transference number of 0.56. Liquidization was
3.9 Conclusion and Future Prospects 77
Fig. 3.6 Structure of a zwitterionic salt for selective ion conductive materials.
explained by the formation of an IL-like domain with low T g between the cationic
part of the zwitterions and the anion of the added salt. Accordingly, the equimolar mixture of these can be regarded as an IL containing a negatively charged tail
(and its counter cation). These zwitterions have much potential for electrochemical
applications.
3.9
Conclusion and Future Prospects
In this chapter, the basic characteristics of ILs have been summarized. These basic
data should be helpful for the use of known ILs for electrochemical purposes and
for the design of ILs with better properties.. With the aid of chemistry, there are
increasing numbers of ILs being discovered and studied; there may be some superILs that remain undiscovered or not yet synthesized. Through these studies we will
be able to achieve better ILs.
Some properties of ILs have been analyzed and collected to construct a database
for future use as a guideline. A serious problem at present is the fluctuation of data.
Even ILs having identical structure have different data reported. These differences
are attributed mainly to trace amounts of contaminants and to different analytical
methods. Accordingly, the construction of an accurate database of highly pure ILs
is the burning issue. Their movement on this issue and a set of accurate data for
ultra pure ILs will be supplied in the near future. After a few years, most physicochemical properties of many ILs will be corrected and then we will be able to
obtain new strategies to accurately predict the properties from their component ion
structures.
Acknowledgement
The following coworkers in our laboratory should be acknowledged for their cooperation: Dr. Wataru Ogihara, Dr. Tomonobu Mizumo, Mr. Yukinobu Fukaya, Miss
Junko Kagimoto and Mr. Masahiro Tamada. Especially, the author would like to
thank Dr. Wataru Ogihara for his considerable contribution to the preparation of
this chapter.
78 3 Physical Properties of Ionic Liquids for Electrochemical Applications
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83
4
Electrodeposition of Metals
Thomas Schubert, Sherif Zein El Abedin, Andrew P. Abbott, Katy J. McKenzie,
Karl S. Ryder, and Frank Endres
Between 1980 and about 2000 most of the studies on the electrodeposition in ionic
liquids were performed in the first generation of ionic liquids, formerly called
“room-temperature molten salts” or “ambient temperature molten salts”. These
liquids are comparatively easy to synthesize from AlCl3 and organic halides such as
1-ethyl-3-methylimidazolium chloride. Aluminum can be quite easily be electrodeposited in these liquids as well as many relatively noble elements such as silver,
copper, palladium and others. Furthermore, technically important alloys such as
Al–Mg, Al–Cr and others can be made by electrochemical means. The major disadvantage of these liquids is their extreme sensitivity to moisture which requires handling under a controlled inert gas atmosphere. Furthermore, Al is relatively noble
so that silicon, tantalum, lithium and other reactive elements cannot be deposited
without Al codeposition. Section 4.1 gives an introduction to electrodeposition in
these first generation ionic liquids.
In the 1990s John Wilkes and coworkers introduced air- and water-stable ionic
liquids (see Chapter 2.2) which have attractive electrochemical windows (up to ±
3 V vs. NHE) and extremely low vapor pressures. Furthermore, they are free from
any aluminum species per se. Nevertheless, it took a while until the first electrodeposition experiments were published. The main reason might have been that purity
was a concern in the beginning, making reproducible results a challenge. Water
and halide were prominent impurities interfering with the dissolved metal salts
and/or the deposits. Today about 300 different ionic liquids with different qualities
are commercially available from several companies. Section 4.2 summarizes the
state-of-the-art of electrodeposition in air- and water-stable ionic liquids. These liquids are for example well suited to the electrodeposition of reactive elements such
as Ge, Si, Ta, Nb, Li and others.
Section 4.3 is devoted to electrodeposition in a special class of deep eutectic
solvents/ionic liquids which are based on well-priced educts such as e.g. choline
chloride. The impressive aspect of these liquids is their easy operation, even under
air, as well as their large-scale commercial availability. One disadvantage has to
be mentioned: the choline chloride-based liquids especially are currently not yet
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
84 4 Electrodeposition of Metals
suited to the electrodeposition of reactive elements such as aluminum or elemental
semiconductors like silicon.
In Section 4.4, finally, troublesome aspects are shortly summarized. An important aspect is that the electrochemical window alone is not sufficient and one
can be pretty surprised if the electroreduction of e.g. TaCl5 rather delivers nonstoichiometric halides instead of the desired tantalum metal. For an electroplating
bath the solution chemistry also plays an important role and a new concept of
additives seems to be necessary.
4.1
Electrodeposition in AlCl3 -based Ionic Liquids
4.1.1
Introduction
Historically, AlCl3 -based ionic liquids were the first to be used for the electrodeposition of metals. As described before, they are easy to synthesize by simple addition of
the Lewis acidic AlCl3 to a 1,3-dialkyl-imidazolium, alkyl-pyridinium or quaternary
ammonium compound under an inert atmosphere.
The main disadvantages of these materials are their corrosiveness and their
instability against air and moisture. Nevertheless, they have a kind of universal
character to dissolve other metal salts.
The Lewis acidity of these materials can be varied by varying the relative amount
of organic salt and AlCl3 : with a molar excess of AlCl3 they are Lewis acidic; with
a molar excess of the organic salt they are Lewis basic. To yield neutral liquids
it is necessary to buffer the 50:50 mol% mixture with NaCl, because minimum
variations from the equimolar composition would shift them towards basic or
acidic compositions [1].
It is well known that the chemistry and electrochemistry of many elements are
influenced significantly by the Lewis acidity of AlCl3 -based ionic liquids.
By far the most studies on metal and alloy deposition have been performed in
AlCl3 -based ionic liquids. In the following subsections all metals are categorized in
groups of the periodic system of the elements.
4.1.2
Group I Metals
Alkali metals have high oxidation–reduction potentials and low atomic masses.
Thus they are attractive candidates for anodes in secondary batteries. In this context, it was shown in a couple of investigations that lithium and sodium can be
electrodeposited from tetrachloroaluminate-based ionic liquids.
4.1.2.1 Electrodeposition of Lithium
Lithium is of particular interest for use as an anode material for secondary batteries
because it has the highest electricity storage density of the active metals.
4.1 Electrodeposition in AlCl3 -based Ionic Liquids
Fig. 4.1 Simplified electrochemical windows of 1-butyl-pyridinium
chloride and 1-ethyl-3-methyl-imidazolium chloride.
The first electrodeposition of lithium from an ionic liquid was reported in 1985
by Lipsztajn and Osteryoung [2]. They were able to deposit lithium from a 1-ethyl3-methyl-imidazolium chloride/aluminum trichloride ionic liquid. They noticed
that a “neutral” ionic liquid, a “neutral basic” ionic liquid (neutral + small excess
of RCl) and a “neutral acidic” ionic liquid (neutral + small excess of AlCl3 ) each
have unique features. Both the basic and the neutral acidic ionic liquids show an
extension of 1.5 V of the electrochemical window, making them interesting for
electrochemical applications.
They found that lithium chloride was not soluble in the neutral but dissolved
in the neutral acidic ionic liquid. From the latter no reduction of lithium was
observed prior to the cathodic limit of those ionic liquids. Thus they prepared first
the neutral acidic ionic liquid with a certain excess of AlCl3 and then added an
equivalent amount of lithium chloride to obtain a LiAlCl4 solution in a neutral
ionic liquid. From that they were able to reduce lithium ions on tungsten, glassy
carbon, and aluminum electrodes.
Piersma et al. demonstrated that lithium can be electrodeposited from 1-ethyl3-methyl-imidazolium tetrachloroaluminate ionic liquid, when lithium chloride
was dissolved in the melt [3]. Platinum, glassy carbon and tungsten were used as
working electrodes with molybdenum and platinum foils as counter electrodes.
At −2.3 V a reduction peak of Li+ is observed and at about −1.6 V the stripping
of lithium occurs. They noticed that the efficiency was much less than 100%. In
addition, they were able to demonstrate that the addition of proton sources like
triethanolamine·HCl widens the electrochemical window and allows the plating
and stripping of lithium (and also sodium).
85
86 4 Electrodeposition of Metals
Fig. 4.2 Cathodic scan cyclic voltammograms of near-Lewis neutral
and LiCl-buffered [EMIM]Cl/AlCl3 ionic liquids at a W working electrode. Scan rate: 100 mV s−1 . (a) Negative scan limit: −2.2 V; (b) negative scan limit: −2.7 V.
4.1.2.2 Electrodeposition of Sodium
The first work in this field was reported by Winnick et al. in 1995 [4]. In order
to design a sodium/iron(II) chloride battery, they examined a 1-ethyl-3-methylimidazolium chloride/aluminum chloride-based system. As described by Lipsztajn
and Osteryoung for lithium it was first necessary to synthesize the acidic ionic
liquid by adding an excess of AlCl3 and then adding an equivalent amount of
sodium chloride as a buffer to obtain again the neutral species.
4.1 Electrodeposition in AlCl3 -based Ionic Liquids
Fig. 4.3 Cyclic voltammogram of neutral buffered, unprotonated melt
at tungsten (a) and 303 stainless steel (b).
For their experiments they used tungsten and 303 stainless steel as substrates.
They used ionic liquids protonated with HCl in order to enhance the platingstripping efficiency. Electrodeposition of sodium from AlCl3 /[EMIM]Cl/NaCl ionic
liquid was not observed because of limited cathodic stability.
The addition of gaseous HCl increased the cathodic stability so that a Na+ /Na
redox was observed: at a pressure of 6.1 Torr a reduction of Na+ at −2.3 V followed
by a stripping wave at −2.1 V on the W electrode was observed.
Piersma et al. were able to enhance the deposition–stripping behavior by adding
triethanolamine hydrochloride instead of gaseous HCl. The addition widened the
electrochemical window and the resulting mixtures were reported to be stable for
a month.
87
88 4 Electrodeposition of Metals
Fig. 4.4 Cyclic voltammogram at tungsten of a neutral buffered, ionic
liquid protonated to a partial pressure of 6.1 Torr HCl.
Finally, Kohl et al. used quaternary ammonium chlorides, e.g. benzyltrimethylammonium chloride, together with AlCl3 and sodium chloride (working electrode:
Pt, counter electrode : Pt wire, Al wire in melt). The addition of SOCl2 was required
in order to reduce Na+ . At −2.4 V the deposition of sodium started and at −1.8 V
the reoxidation was observed [5].
4.1.3
Group II Metals
From the elements of this group, magnesium is probably the most interesting
to deposit on other metals because of its property of forming dense layers of the
corresponding oxides.
To date no method has been published that describes the electrodeposition of
magnesium or any other Group II metal from tetrachloroaluminate-based ionic
liquids.
4.1.4
Group III Metals
4.1.4.1 Electrodeposition of Aluminum and Aluminum Alloys
The electrodeposition of aluminum has enormous potential in industrial applications. The main reason for this is that aluminum reacts with oxygen to form dense
layers of aluminum oxides, protecting metals from corrosion. By far most of the
publications concerning the electrodeposition of metals from tetrachloroaluminatebased ionic liquids focus on aluminum.
In this context, Osteryoung and Robinson were the first, in 1980, when they
described the electrodeposition of aluminum on platinum and glassy carbon from
4.1 Electrodeposition in AlCl3 -based Ionic Liquids
an acidic composition of butylpyridinium chloride and AlCl3 where 50% benzene
was added [6]. The reduction was observed at −0.43 V. From basic reaction mixtures
they were not able to observe a deposition.
Osteryoung and Welch demonstrated by coulometry using a tungsten electrode that the deposition process of aluminum is reversible (in a butylpyridinium
chloride/AlCl3 ionic liquid) [7]. The deposition occurred at −0.43 V and the oxidation was observed at −0.22 V. In addition, they determined the rate of the corrosion
process to be 1 × 10−11 mol cm−2 s−1 . In a bulk deposition they were able to deposit
aluminum on a brass substrate at a thickness up to 15 µm.
Lay and Skyllas-Kazacos were the first to describe a deposition from imidazoliumbased tetrachloroaluminate ionic liquid [8]. On glassy carbon, aluminum was deposited at −0.2 V (instead of −0.43 V for the pyridinium-based system of Osteryoung and Welch). Furthermore, they were able to show that the deposition process
has complicated kinetics and is not simply controlled by diffusion. Using a tungsten
electrode they were able to demonstrate in chronopotentiometric measurements
that initially a potential of −0.65 V is necessary due to the nucleation process, but
after reaching the barrier the potential drops below −0.2 V.
Hussey et al. carried out an aluminum bulk deposition on copper foil using a
Lewis acidic aluminum chloride 1-ethyl-3-methyl-imidazolium chloride-based ionic
liquid [9]. The thickness of the observed deposits were in the range 24–30 µm. Without additives the deposits were not shiny and only poorly adherent. The addition of
benzene enhanced the quality of the deposit. XRD measurements confirmed that
the composition of the deposits was 100% aluminum metal.
Very significant investigations concerning important parameters for the commercialization were performed by Abbott et al. They used benzyltrimethylammonium chloride/AlCl3 instead of 1-ethyl-3-methyl-imidazolium chloride/AlCl3
to deposit aluminum on a number of substrates [10]. The reason for using benzyltrimethylammonium chloride was that this material is less water sensitive, easier
to purify, has greater thermal stability and is potentially more cost effective than
the materials used before. Surprisingly, when using an iron electrode an underpotential deposition of aluminum at +0.20 V was observed, which was not the case
on aluminum (−0.20 V) and platinum (−0.25 V) substrates. In the corresponding
cyclovoltammogram a nucleation loop was observed for Al and Pt, which suggests a
kinetic nucleation control of the deposition. A test in a hull cell was also performed
on a nickel foil. It showed that the brightest and most uniform deposit was obtained
at 5.1 × 10−5 A cm−2 .
Endres et al. were able to deposit nanocrystalline aluminum from an aluminum
chloride/1-butyl-3-methyl-imidazolium chloride-based ionic liquid (molar ratio:
55/45 mol%) and to characterize it by using XRD and TEM [11]. Figure 4.5 shows
the corresponding XRD pattern.
The size distribution is shown in Figure 4.6. The grain size was determined to
be 12 ± 1 nm.
Furthermore, Endres et al. were also able to deposit aluminum alloys such
as Al–Mn, which are widely used in the automotive and aviation industries for
lightweight construction. The deposition was performed from a Lewis acidic ionic
89
90 4 Electrodeposition of Metals
Fig. 4.5 XRD pattern of nanocrystalline Al with a grain size of 12 ± 1 nm.
liquid (as described above), where MnCl2 was added. The average grain size of the
deposits was 26 ± 1 nm.
4.1.4.2 Electrodeposition of Indium
Sun et al. reported the electrodeposition of indium on glassy carbon, tungsten
and Nickel. In basic chloroaluminates, elemental indium is formed via one threeelectron reduction step from the [InCl5 ]2− complex [12]. Furthermore, Carpenter
reported the deposition of an indium(I) species [13].
Fig. 4.6 Size distribution of nanocrystalline Al TEM image.
4.1 Electrodeposition in AlCl3 -based Ionic Liquids
4.1.4.3 Electrodeposition of Gallium
The electrodeposition of gallium is of interest for its extraction and purification
and for the production of III-V semiconductors. Sun et al. were the first to report
the electrodeposition of gallium from Lewis acidic aluminum chloride–1-ethyl3-methyl-imidazolium chloride melts (ratio 60:40 mol%) on tungsten and glassy
carbon in 1999 [14] The Ga(I) species was introduced by anodization of the gallium
metal.
Using a tungsten electrode the electroreduction was observed at +0.255 V. At
a deposition temperature of 30 ◦ C the gallium deposits were liquid, covering the
tungsten wire. If removed from the electrolyte they solidified in droplet-like form.
The Ga(I)/Ga(III) electrode reaction exhibits slow charge transfer kinetics with an
anodic transfer coefficient and a standard heterogeneous rate constant of 0.24 and
3.16 × 10−4 cm s−1 , respectively, at tungsten.
Using glassy carbon they observed a three-dimensional nucleation of gallium,
with diffusion-controlled growth of the nuclei. The diffusion coefficients for the
Ga(III) and the Ga(I) species were 2.28 × 10−7 and 9.12 × 10−7 cm2 s−1 , respectively.
4.1.5
Group IV Metals
4.1.5.1 Electrodeposition of Tin
Pitner and Hussey studied the electrochemistry of tin in acidic and basic AlCl3 /1ethyl-3-methyl-imidazolium chloride-based ionic liquids by using voltammetry and
chronoamperometry at 40 ◦ C [15]. They reported that the Sn(II) reduction process
is uncomplicated at a platinum substrate, where in the acidic ionic liquid the
reduction wave was observed at +0.46 V on the Pt electrode and the oxidation
at +0.56 V. When they used a gold electrode instead of platinum, they observed
an underpotential deposition of a tin monolayer and an additional underpotential
deposition process that was attributed to the formation of tin–gold alloy at the
surface. The deposition of tin on glassy carbon was controlled by nucleation.
The formal potentials of the Sn(II)/Sn couple in the 66.7/33.3 and
44.4/55.6 mol% ionic liquids were determined to be 0.55 ± 0.01 V and −0.85 ±
0.03 V, respectively, vs. Al(III)/Al in the 66.7/33.3 mol% ionic liquid, and the diffusion coefficients of Sn(II) were determined to be (5.3 ± 0.7) × 10−7 and (5.1 ± 0.6)
× 10−7 cm2 s−1 , respectively.
4.1.6
Group V Metals
4.1.6.1 Electrodeposition of Antimony
Habboush and Osteryoung were the first to describe the electrodeposition of a
Group V metal from AlCl3 /1-butyl-pyridinium chloride-based ionic liquids. As antimony sources they used SbCl3 or Sb-rods, dissolved by anodic dissolution [16]. For
the composition AlCl3 :BuPyCl (0.8:1) a deposition of Sb was observed at −0.885 V
91
92 4 Electrodeposition of Metals
and dissolution at −0.420 V and for the solution composed of AlCl3 :BuPyCl (1.5:1)
a deposition of Sb occurred at +0.53 V and anodic dissolution at +1.11 V.
They remarked that SbCl2 + was the dominant species in the acidic ionic liquids.
The reduction of this species on glassy carbon exhibited irreversible behavior. In the
basic melts SbCl4 − and SbCl6 − were believed to be the dominant species. In basic
media the reduction of Sb(III) to the metal on glassy carbon was also irreversible
while its oxidation to Sb(V) showed quasi-reversible behavior.
In another publication Lipsztain and Osteryoung used imidazolium-based ionic
liquids to study the behavior of Sb(III) under conditions where the unbuffered
properties of a neutral ionic liquid played an important role [17].
4.1.7
Group VI Metals
4.1.7.1 Electrodeposition of Tellurium
Sun et al. used basic melts of 1-ethyl-3-methyl-imidazolium chloride and aluminum
chloride of different molar ratios to dissolve TeCl4 [18]. At −0.68 V reduction of
tellurium was observed, which was clearly controlled by the nucleation/growth rate.
The bulk deposition led only to poorly adherent powder which was confirmed to be
Te◦ by XRD.
4.2
Electrodeposition of Metals in Air- and Water-stable Ionic Liquids
4.2.1
Introduction
Ionic liquids have attracted extensive attention since they have extraordinary physical properties, superior to those of water or organic solvents. They are usually
nonvolatile, nonflammable, less toxic, good solvents for both organics and inorganics and can be used over a wide temperature range (see Chapters 2 and 3). Moreover,
ionic liquids have quite large electrochemical windows, up to 6 V, and hence they
give access to elements which cannot be electrodeposited from aqueous or organic
solutions. Another advantage of ionic liquids is that problems associated with hydrogen ions in conventional protic solvents, for example hydrogen embrittlement,
can be eliminated as most of them are aprotic.
Chloroaluminate ionic liquids are regarded as the first generation of ionic liquids. However, their hygroscopic nature has delayed progress in many applications
since they must be prepared and handled under an inert gas atmosphere. Thus,
the synthesis of air- and water-stable ionic liquids, which are considered as the
second and third generations of ionic liquids, has attracted further interest in the
use of ionic liquids in various fields. Unlike the chloroaluminate ionic liquids,
these ionic liquids can be prepared and safely stored outside an inert atmosphere.
Generally, they are water insensitive. However, water-containing [BMIM]PF6 can
4.2 Electrodeposition of Metals in Air- and Water-stable Ionic Liquids 93
be pretty aggressive due to the formation of HF by decomposition of the ionic liquid
in the presence of water. Therefore, ionic liquids based on more hydrophobic anions such as trifluoromethanesulfonate (CF3 SO3 − ), bis (trifluoromethanesulfonyl)
amide [(CF3 SO2 )2 N− ] and tris (trifluoromethanesulfonyl) methide [(CF3 SO2 )3 C− ]
have been developed [19–21]. These ionic liquids have received extensive attention,
not only because of their low reactivity towards water but also because of their
large electrochemical windows. In general, the wide electrochemical windows of
the ionic liquids have opened the door to the electrodeposition of reactive elements,
such as Al, Ta, Si, Se and others, which cannot be obtained from aqueous solutions
at moderate temperatures. For more information on the use of ionic liquids as
solvents for electrodeposition of metals and semiconductors, we refer to recently
published review articles [22–25]. We have reviewed the electrodeposition of metals
and semiconductors in the most popular air- and water-stable ionic liquids in a
short review [23]. In this section we present a review of the recent efforts for the
electrodeposition of less reactive metals (such as Zn, Cu, Cd, Cr, Pd, Ag, Pt and
Sb), highly reactive metals (such as Al, Mg and Li) and, finally, refractory metals
(such as Ta and Ti) in air- and water-stable ionic liquids.
4.2.2
Electrodeposition of Less Reactive Metals
Most of the metals that can be electrodeposited from aqueous solutions can also
be electrodeposited from ionic liquids. As many ionic liquids are environmentally
friendly, they are considered as suitable alternatives for poisonous plating baths.
Furthermore, as ionic liquids have very low vapor pressures (often between 10−11
and 10−10 mbar at or near room temperature, at 100 ◦ C the vapor pressure depends on the liquid, in the region of 10−6 –10−4 mbar) they could principally be
used in open galvanic baths at variable temperatures without releasing harmful
vapors which reduces the amount of volatile organic compounds released into the
atmosphere. Another advantage of using ionic liquids instead of aqueous baths is
that their thermal stability makes it easier to electrodeposit metals through direct
electrodeposition without subsequent annealing. Moreover, electrodeposition carried out in aqueous solutions is often complicated by problems involving hydrogen
embrittlement and low current efficiency. As a result, ionic liquids wherein the coevolution of hydrogen is excluded are good alternatives to aqueous plating baths.
Examples of some important metals that can be electrodeposited from aqueous
electrolytes and ionic liquids will be presented below.
4.2.2.1 Zinc
Zinc and its alloys are good materials for corrosion-resistant coatings and they are
widely used in the automobile industry. The electrodeposition of zinc or its alloys is
normally performed in aqueous electrolyte solutions. However, zinc and its alloys
can be obtained in improved quality from ionic liquids. It was shown that Lewis
acidic ZnCl2 –[EMIM]Cl (1-ethyl-3-methylimidazolium chloride) liquids in which
94 4 Electrodeposition of Metals
the concentration of ZnCl2 is higher than 33 mol%, are potentially useful for the
electrodeposition of zinc and zinc-containing alloys [26–28], see also Chapters 4.3
and 5. The cathodic electrochemical window of these liquids is determined by the
reduction of Zn(II) to Zn metal. As a result, the electrodeposition of Zn and its
alloys is possible from these melts. As these ionic liquids are aprotic solvents, hydrogen embrittlement is excluded. Such systems and others under development
may allow deposition of zinc on safety-relevant steel supports, with high-strength
steel qualities, which are not allowed to be zincated in aqueous solutions. Moreover, in contrast to the AlCl3 -based ionic liquids (see Chapters 2.1 and 4.1), the
ZnCl2 –[EMIM]Cl liquid does not react vigorously with moisture and is thus easier
to handle. Recently Abbott et al. [29] reported that Zn can be electrodeposited from
a solution of ZnCl2 in urea and ethylene glycol/choline chloride-based ionic liquid
(see Chapter 4.3).
4.2.2.2 Copper
Copper is a widely used metal with extensive industrial applications, especially in
the semiconductor industry. Almost all connections on semiconductor chips are
made with copper due to its low electrical resistance, good mechanical properties
and high corrosion resistance. Therefore, the electrochemical deposition of Cu
becomes more attractive in the semiconductor manufacturing industry because
of its low processing temperature, high selectivity and low cost. One problem in
semiconductor technology is that the tantalum diffusion barrier, on which copper is deposited, reacts to tantalum oxide at the surface. The electrodeposition
of copper has been intensively investigated in chloroaluminate ionic liquids (see
for example Ref. [30]). Sun et al. [31] demonstrated that Cu can be electrodeposited in a basic chloride containing 1-ethyl-3-methyl imidazolium tetrafluoroborate ([EMIM]BF4 ). Murase et al. [32] stated that Cu can be electrodeposited in the
air- and water-stable ionic liquid trimethyl-n-hexylammonium bis (trifluoromethylsulfonyl) amide and the Cu deposition and dissolution involved one-electron redox
reactions.
Recently, we reported that nanocrystalline copper with an average crystallite
size of about 50 nm can be obtained without additives in the ionic liquid 1-butyl1-methylpyrrolidinium bis (trifluoromethylsulfonyl) amide ([BMP]Tf2 N) [33]. Because of the limited solubility of many tested copper compounds in the ionic liquid
[BMP]Tf2 N, copper cations were introduced into the ionic liquid by anodic dissolution of a copper electrode [33]. The SEM micrograph in Figure 4.7 shows the
surface morphology of such an electrodeposited copper layer on a gold substrate
obtained at a constant potential of −0.250 V (vs. Pt) for 2 h in the ionic liquid
[BMP]Tf2 N containing 60 mmol L−1 of Cu(I) at 25 ◦ C. As seen, the deposit is dense
and contains fine crystallites with average sizes of about 50 nm. Interestingly, the
deposited copper is nanocrystalline without any additive. The electrodeposition of
nanocrystalline copper is quite interesting as a coating since nano-Cu has excellent
mechanical and electronic properties, superior to those of microcrystalline copper,
furthermore nano-Cu is an important catalyst [34].
4.2 Electrodeposition of Metals in Air- and Water-stable Ionic Liquids 95
Fig. 4.7 SEM micrograph of nanocrystalline copper obtained potentiostatically on Au in the ionic liquid [BMP]Tf2 N containing 60 mmol L−1
Cu(I) at a potential of −0.25 V (vs. Pt) for 2 h at room temperature.
4.2.2.3 Cadmium
Despite environmental and health issues, cadmium is still an important metal
because of the wide variety of its applications, e.g. in solar cells [35] and rechargeable
batteries [36]. It was reported [37] that Cd can be electrodeposited from a basic
1-ethyl-3-methyl imidazolium tetrafluoroborate [EMIM]BF4 ionic liquid containing
CdCl2 . The cadmium electrodeposits were found to be very pure and adhered well
to the tungsten substrate. Furthermore, Cd can also be deposited in the acidic zinc
chloride-1-ethyl-3-methyimidazolium chloride (ZnCl2 –[EMIM]Cl) ionic liquid [38].
At a more negative deposition potential, zinc can be codeposited with cadmium.
The Cd content in the Cd–Zn electrodeposits can be increased by increasing the
Cd(II) concentration or by increasing the deposition temperature [38].
4.2.2.4 Chromium
The electrodeposition of chromium in a mixture of choline chloride and
chromium(III) chloride hexahydrate has been reported recently [39]. A dark green,
viscous liquid is obtained by mixing choline chloride with chromium(III) chloride
hexahydrate and the physical properties of this deep eutectic solvent are characteristic of an ionic liquid. The eutectic composition is found to be 1:2 choline
chloride/chromium chloride. From this ionic liquid chromium can be electrodeposited efficiently to yield a crack-free deposit [39]. Addition of LiCl to the choline
chloride–CrCl3 ·6H2 O liquid was found to allow the deposition of nanocrystalline
black chromium films [40]. The use of this ionic liquid might offer an environmentally friendly process for electrodeposition of chromium instead of the current
chromic acid-based baths. However, some efforts are still necessary to get shining
96 4 Electrodeposition of Metals
chromium deposits which can compete with the conventional Cr(VI)- or Cr(III)based aqueous galvanic processes.
4.2.2.5 Palladium
Palladium is employed in a number of industrial applications and fundamental
studies because of its high catalytic activity for many chemical reactions, e.g. its
ability to absorb hydrogen [41]. On the other hand, due to hydrogen absorption,
only brittle Pd deposits can be obtained in aqueous solutions. The advantage of
performing electrodeposition of Pd in ionic liquids is that hydrogen evolution
does not occur. Sun et al. demonstrated that Pd and some of its alloys, namely
Pd–Ag [42], Pd–Au [43] and Pd–In [44], can be obtained from the basic 1-ethyl-3methylimidazolium chloride/tetrafluoroborate ionic liquid. Compact alloy deposits
were obtained and the Pd content in the deposits increased with the increase in Pd
mole fraction in the plating bath.
4.2.2.6 Silver
The electrodeposition of silver from chloroaluminate ionic liquids has been studied
by several authors [45–47]. Katayama et al. [48] reported that the room-temperature
ionic liquid 1-ethyl-3-methylimidazolium tetrafluoroborate ([EMIM]BF4 ) is applicable as an alternative electroplating bath for silver. The ionic liquid [EMIM]BF4 is
superior to the chloroaluminate systems since the electrodeposition of silver can
be performed without contamination of aluminum. Electrodeposition of silver in
the ionic liquids 1-butyl-3-methylimidazolium tetrafluoroborate ([BMIM]BF4 ) and
1-butyl-3-methylimidazolium hexafluorophosphate was also reported [49]. Recently
we showed that isolated silver nanoparticles can be deposited on the surface of the
ionic liquid 1-butyl-3-methylimidazolium trifluoromethylsulfonate ([BMIM]TfO)
by electrochemical reduction with free electrons from low-temperature plasma [50]
(see Chapter 10). This unusual reaction represents a novel electrochemical process,
leading to the reproducible growth of nanoscale materials. In our experience silver
is quite easy to deposit in many air- and water-stable ionic liquids.
4.2.2.7 Platinum
Films or nanoparticles of platinum are of particular interest since they are important catalysts for many chemical reactions. The electrodeposition of platinum
in the ionic liquids [BMIM]BF4 and [BMIM]PF6 has been reported [51]. The Pt
deposit was shiny, dense and contained nanocrystals with sizes less than 100 nm.
Furthermore, the deposited Pt films obtained in the ionic liquids exhibited higher
catalytic performance for the electroxidation of methanol compared with the Pt
films obtained in HClO4 aqueous solution [51]. The electrodeposition of PtZn from
a Lewis acidic 40–60% ZnCl2 –[EMIM]Cl containing PtCl2 has also been reported
[52]. Similar to the electrodeposition of palladium there is no hydrogen evolution
during platinum deposition in ionic liquids which can alter the quality of platinum
deposited in aqueous solutions.
4.2 Electrodeposition of Metals in Air- and Water-stable Ionic Liquids 97
4.2.2.8 Antimony
Antimony is a brittle silvery-white metal. Although the unalloyed form of antimony
is not often used in industry, alloys of antimony have found wide commercial applications. The integration of antimony gives certain desirable properties, such as
increased corrosion resistance and hardness. Moreover, antimony is also the component of some semiconductors such as InSb and InAs1–x Sbx . Sb electrodeposits
with good adherence were obtained in a water-stable 1-ethyl-3-methylimidazolium
chloride-tetrafluoroborate ([EMIM]Cl-BF4 ) room-temperature ionic liquid [53]. Furthermore, it was stated that a crystalline InSb compound can be obtained through
direct electrodeposition in the ionic liquid [EMIM]Cl-BF4 containing In(III) and
Sb(III) at 120 ◦ C [54]. It is just a question of time until antimony electrodeposition
is reported in the third generation of ionic liquids.
4.2.3
Electrodeposition of Reactive Metals
In this section we will show that air- and water-stable ionic liquids can be used for
the electrodeposition of highly reactive elements which cannot be obtained from
aqueous solutions, such as aluminum, magnesium and lithium, and also refractory
metals such as tantalum and titanium. Although these liquids are no longer airand water-stable when AlCl3 , TaF5 , TiCl4 and others are dissolved, quite interesting
results can be obtained in these liquids.
4.2.3.1 Electrodeposition of Aluminum
As is known, the commercial production of aluminum is carried out by electrolysis
of molten cryolite (Na3 AlF6 ) in which aluminum oxide is dissolved at an elevated
temperature of about 1000 ◦ C [55]. This method is still the main industrial method
for primary aluminum production. However, it is not suitable for coating other
metals with a layer of aluminum since the electrolysis is performed at a temperature where Al is liquid. Nowadays, there are various methods for aluminum coating
such as, hot dipping, thermal spraying, sputter deposition, vapor deposition and
electroplating in e.g. organic solvents. The electroplating process offers some advantages: the deposits are usually adherent and do not affect the structural and
mechanical properties of the substrate. Furthermore, the thickness and the quality of the deposits can be adjusted by controlling the experimental parameters.
Moreover, the electroplating process is rather cost-efficient, since it is performed
at moderate temperature.
Because of its high reactivity (−1.67 V vs. NHE), the electrodeposition of aluminum from aqueous solutions is not possible. Therefore, electrolytes for Al deposition must be aprotic, such as ionic liquids or organic solvents. The electrodeposition of aluminum in organic solutions is commercially available (SIGAL-process
[56, 57]) but due to volatility and flammability there are some safety issues. Therefore, the development of room-temperature ionic liquids in recent years has resulted
in another potential approach for aluminum electrodeposition. Many papers have
been published on the electrodeposition of aluminum from chloroaluminate (first
98 4 Electrodeposition of Metals
generation) ionic liquids [58–68]. Although high quality Al deposits can be obtained
using such liquids, a main disadvantage of them is that they are extremely hygroscopic and thus must be strictly handled under inert gas conditions. Furthermore,
the organic halides are very difficult to dry. Therefore, the electrodeposition of
aluminum in less reactive air- and water-stable ionic liquids is of great interest.
Quite recently we reported for the first time that nano- and microcrystalline
aluminum can be electrodeposited in three different air- and water-stable ionic
liquids, namely 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl) amide
[BMP]Tf2 N, 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl) amide
[EMIM]Tf2 N and trihexyl-tetradecyl phosphonium bis(trifluoromethylsulfonyl)
amide (P14,6,6,6 Tf2 N) [69, 70]. It was found that the ionic liquids [BMP] Tf2 N
and [EMIM] Tf2 N form biphasic mixtures in an AlCl3 concentration range
1.6–2.5 mol L−1 and 2.5–5 mol L−1 , respectively [70]. Moreover, the electrodeposition of aluminum at room temperature occurs only from the upper phase at
AlCl3 concentrations ≥ 1.6 mol L−1 and ≥ 5 mol L−1 in the ionic liquids [BMP]Tf2 N
and [EMIM] Tf2 N, respectively. The biphasic behavior of such liquids was first reported by Wasserscheid [71], but a comprehensive understanding of the aluminum
species in the phases is still missing. Interestingly, we have found that Al can only
be electrodeposited from the upper phase of the biphasic mixture. This means that
the reducible aluminum-containing species exists only in the upper phase of the
biphasic mixtures and hence the electrodeposition of Al occurs only from the upper
phase.
In the case of the ionic liquid [BMP]Tf2 N, shiny, dense and adherent deposits
with very fine crystallites in the nanometer regime can be obtained without any
addition of organic brighteners or use of pulse plating techniques (Figure 4.8(a)).
In contrast, coarse cubic-shaped aluminum particles in the micrometer regime are
obtained in the ionic liquid [EMIM]Tf2 N (Figure 4.8(b)). As the temperature and
electrochemical parameters were varied it is unlikely that this observation is due to
viscosity effects alone. Probably, the [BMP]+ cation acts as a grain refiner and plays
its role by adsorption on the substrates and on growing nuclei, thus hindering the
further growth of crystallites. Many more experiments are required to elucidate the
effect of the ionic liquid on the deposit quality.
4.2.3.2 Electrodeposition of Magnesium
Magnesium and its alloys offer a high potential for use as lightweight structural
materials in automotive and aircraft applications. As magnesium is a very reactive metal (E ◦ = −2.37 V vs. NHE), it can be only obtained from aprotic electrolytes. It is worth noting that the electrodeposition of magnesium in organic
electrolytes or in ionic liquids is feasible but not straightforward. Recently, it
was claimed in several papers (with similar data) that magnesium can be electrodeposited from the ionic liquid 1-butyl-3-methylimidazolium tetrafluoroborate
[BMIM]BF4 using magnesium trifluoromethylsulfonate [Mg(CF3 SO3 )2 ] as a source
of magnesium [72–75]. Apart from the comparatively low reduction stability of imidazolium ions with magnesium deposition, there is no hard evidence presented
for the deposition of metallic magnesium. We ourselves could not electrodeposit
4.2 Electrodeposition of Metals in Air- and Water-stable Ionic Liquids 99
Fig. 4.8 (a) SEM micrograph of an electrodeposited Al film on gold formed in the
upper phase of the mixture AlCl3 /[BMP]
Tf2 N after potentiostatic polarization at
−0.45 V (vs. Al) for 2 h at 100 ◦ C. (b) SEM
micrograph of an electrodeposited Al film
on gold made in the upper phase of the
mixture AlCl3 /[EMIm] Tf2 N after potentiostatic polarization at −0.05 V (vs. Al) for 2 h
at 100 ◦ C.
magnesium using the recipes described in the mentioned papers. In part we
had solubility problems with water-free ionic liquids, and a deposit forms which
seems to contain mainly decomposition products. On the other hand the reduction stability of ionic liquids with tetra-alkylammonium or pyrrolidinium cations
(∼ −3 V vs. NHE) should, thermodynamically, be sufficient to allow magnesium
deposition. In a recent paper, we have tried to electrodeposit magnesium in
the ionic liquids 1-butyl-1-methyl-pyrrolidinium bis(trifluoromethylsulfonyl) amide
and 1-butyl-1-methylpyrrolidinium trifluoromethylsulfonate ([BMP]TfO) using a
Grignard reagent and magnesium perchlorate as sources of magnesium, respectively [76]. Pyrrolidinium ions are cathodically about 700 mV more stable than
100 4 Electrodeposition of Metals
imidazolium ions; furthermore, it is well known that magnesium can be electrodeposited from Grignard compounds in ether solvents [77, 78]. It was found that the
electroreduction of Grignard reagent in the ionic liquid [BMP]Tf2 N might lead to
the formation of thin Mg films which, under air, are subject to oxidation to magnesium oxide and hydroxide. Furthermore, the reduction of Mg(ClO4 )2 in [BMP]TfO
is followed by an anodic process showing typical stripping peak behavior; however,
the current efficiency for magnesium deposition is not very high. The electrodeposition of magnesium in ionic liquids should, thermodynamically, be possible.
Nevertheless more effort is required to find a suitable ionic liquid and suitable
magnesium precursors for a technically relevant process.
4.2.3.3 Electrodeposition of Lithium
Lithium metal is of particular importance as an anode material for high energy batteries. Lithium batteries are used widely in portable electronic devices and electric
vehicles. They show the highest energy density among the applicable chemical and
electrochemical energy storage systems (up to 180 Wh kg−1 ). Due to its high reactivity (E ◦ = −3.05 V vs. NHE), lithium cannot be electrodeposited from any aqueous
electrolytes. The electrodeposition of lithium in the ionic liquid [BMP]Tf2 N containing Li(Tf2 N) was reported by Katayama [79] and MacFarlane [80]. The latter
group investigated the cycling properties (repetitive deposition and stripping) of
lithium in the ionic liquid [BMP]Tf2 N containing Li(Tf2 N). It was shown that uniform lithium deposit morphology over many cycles can be achieved at moderate
current densities. Cycling efficiencies exceeding 99% were obtained [80]. However,
the Tf2 N ion breaks down in the presence of lithium. On the one hand, the decomposition products stabilize the Li, on the other hand anion decomposition leads to a
limited lifetime of such secondary batteries, therefore much more effort is required
to make Li secondary batteries based on ionic liquids.
4.2.3.4 Electrodeposition of Tantalum
High-temperature molten salts were found to be efficient baths for the electrodeposition of tantalum [81–86]. Senderoff and Mellors reported the first results on the
electrodeposition of Ta using the ternary eutectic mixture LiF–NaF–KF as a solvent
and K2 TaF7 as a source of Ta at temperatures between 650 and 850 ◦ C [81, 82].
Despite enormous importance, these baths have many technical and economic
problems, such as loss in the current efficiency of the electrolysis process due to
the dissolution of metal after its deposition [87] and the expected corrosion problems at high temperatures. Furthermore, from a practical point of view, molten
salts are hardly suited for the coating of sensitive materials like NiTi shape memory
alloy with tantalum since the electrolysis process is performed at too high temperatures. With ionic liquids a technical electroplating processes might be performed
at moderate temperature.
Recently, we reported for the first time that tantalum can be electrodeposited as
thin layers in the water and air stable ionic liquid 1-butyl-1-methyl pyrrolidinium
bis (trifluoromethylsulfonyl) amide at 200 ◦ C using TaF5 as a source of tantalum
[88]. The quality of the deposit was found to be improved on addition of LiF to the
4.2 Electrodeposition of Metals in Air- and Water-stable Ionic Liquids 101
Fig. 4.9 (a) SEM micrograph of the electrodeposit formed potentiostatically on Pt in ([BMP]Tf2 N) containing 0.25 M TaF5 and 0.25 M LiF
at a potential of −1.8 V for 1 h at 200 ◦ C. (b) XRD patterns of the deposited layer.
deposition bath. The SEM micrograph of the Ta electrodeposit, Figure 4.9(a), made
potentiostatically at −1.8 V in ([BMP]Tf2 N) containing 0.25 M TaF5 and 0.25 M LiF
on a Pt electrode at 200 ◦ C for 1 h shows a smooth, coherent and dense layer. XRD
patterns of the electrodeposit clearly show the characteristic patterns of crystalline
tantalum, Figure 4.9(b). We would like to point out clearly that tantalum deposition
is not straightforward in ionic liquids, especially at low temperatures. Under the
wrong conditions (high current density) mainly subvalent XRD-amorphous tantalum species like [Ta6 Cl12 ]2+ are obtained. There seems to be a limiting current
density above which one obtains only subhalides and below which thin crystalline
tantalum layers (Figure 4.9(a)) are obtained. Currently we are able to deposit about
1 µm thick tantalum layers. In our opinion thicker deposits are feasible, but the
102 4 Electrodeposition of Metals
development of a technical process will require some effort. Some problems in the
electrodeposition of refractory and rare earth elements will be presented in Section
4.4.
Furthermore, we showed that adherent, dense and uniform layers of Ta can be
electrodeposited on NiTi alloy in the ionic liquid [BMP]Tf2 N containing 0.25 M TaF5
and 0.25 M LiF at 200 ◦ C [89]. NiTi alloys are widely used as orthodontic wires, selfexpanding cardiovascular and urological stents, and bone fracture fixation plates
and nails [90–92]. The biocompatibility of NiTi implants depends on their corrosion
resistance. The major risk associated with NiTi implants is the breakdown of the
passive film which occurs owing to the aggressiveness of human body fluids,
leading to a release of Ni ions that may cause allergic, toxic and carcinogenic
effects [93–95]. We found that the electrodeposition of only a 500 nm thick film
of Ta on NiTi alloy improves its corrosion resistance considerably, leading to a
decreased release of Ni ions into solution which enhances its biocompatibility [89].
Furthermore, we think that micrometer thick Ta layers, e.g. on stents to improve
the X-ray contrast, are possible.
4.2.3.5 Electrodeposition of Titanium
Titanium owes its great importance to its excellent mechanical and corrosion
performance. As for most refractory metals, high-temperature molten salts are
considered as the most efficient baths for titanium electrodeposition. Recently,
there was an attempt to electrodeposit titanium at room temperature in the airand water-stable ionic liquid 1-butyl-3-methyl-imidazolium bis (trifluoromethylsulfonyl) amide [BMIM] Tf2 N containing TiCl4 as a source of titanium. Using in situ
STM there were hints that titanium may be electrodeposited in ultrathin layers [96].
Our own experience has shown that attempts to deposit micrometer thick titanium
deposits with the recipe in Ref. [96] fail. Instead of elemental titanium soluble, polymeric subvalent titanium halide species are obtained. In situ electrochemical quartz
crystal microbalance (EQCM) measurements show that there is a tremendous increase in viscosity during TiCl4 electroreduction, furthermore Tf2 N breakdown (see
Refs. [97] and [98]) might alter titanium deposition. Thermodynamically, Ti deposition should be possible in thick layers in ionic liquids, but the right ionic liquid
and especially the right titanium precursors still have to be found. An idea might
be to make Ti(Tf2 N)4 or similar compounds for titanium electrodeposition.
4.2.4
Summary
In this chapter we have briefly discussed the high potential of air- and waterstable ionic liquids as electrolytes for metal deposition. Their extraordinary physical
properties, superior to those of water or organic solvents, and their stability, open
the door to the electrodeposition of many metals. Some advantages of air- and waterstable ionic liquids in electrodeposition are that they are quite easy to purify and
handle and in most cases they do not decompose under environmental conditions.
They can have pretty wide electrochemical windows of up to 6 V, and hence they
4.3 Deposition of Metals from Non-chloroaluminate Eutectic Mixtures 103
give access to reactive metals which cannot be electrodeposited from aqueous or
organic solutions. This branch of electrodeposition is quite novel and will require
much more effort to develop technical processes. However, there is also a price to
pay. In our experience only rarely can the know-how from aqueous electrochemistry
be transferred to ionic liquids. A quick success, especially with refractory and rare
earth metals, is unlikely as cluster chemistry has to be considered.
4.3
Deposition of Metals from Non-chloroaluminate Eutectic Mixtures
The preceding chapters have shown that the majority of metals can now be electrodeposited from ambient-temperature ionic liquids. However, this does not necessarily mean that the liquid with the widest potential window will negate the use
of all other ionic liquids. Rather, it is most likely that ionic liquids will be taskspecific with discrete anions being used for metals that cannot be electrodeposited
from aqueous solutions such as Al, Li, Ti, V and W. Type I eutectics will probably be the most suitable for Al, Ga and Ge. Type II eutectics are most suitable
for Cr and Type III are most suited to Zn, Cu, Ag and associated alloys. Type III
will also find application in metal winning, oxide recycling and electropolishing.
To date most practically important metals have been electrodeposited from ionic
liquids and a comprehensive review is given in articles by Abbott [99] and Endres
[100–102].
In this chapter we will concentrate on the deposition of metals from eutecticbased ionic liquids. These have been developed since the end of the 1990s, primarily
by our group and that of Sun. Figure 4.10 shows just some of the metals that
Fig. 4.10 A range of metal and metal alloy coatings deposited electrolytically from type II (Cr) and type III (Ni, Cu Zn Sn, Ag) choline
chloride-based ionic liquids.
104 4 Electrodeposition of Metals
have been deposited from these types of eutectics. The principles underlying the
eutectic-based liquids are all the same although clearly speciation is dominated
by the Lewis or Brønsted acidity of the components. The electrochemistry of the
liquids is largely unaffected by the cation although this does have a significant
effect upon the physical properties, most noticeably the phase behavior, viscosity
and conductivity.
A significant number of studies have characterized the physical properties of
eutectic-based ionic liquids but these have tended to focus on bulk properties such
as viscosity, conductivity, density and phase behavior. These are all covered in
Chapter 2.3. Some data are now emerging on speciation but little information is
available on local properties such as double layer structure or adsorption. Deposition mechanisms are also relatively rare as are studies on diffusion. Hence the
differences between metal deposition in aqueous and ionic liquids are difficult to
analyse because of our lack of understanding about processes occurring close to
the electrode/liquid interface.
One issue is that most metal complexes formed in ionic liquids are anionic and
these will have a significant effect on viscosity and mass transport. The effect of
metal ion concentration on reduction current will therefore not be linear. Relative
Lewis acidity will affect mass transport, ionic strength and speciation and accordingly the nucleation and growth mechanism of metals would be expected to be
concentration dependent.
The only models that exist for double layer structure in ionic liquids suggest
a Helmholz layer at the electrode/solution interface [103, 104]. If the reduction
potential is below the point of zero charge (pzc) then this would result in a layer
of cations approximately 5 Å thick across which most of the potential would be
dropped, making it difficult to reduce an anionic metal complex. Hence, the double
layer models must be incorrect.
Electrodeposition using eutectic-based ionic liquids has almost exclusively used
quaternary ammonium halides with metal halides primarily in the chloride form.
Aqueous plating solutions rarely use chlorides as they tend to yield black powdery
deposits and the inclusion of halides into metallic coatings is seen as undesirable
due to the possibility that it can lead to the breakdown of passivating layers and
exacerbate corrosion. The morphology issue is thought to be due to the ease of
nucleation from halide salts which leads to large numbers of small nuclei forming
at the electrode surface. Lewis basic anions cannot be circumvented for eutecticbased ionic liquids as they need to be good ligands to interact strongly with the
Lewis acid. The question that needs to be posed is whether chloride ions actually
cause a problem when their activity is negligible due to the presence of a strong
Lewis acid. The issue that needs to be addressed is that Type I and II eutecticbased ionic liquids necessarily have high concentrations of metal chlorides and will
tend to promote nucleus formation. In many cases the working concentration is
5 to 10 mol dm−3 which, although seemingly high, is not overly different to many
aqueous plating solutions. Further ionic liquid formulation needs to address how
nucleation can be suppressed while growth is supported.
4.3 Deposition of Metals from Non-chloroaluminate Eutectic Mixtures 105
4.3.1
Type I Eutectics
4.3.1.1 Chlorozincate Ionic Liquids
In general the potential windows are not as wide as those for the haloaluminates
or the discrete anions and they tend to be limited by the deposition of metal at the
cathodic limit and the evolution of chlorine at the anodic limit. Since ionic liquids
are aprotic solvents, hydrogen evolution and hydrogen embrittlement that often
occur in aqueous baths are circumvented in these liquids. Moreover, because of
their thermal stability, these ionic liquids make it easier to electrodeposit crystalline
metals and semiconductors through direct electrodeposition without subsequent
annealing.
From a practical perspective the chlorozincate liquids are easier to make and
handle than the corresponding chloroaluminates as they are less susceptible to
hydrolysis. As with the aluminum-based liquids the electrochemistry is dominated
by the complex anions present in the liquid, which depend upon the composition
and the relative Lewis acidity. There is some evidence, however, that hydrolysis
of zinc Lewis basic melts does occur as Hsiu et al. used fast atom bombardment mass spectroscopy (FAB MS) to show that some zinc species do contain
oxygen [105].
The same group studied the potential limits of [EMIM]Cl/ZnCl2 in the molar ratio range 3:1 to 1:3 [105]. It was found that in the Lewis acidic region (excess ZnCl2 )
the potential window was ca. 2 V; the negative potential limit is due to the deposition
of metallic zinc and the positive potential limit is due to the oxidation of the chlorozincate complexes to form chlorine. In the Lewis basic region the potential window
could be as large as 3 V, corresponding to the cathodic reduction of [EMIM]+
and the anodic oxidation of Cl− , which is similar to the basic chloroaluminate
melts. The potential windows for the Lewis acidic ionic liquids are surprisingly
close to the difference in the standard cell potentials for the corresponding half◦
◦
= 2.1 V). This suggests that
− E Cl
cell reactions in aqueous solutions (E Zn
2+
/Zn
/Cl−
2
while the reduction potential for Zn2 Cl5 − will be shifted with respect to the Zn2+ /Zn
couple the oxidation of Cl− to Cl2 will be affected by the same amount.
In the Lewis acidic melts underpotential deposition (UPD) of zinc was observed
on Pt and Ni electrodes. The potential window and UPD of zinc in Lewis acidic
choline chloride (ChCl):ZnCl2 was found to be exactly the same as the corresponding [EMIM]Cl system, suggesting that the cation has little or nothing to do with the
electrochemistry of the liquid.
An in depth study of the deposition mechanism was carried out by Sun et al. who
studied the 1:1 [EMIM]Cl/ZnCl2 system at various temperatures on glassy carbon
(GC), nickel and platinum electrodes [106]. The GC electrode required the largest
overpotential for deposition. The stripping process showed a single peak on GC,
whereas on Ni two oxidation processes were observed, separated by ca. 0.6 V. It was
proposed that the more positive oxidation process corresponded to the dissolution
of an intermetallic compound formed during electrodeposition.
106 4 Electrodeposition of Metals
Chronoamperometry on the GC and Ni electrodes at 50 ◦ C showed that the electrodeposition of zinc proceeded by a three-dimensional instantaneous nucleation
and growth process. The results also suggested that the growth process is under
mixed diffusion and kinetic control. The zinc deposits formed by bulk electrolysis
consisted of hexagonal grains with a size of 4–5 µm. These crystals were covered
by numerous small needles which could be subsequently removed, however the
underlying grains showed good adherence to the substrate. Deposits formed at
larger overpotentials were poorly adherent flakes and increasing the temperature
to 80 ◦ C also had little effect upon the morphology.
Sun also studied the effect of adding a diluent to a liquid to improve mass transport [106]. Propylene carbonate was added from 20 to 60% (v/v) and found to have
little effect on the voltammetry. Chronoamperometry showed the same instantaneous three-dimensional growth under mixed diffusion and kinetic control. The
propylene carbonate did, however, lead to an improvement in deposit morphology
compared to the neat melt and no small needles were seen in any of the deposits.
The grain size was affected by the deposition potential, indicating that nucleation
density increases with increased overpotential. The grain size was also affected by
diluent concentration with larger grains forming with higher propylene carbonate
content although this also increased the grain size distribution. Grain size could
also be increased by increasing the temperature of the melt.
Iwagishi studied the deposition of zinc from Lewis basic [EMIM]Br/ZnBr2 eutectics at 120 ◦ C and investigated the effect of adding ethylene glycol as a diluent
[107, 108]. Analysis of the choronoamperometric current–time transients indicated
that the overpotential was related to the progressive nucleation with diffusioncontrolled growth of the nuclei. The nucleation loop observed using cyclic voltammetry disappeared on adding more than 45 mol% ethylene glycol to the ionic liquid. The cathodic current increased with increasing ethylene glycol content and
it was proposed that this promoted the dissociation of [EMIM]Br to [EMIM]+
cation and Br− , and accordingly the concentration of ZnBr4 2− in the liquid was
increased.
The study was further extended to investigate the effect of a range of dihydric
alcohols (ethylene glycol, 1,3-propanediol, 1,2-butanediol and 1,3-butanediol). The
addition of the dihydric alcohol improved the smoothness and color of the deposits
and also increased the cathodic current efficiency at high current density. Of the
four dihydric alcohols, ethylene glycol gave the best results.
The same group studied the effect of water on zinc deposit morphology
in the same ionic liquid in the [EMIM]Br/ZnBr2 (70:30 mol%). Smooth layers
of silver-colored Zn were obtained at cathodic current densities < 100 A m−2 ,
whereas smooth grey Zn layers were electrodeposited at cathodic current densities
>150 A m−2 . The liquid with a water content of less than 10 ppm was superior to
the liquid with a water content of 400 ppm in cathodic current efficiency, smoothness, and metallic luster. The [EMIM]Br/ZnBr2 –ethylene glycol ternary systems
were also studied with a water content of 30 ppm. The cathodic current efficiencies
were all 100%, even at a current density as high as 300 A m−2 , in Lewis basic liquids
with an EG content of between 30 and 75 mol%.
4.3 Deposition of Metals from Non-chloroaluminate Eutectic Mixtures 107
Abbott et al. studied the deposition of zinc from a 1:2 choline chloride
(ChCl):ZnCl2 ionic liquid [109] at 60 ◦ C and found deposits with a similar morphology to that shown by Sun. The optimum current density was found to be
between 2 and 5 A m−2 and higher current densities led to powdery, non-adherent
deposits. This is due primarily to the high viscosity and low conductivity of the
choline-based liquids. The current plating efficiency in this liquid was found to
be effectively 100% and the deposition process was shown to be almost totally
reversible, with only the UPD material remaining on the surface.
Chlorozincate liquids have also been studied for the deposition of numerous
zinc-containing alloys including Pt–Zn, Zn–Fe, Sn–Zn and Cd–Zn alloys. These
alloys will all be discussed in greater detail in Chapter 5.
One explanation for the change in deposit morphology with time observed by Sun
and others [105–108] could be the structure of the double layer during deposition.
As the chlorozincate anions are reduced at the electrode surface the liquid close to
the electrode surface will become more Lewis basic
−
−
i.e. Zn2 Cl−
5 + 4 e → 2Zn + 5Cl
and this will in turn affect the composition of the zinc-containing species
−
−
Zn2 Cl−
5 + Cl → 2ZnCl3
The effect of this will be exacerbated if the liquid is viscous. It can be seen from the
above studies that the nucleation and growth mechanisms are dependent upon the
Lewis acidity of the liquid and this may help to explain the growth of needle-shaped
crystals on top of the hexagonal crystallites. The addition of a diluent decreases the
viscosity and could allow the chloride ions to diffuse away.
4.3.1.2 Other Type I Eutectics
Chlorostannate and chloroferrate [110] systems have been characterized but these
metals are of little use for electrodeposition and hence no concerted studies have
been made of their electrochemical properties. The electrochemical windows of the
Lewis acidic mixtures of FeCl3 and SnCl2 have been characterized with ChCl (both
in a 2:1 molar ratio) and it was found that the potential windows were similar to those
predicted from the standard aqueous reduction potentials [110]. The ferric chloride
system was studied by Katayama et al. for battery application [111]. The redox
reaction between divalent and trivalent iron species in binary and ternary molten
salt systems consisting of 1-ethyl-3-methylimidazolium chloride ([EMIM]Cl) with
iron chlorides, FeCl2 and FeCl3 , was investigated as possible half-cell reactions for
novel rechargeable redox batteries. A reversible one-electron redox reaction was
observed on a platinum electrode at 130 ◦ C.
Elemental gallium has been electrodeposited from chlorogallate ionic liquids
formed between [EMIM]Cl and GaCl3 [112]. The direct electrodeposition of GaAs
from ionic liquids was studied mainly by two groups. Wicelinski et al. [113] used
an acidic chloroaluminate liquid to co-deposit Ga and As. However, it was reported
108 4 Electrodeposition of Metals
that Al underpotential deposition on Ga occurred. Carpenter and Verbrugge studied
the deposition of GaAs using an ionic liquid based on GaCl3 to which AsCl3 was
added [112, 114]. Unfortunately poor quality deposits were obtained and both pure
As and Ga were present in the deposits but they did show that semiconductor
deposition was possible and thermal annealing could improve the quality of the
deposits.
The same principle has been used for the deposition of InSb alloys [115]. An ionic
liquid based on InCl3 is made and SbCl3 is added. InSb alloy was electrodeposited
but there was also some elemental In and Sb in the deposits. The In:Sb ratio could
be varied by altering the deposition potential.
4.3.2
Type II Eutectics
Type I eutectics only form liquids at ambient temperatures with metal salts that
melt below about 400 ◦ C (e.g. ZnCl2 /ChCl). This is related to the lattice energy of
the salt and its ability to interact with the quaternary ammonium salt. There is
a limited number of such salts and these preclude several of the technologically
important metals being incorporated in eutectic-based ionic liquids. In general, it
is metal salts with tetrahedral geometries that have lower lattice energies. One way
of decreasing the lattice energy, especially of octahedrally coordinated metal salts,
is to use hydrate salts, as these have relatively low melting points.
A wide variety of hydrated salt mixtures with choline chloride have been found
to form these ionic liquids, including CrCl3 ·6H2 O, CaCl2 ·6H2 O, LaCl3 ·6H2 O
CoCl2 ·6H2 O, LiNO3 ·4H2 O and Zn(NO3 )2 ·4H2 O [116] and hence this technology
could be generic to the deposition of a range of metals and alloys. However,
only chromium and cobalt have been deposited from these liquids. These liquids could be viewed as concentrated aqueous solutions, but because the ionic
strength is extremely high the water molecules are strongly coordinated to the
ions and hence they are difficult to reduce at the electrode surface. Accordingly
the deposition of metals such as chromium can be carried out with high current
efficiencies [117].
Figure 4.11 shows the voltammetry of a 1ChCl: 2 CrCl3 ·6H2 O and the corresponding 1ChCl: 2 CoCl2 ·6H2 O. It is immediately apparent that the two metals
have different electrochemical behavior. Chromium is reduced via a Cr(II) state
which forms an insoluble intermediate at the electrode–solution interface and
the deposition process is irreversible. The cobalt analogue is quasi-reversible as
there is a ca. 500 mV difference between the deposition potential and the stripping potential. Metal can be deposited from both eutectic mixtures. In the case of
chromium a dull, metallic-looking coating is obtained from the bulk electrolysis
of 1ChCl: 2 CrCl3 ·6H2 O at − 4 V and 60 ◦ C [116]. The corresponding experiment
using CoCl2 ·6H2 O yields a bright adherent metallic film.
The addition of up to 10 wt.% LiCl, however, leads to a matt black chromium
layer. The film has an amorphous morphology that is free of surface cracks, unlike
4.3 Deposition of Metals from Non-chloroaluminate Eutectic Mixtures 109
Fig. 4.11 Cyclic voltammetry of a 1ChCl: 2 CrCl3 .6H2 O and the corresponding 1ChCl: 2 CoCl2 ·6H2 O ionic liquids on a Pt microelectrode at
a sweep rate of 20 mV s−1 . Cr data at 60 ◦ C and Co data at 25 ◦ C.
samples deposited using aqueous Cr(III) and Cr(VI) which are highly crystalline
and have a highly cracked surface [116]. Cross-sectional analysis of the film produced using an ionic liquid showed that it was homogeneous with no structural
characteristics, even under the highest magnification. Relatively fast deposition
rates could be obtained (up to 60 µm h−1 ) although deposits thicker than about
30 µm tended to become quite powdery and less adherent. Chromium films were
deposited onto 304 stainless steel but the substrate was found to corrode relatively
quickly (<6 h in a salt spray test), presumably because of the high chloride content
of the film. When the samples were electrolysed at + 2 V in a 0.1 M KNO3 solution for 2 min they were found to have excellent corrosion resistance. The samples
withstood > 1600 h in a salt spray test without any visible signs of corrosion. The
black coloration of the film was due to the deposits being nanoparticulate and XRD
analysis shows that the 110 and 211 were the predominant crystal faces present.
The deposit thickness, adherence and morphology could be further improved using pulse-plating. A range of brighteners and additives used in aqueous plating
baths was tested, but none was found to cause an improvement in the deposit
morphology.
The relatively high viscosity of these ionic liquids allows improved stability of
particulate suspensions. Consequently, deposition of a range of Cr composites
using Si3 N4 , BN, Al2 O3 or particulate (0.3–1.0 µm) diamond is facile. The scanning
electron micrographs presented in Figure 4.12 show a Cr composite deposited
from a type II Cr/ChCl liquid containing 5 wt.% diamond powder. The crystals
of C (diamond) are clearly visible, held within a Cr metal matrix. EDAX analysis
confirms large amounts of both Cr and C. Similar results were obtained with Si3 N4
and Al2 O3 . The particles do not aggregate upon deposition in the film but rather
remain as discrete entities, suggesting that they are just dragged onto the surface
as the metal deposits.
110 4 Electrodeposition of Metals
Fig. 4.12 (a) and (b) Scanning electron micrographs of a Cr composite deposited from a type II Cr/ChCl liquid containing 5 wt.% diamond
powder.
4.3.3
Type III Eutectics
Eutectics formed between quaternary ammonium salts and hydrogen bond donors
(HBD) have potential windows that tend to be controlled by the stability of the
carboxylic acid, amide or alcohol. In general the potential windows depend upon
the pK a of the HBD. Figure 4.13 shows the potential windows of eutectics formed
between ChCl with ethylene glycol, urea and malonic acid.
The potential windows are significantly smaller than some imidazolium-based
liquids with discrete anions, however, the windows are sufficiently wide for metals
such as zinc and nickel to be electrodeposited with high current efficiencies. The
potential windows are naturally wider on metals that are less catalytic than Pt. Type
III eutectics have the advantages that they are relatively benign and inexpensive
and can thus be applied to large scale processes.
4.3 Deposition of Metals from Non-chloroaluminate Eutectic Mixtures 111
Fig. 4.13 Cyclic voltammograms of (a) urea/ChCl, (b) ethylene glycol/ChCl and (c) malonic acid/ChCl ionic liquids. Voltammograms
were acquired at room temperature (20 ◦ C), using a Pt disk (2 mm
diameter) working electrode, a Ag wire reference electrode and a potential scan rate of 50 mV s−1 .
The deposition of zinc, tin and zinc–tin alloys has been studied in eutectics based
on choline chloride with ethylene glycol and urea [118]. SEM images of Zn–Sn alloys
deposited from choline chloride with ethylene glycol and urea liquids are shown in
Figure 4.14. The morphologies are clearly very different. The electrochemistry of
the different metals was affected by the different HBDs and this was assigned to the
different speciation of the metals in the liquids. In urea the only zinc-containing
species is ZnCl3 − whereas in ethylene glycol ZnCl3 − , Zn2 Cl5 − and Zn3 Cl7 − were
detected. This is because urea acts as a far stronger ligand for ZnCl3 − than ethylene
glycol [118]. The zinc deposits obtained from electrolysis of both the ethylene glycol
and urea-based liquids were similar; dull grey metallic films with good adherence.
The SEM images were similar to those reported by Sun et al. [106], which is
unsurprising given that a type III eutectic with high ZnCl2 concentrations is very
similar in composition to a chlorozincate melt to which ethylene glycol has been
added as a diluent. What is clear however is that the diluent is affecting not only the
viscosity but also the speciation, which will accordingly change the nucleation and
growth processes.
The HBD also affects the way in which alloys form. Voltammetry and XRD
showed that when both Zn and Sn were present in the same melt a homogeneous
ZnSn alloy phase was formed when urea was used as the HBD whereas ethylene
glycol caused separate zinc and tin phases to form.
As with the type II metal-based liquids (e.g. Cr) particulate suspensions are
stabilized by the relatively high viscocity. A dispersion of 3 wt.% Al2 O3 was made
in type III eutectics and mild agitation was sufficient to retain the alumina as
112 4 Electrodeposition of Metals
Fig. 4.14 Scanning electron micrographs obtained by the electrolysis
of 0.5 M ZnCl2 /0.05 M SnCl2 in (a) 1ChCl: 2urea and (b)1ChCl: 2EG,
both at a current density of 10 mA cm−2 for 1 h.
a homogeneous dispersion. EDAX and SEM analysis showed the inclusion of
approximately 1 wt.% Al2 O3 in the film.
Recent work has focussed on converting these liquids into practical plating solutions by investigating ways of improving the morphology of the deposits. In
aqueous solutions a range of compounds is routinely added to act as brighteners.
These are thought to work by either shifting the redox potential of the metal through
complexation or by hindering metal nucleation and growth at the electrode surface.
It was shown that the anion of the metal salt has a significant effect on the reduction potential for the Cu2+ /Cu+ and Cu+ /Cu couples in a ChCl: 2urea eutectic,
even though it was only present in a small concentration (20 mM) compared to the
4.3 Deposition of Metals from Non-chloroaluminate Eutectic Mixtures 113
chloride ion (>1 M). The anion can also change the trend in Cu2+ /Cu+ redox with
respect to that of the Cu+ /Cu couple [99].
The addition of well known complexing agents such as ethylene diamine and
EDTA can have a more significant effect on redox potentials where the position of
stripping potentials can be shifted by over 250 mV. The complexing agents make
it more difficult to reduce the metal and hinder nucleation, which leads to less
nuclei formation and allows the crystals to grow larger before they encounter a
neighboring grain. Bulk deposition from an ionic liquid containing just CuCl2
produces black, powdery deposits whereas the addition of a complexing agent can
lead to lustrous copper deposits [99]. The addition of strong complexing agents to
ionic liquids may not be trivial, however, as it will also affect the charge on the
metal center and the interaction between the metal center and the halide anion
of the ammonium salt, influencing the phase behavior and viscosity. In Type III
eutectics the possibility exists to choose a Brønsted acid that could act as a built-in
brightener.
Most deposition experiments in ionic liquids have been carried out primarily
using halide salts, which is different to most aqueous processes. Oxides have been
found to dissolve in high concentrations in Type III eutectics [119, 120] and these
have been studied for the electrowinning of metals from ores or waste materials
[121]. In principle, however. they could be used for metal plating as they produce
coatings with morphologies similar to those obtained using chlorides. One issue
that arises is the speciation of the oxide following dissolution. FAB-MS data have
shown that some metals, particularly with acid-based hydrogen bond donors revert
to the halometalate complexes, e.g. CuO gives CuCl3 − [120]. Urea-based liquids
give complexes where the oxygen is still attached to the metal center e.g. ZnO gives
[ZnOCl·urea]− [119].
4.3.4
Future Developments
From an academic standpoint there are numerous fundamental issues that still
need to be addressed, including the effect of speciation on the mechanism of nucleation and growth. An understanding of the double layer structure and processes
occurring during deposition is essential for an informed choice of suitable metal
salts and the design of brighteners. Diluents to reduce the liquid viscosity and
make them easier to handle will also have to be identified. It is highly probable that
practical plating solutions will be a complex mixture of salts, viscosity modifiers,
brighteners and wetting agents, analogous to current aqueous plating solutions.
To develop practical plating systems information about the long term stability of
the ionic liquids under high applied current densities needs to be determined. The
effect of adsorbed moisture on deposit morphology also needs to be ascertained as
practical liquids will have to be as robust as possible.
114 4 Electrodeposition of Metals
4.4
Troublesome Aspects
4.4.1
Deposition of Reactive Elements
The deposition of anti-corrosion coatings is one of the main goals of electrochemical
research. The majority of useful metals for this application have extremely negative reduction potentials. Metals such as vanadium, niobium, tantalum, titanium,
magnesium and others are precluded from deposition in ambient temperature
systems due to the narrow potential window of potential solvents. They should,
however, be easily electrodeposited from ionic liquids. The cathodic limits of e.g.
the 1,1-dialkylpyrrolidinium-based ionic liquids are below the electrodeposition
potential for lithium (around – 3 V vs. NHE), thus it should be trivial to deposit
vanadium (−1.17 V), titanium (−1.21 V) and magnesium (−2.34 V), to mention a
few. However, thus far there has been no convincing report on the electrodeposition of magnesium or titanium available in the literature, and we ourselves needed
almost two years to find a suitable way to electrodeposit crystalline tantalum in
micrometer thick layers. In the case of tantalum deposition we started initially with
1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide, a liquid which
has a cathodic limit of about −3 V vs. NHE on Au(111), in which we dissolved
TaCl5 (0.1–0.3 mol l−1 ). Indeed we got a relatively simple cyclic voltammogram
with two reduction processes and poorly defined oxidation reactions.
We were quite optimistic in the beginning as the second reduction process
corresponds to the formation of a black deposit which was potentially the first
electrochemical route to make thick tantalum layers. After having washed off all
ionic liquid from the sample we were already a bit sceptical as the deposit was
quite brittle and did not look metallic. The SEM pictures and the EDX analysis
supported our scepticism and the elemental analysis showed an elemental Ta/Cl
ratio of about 1/2. Thus, overall we have deposited a low oxidation state tantalum
choride. Despite the initial disappointment we were still eager to obtain the metal
and found some old literature from Cotton [122], in which he described subvalent
clusters of molybdenum, tungsten and tantalum halides. In the case of tantalum the
well-defined Ta6 Cl12 2+ complex was described with an average oxidation number of
2.33 and thus with a Ta/Cl molar ratio very close to 1/2. Such clusters are depicted
in Figure 4.15.
In these clusters tantalum atoms are bound to other tantalum atoms and are
also edge bridged via halide. As our deposit was completely amorphous without
any XRD peak we concluded that it did not consist of crystalline tantalum but
rather of such clusters. We varied the electrode potential for deposition and tried
deposition with very low constant current densities, but in no case was crystalline
tantalum obtained. Thus, the electrochemical window of our liquid was surely
wide enough, but for some reason the electrodeposition stopped before Ta(0) was
obtained. When we studied the literature dealing with metal clusters we found that
the cluster chemistry with fluoride seems to be less comprehensive. Consequently
4.4 Troublesome Aspects 115
Fig. 4.15 Some metal/halide clusters described in 1969 by Cotton, Ref. [122].
we tried the deposition in the same liquid with TaF5 (0.25 mol l−1 ) as a source of
tantalum. In Figure 4.16 the cyclic voltammogram of TaF5 in the above-mentioned
liquid is shown at three different temperatures [123].
The cyclic voltammogram is quite similar to the voltammogram of TaCl5 and
consists mainly of two reduction processes. There is no visible surface process at
Fig. 4.16 Cyclic voltammogram of TaF5 in 1-butyl-1methylpyrrolidinium bis(trifluoromethylsulfonyl)amide at variable temperature. The second cathodic process is correlated to the deposition
of a black amorphous material.
116 4 Electrodeposition of Metals
Fig. 4.17 SEM picture of a “tantalum” deposit made from TaF5 in
1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide. The
deposit is obviously amorphous, with EDX a Ta/F ratio of 4/1 is obtained.
the first cathodic peak, at the second cathodic peak a black deposit again forms.
Unfortunately the deposit is still XRD amorphous, and the SEM picture does not
show a crystalline material, Figure 4.17.
However, the EDX analysis made us more optimistic as we got an atomic Ta/F
ratio of 4/1. Our assumption that with fluoride we would get a lower amount of subvalent tantalum fluorides seemed to be right. Furthermore, underneath the black
brittle deposit there was always a thin shining layer which looked metallic. An in
situ STM study showed that triangularly shaped crystals grew on the nanoscale
with a typically metallic behavior in the tunnelling spectrum [123] which, taken
both together, is quite unusual for an amorphous deposit. Upon addition of LiF
we finally found (see Chapter 4.2) parameters with which we could deposit crystalline tantalum layers with thicknesses of about 1 µm. The key parameter was a
low current density for the electrodeposition. With high current densities in the
10 mA cm−2 region we mainly got an amorphous deposit, whereas with current
densities of about 10 µA cm−2 we got metallic tantalum, although the deposition
rate was naturally slow. At least three aspects have to be considered:
1. It is known in inorganic chemistry that all refractory and rare earth elements
tend to form subvalent halides from their iodides, bromides and chlorides.
2. In the case of tantalum deposition the addition of LiF was required. On the one
hand the Li+ ion might destabilize the Ta–F bond and thus facilitate the deposition of tantalum, on the other hand Li+ ions might influence the electrochemical
double layer and facilitate charge transfer. We are about to perform in situ STM
studies, the results will be reported in the peer reviewed literature.
4.4 Troublesome Aspects 117
3. Mass transport may influence material growth in ionic liquids. 1-Butyl-1methylpyrrolidinium bis(trifluoromethylsulfonyl)amide, for example, is, at room
temperature, about 60 times more viscous than water. At temperatures above
150 ◦ C its viscosity is similar to most molecular solvents at ambient conditions.
Indeed, temperatures between 150 and 200 ◦ C were best to deposit tantalum
from TaF5 in the presence of LiF. One has to keep in mind that the deposition of tantalum from TaF5 or an anionic complex delivers one Ta atom and
5–7 fluorides. If the deposition is too fast F− may not diffuse rapidly enough
from the surface to the bulk of the solution and may be trapped in the deposit.
This might explain why we only got crystalline tantalum layers at low current
densities.
In our opinion non-stoichiometric metal halide compounds have to be expected
for the electrodeposition of refractory and rare earth metals if the deposition is performed from halides as precursors. The electrodeposition of these metals requires
tailor-made metal salts.
4.4.2
Viscosity/Conductivity
The view of many electroplaters is that the viscosity of ionic liquids is too high and
the specific conductivity too low to be viable for the deposition of metals. At room
temperature it is surely right that many liquids are viscous compared with aqueous
solutions. However, it is totally neglected that the situation changes completely
when the liquids are heated. Even at a moderate temperature, i.e. 100 ◦ C, many
liquids have viscosities of only a few mPa s, quite similar to water, thus reducing
IR drops in electrochemical experiments considerably. From our point of view we
recommend: “If it doesnt work at room temperature, just heat it up!”. As some ionic
liquids have practical thermal windows as high as 200–300 ◦ C they can be regarded
as the missing link to high-temperature molten salts [124]. Methods for estimating
maximum process operating times and temperatures have been developed [125].
Thus, variation of temperature is rather a benefit and we ourselves were able to show
that at T > 100 ◦ C the grey phase of selenium can be electrodeposited exclusively
[126]. Furthermore one should take into account that there is constant progress in
the synthesis of ionic liquids, thus it can be expected that liquids with viscosities
around 5 mPa s at room temperature (just 5 times more than water) will be available
in future.
4.4.3
Impurities
It is commonly accepted in the ionic liquids community that the purification of
ionic liquids can be relatively complex. Currently, they cannot be distilled at reasonable rates, crystallized or sublimed. Thus, the only reasonable solution is to synthesize them from high quality starting materials. Apart from organic impurities
118 4 Electrodeposition of Metals
(decomposition products of anions and/or cations, side products) halide, Li+ and
K+ are common inorganic impurities. Li+ and K+ can be found in the 1000 ppm
regime if the liquids are made by metathesis reaction from metal salts and organic halides. Sometimes even low amounts of impurities washed off from silica
or alumina (often used to remove the yellowish color of ionic liquids) can be found
in the liquids. These impurities can only be removed by extensive washing with
highest quality water or by an electrochemical treatment with separated cathodic
and anodic compartments. Water is introduced during the washing process, but
usually it can easily be removed by putting the liquids under vacuum at elevated
temperature. Water concentrations of 3 ppm and below are easily obtained for 1butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide with this method.
For a technical electrochemical process 10 ppm of Li+ or K+ might be negligible, for fundamental electrochemical studies on the nanoscale with the in situ
STM (Chapter 9) it is rather a nightmare if there is underpotential deposition
of lithium in a potential regime where the deposition of e.g. silicon is expected.
It took us four months to confirm results for silicon electrodeposition because
of the contamination of one of our liquids with Li+ (see also Refs. [127, 128]).
As will be shown in a later paper the decomposition of the organic cation of an
ionic liquid (e.g. by high galvanostatic pulses) can also strongly alter the morphology of materials. Whereas the electrodeposition in 1-ethyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)amide usually delivers a microcrystalline aluminum, a
nanocrystalline deposit is obtained if the deposition is performed during cathodic
breakdown of the imidazolium cation. This is also a kind of in situ made impurity
which can strongly alter electrodeposition.
Although almost all of the modern ionic liquids are per se air- and water-stable one
has to bear in mind that upon addition of SiCl4 , TaF5 , SeCl4 , AlCl3 and other moisture sensitive compounds the resulting solutions are no longer water-stable. Thus,
inert gas conditions are required to get reproducible results. From our experience
only ultrapure ionic liquids should be employed for fundamental electrochemical
studies unless the influence of impurities has been understood. Our own experience and that of many other groups have shown that even a few ppm of impurities
can strongly alter fundamental studies.
4.4.4
Additives
There is a phalanx of different additives available in aqueous electroplating. We
have often had the experience that aqueous electroplaters add personal additive
recipes to ionic liquids and are surprised, or even disappointed, that they do not
work. One has to bear in mind that all known additives were developed over several
decades and their mode of action in aqueous solutions is still not fully understood.
How can it be expected that one can just replace water by the ionic liquid and get the
same or even better results? In our opinion a deep understanding of cation/anion
4.4 Troublesome Aspects 119
interactions with dissolved substances is required to develop additives that are
suited for ionic liquids.
4.4.5
Cation/Anion Effects
We were the first to find that there are cation/anion effects in the electrodeposition of metals. In the case of aluminum we found that it is deposited as a nanomaterial in 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide from
AlCl3 , whereas it is deposited as a microcrystalline material under quite similar conditions in 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide
[129]. The likely explanation is that the pyrrolidinium ion interferes with the electrode surface and the growing nuclei, thus hindering crystal growth [130]. Maybe
these cation/anion effects explain why in first generation ionic liquids there are
about 100 papers on the electrodeposition of Al and its alloys from AlCl3 and
1-ethyl-3-methylimidazolium chloride and only a few with tetraalkylammonium
chlorides/AlCl3 . Our own experiments have shown that the deposition of Al from
1-butyl-1-methylpyrrolidinium chloride/AlCl3 delivers a crystalline but rather flakelike black product. Thus it might be a bad choice to employ cheap liquids for electrodeposition. In our opinion a deep understanding of cation/anion interferences
is required and one should be aware of unexpected effects by just employing a
different ionic liquid.
4.4.6
Price
A main reproach is that the cost of ionic liquids is too high at the moment. Indeed,
1 kg of ultrapure 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide
costs up to 2000 € . One has to bear in mind that currently one pays more or less
the salary of the technician in the laboratory who synthesizes the liquid from the
educts. If a large scale production line was available, operated automatically, the
costs would be reduced drastically. There is a dispute in the community about what
future prices will be. Currently it is believed that in a few years from now the prices
will start at about 10 € per liter for standard ionic liquids with prices up to 10 000 €
per kg for tailor-made “research liquids”. The first generation ionic liquids based
on AlCl3 and dialkylimidazolium chlorides are candidates for such comparatively
cheap liquids and a price between 10 and 15 € per kg is conceivable. One should
not forget that ionic liquids have practically no vapor pressure and that they can
easily be recycled, as shown in Chapter 11.4. Thus the overall costs for a process
will decide whether an ionic liquids process will be established or not. In the case
of Al electrodeposition there would be an immediate advantage of ionic liquids: in
contrast to the SIGAL process where Al is deposited from explosive alkyl-aluminum
compounds, thus strictly requiring inert gas, dry air would be sufficient in the case
120 4 Electrodeposition of Metals
of an ionic liquids process. It is likely that the overall costs would be at the same
level or even lower.
4.4.7
One Liquid for All Purposes?
The dream would be that there will be – like water – one ionic liquid that is suited
as a general liquid for all electrochemical reactions. It cannot be excluded that
such a liquid will be produced in the future, but at present the field is in rather
a developmental state. We ourselves were pretty surprised when we realized that
the cation of an ionic liquid can have a dramatic effect on the electrodeposition
of metals. A deeper understanding of ionic liquids will be required before ionic
liquids become standard electrolytes for electroplating.
The motivation of this chapter was to show that despite the enormous prospects
of ionic liquids in electrodeposition some troublesome aspects have to be expected.
Apart from the purity and price of ionic liquids the optimum temperature for any
process has to be found. Furthermore, suitable additives for electrodeposition will
have to be developed and cation/anion effects that can strongly alter the morphology
of deposits have to be expected. Finally, the electrochemical window alone is not
the only factor that needs to be considered for the deposition of reactive metals.
Suitable precursors will have to be tailor-made and it is our personal opinion that
the electrodeposition of metals like Mg, Ti, Ta and Mo may not be possible from
metal halides but rather metal bis(trifluoromethylsulfonyl)amide salts and other
ones may be more suitable.
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125
5
Electrodeposition of Alloys
I.-Wen Sun, and Po-Yu Chen
5.1
Introduction
Electrodeposition of alloys is an important subject as alloys often provide properties
superior to those of single-metal electrodeposits. The electrodeposited alloys can
be more corrosion resistant, more wear resistant, better in catalytic properties and
better in magnetic properties. Similar to the deposition of pure metals, the properties of the electrodeposited alloys can be varied by experimental factors such as
plating bath composition, current density, overpotential and temperature. Furthermore, applying pulsed electrodeposition allows one to influence the grain size of
the deposits (see Chapter 9). The interest in the investigation of electrodeposition
of alloys is increasing, quite simply because the number of alloy combinations is
vast. While aqueous plating baths are widely employed for the electrodeposition
of alloys, the limited electrochemical window and hydrogen evolution problems
have considerably restricted the number of alloys that can be electrodeposited from
aqueous plating baths.
Over the past two decades, ionic liquids (ILs) have attracted considerable interest
as media for a wide range of applications. For electrochemical applications they
exhibit several advantages over the conventional molecular solvents and high temperature molten salts: they show good electrical conductivity, wide electrochemical
windows of up to 6 V, low vapor pressure, non-flammability in most cases, and
thermal windows of 300–400 ◦ C (see Chapter 4). Moreover, ionic liquids are, in
most cases, aprotic so that the complications associated with hydrogen evolution
that occur in aqueous baths are eliminated. Thus ILs are suitable for the electrodeposition of metals and alloys, especially those that are difficult to prepare in an
aqueous bath. Several reviews on the electrodeposition of metals and alloys in ILs
have already been published [1–4]. A selection of published examples of the electrodeposition of alloys from ionic liquids is listed in Table 5.1 [5–40]. Ionic liquids
can be classified into water/air sensitive and water/air stable ones (see Chapter 3).
Historically, the water-sensitive chloroaluminate first generation ILs are the most
intensively studied. However, in future the focus will rather be on air- and waterstable ionic liquids due to their variety and the less strict conditions under which
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
126 5 Electrodeposition of Alloys
Table 5.1 Metal alloys that have been electrodeposited from ionic liq-
uids.
Alloy
Ref.
Ionic liquid
Al–Nb
Al–Ni
Al–Co
Al–Cr
Al–Cu
Al–Mn
Al–La
Al–Ag
Al–Ti
Al–Mo
Al–Zr
Al–Pt
Al–Mg
Al–Mo–Mn
Al–Cr–Ni
Zn–Cu
Zn–Cd
Zn–Sn
Zn–Co
Zn–Fe
Zn–Ni
Zn–Mg
Zn–Pt
Pt–Zn
Au–Zn
Ag–Zn
Nb–Sn
Pd–Au
Pd–Ag
Pd–In
In–Sn
Cu–Sn
Zn–Mn
[5]
[6]
[7–10]
[11]
[12]
[13]
[14]
[15]
[16, 17]
[18]
[19]
[20]
[21]
[22]
[23]
[24]
[25]
[26]
[27, 28]
[29]
[30]
[31]
[32]
[34]
[35]
[36]
[37]
[38]
[39]
[40]
[41]
[42]
[46]
[EMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3 , [BP]+ Cl− /AlCl3
[BP]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3 , [BMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3
[BTMA]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /AlCl3
[EMIM]+ Cl− /ZnCl2
[EMIM]+ Cl− /ZnCl2
[EMIM]+ Cl− /ZnCl2
[BP]+ Cl− / ZnCl2 , [EMIM]+ Cl− /ZnCl2
[EMIM]+ Cl− /ZnCl2
[EMIM]+ Cl− /ZnCl2 /NiCl2
[EMIM]+ Br− /ZnBr2 /MgBr2 /EG
[EMIM]+ Cl− /ZnCl2
[EMIM]+ Cl− /ZnCl2
[EMIM]+ Cl− /ZnCl2
[EMIM]+ Cl− /ZnCl2
[EMIM]+ Cl− /SnCl2 /NbCl5
[EMIM]+ BF4 −
[EMIM]+ BF4 −
[EMIM]+ BF4 −
[EMIM]+ BF4 −
[TMHA]+ Tf2 N−
[TBMA]+ Tf2 N−
they can be handled. The principles of alloy electrodeposition are beyond the scope
of this book therefore for an overview on alloy electrodeposition the book by Brenner
[44] is recommended.
5.2
Electrodeposition of Al-containing Alloys from Chloroaluminate Ionic Liquids
The Lewis acidic chloroaluminate ILs are suitable for the electrodeposition of
aluminum-containing alloys. Many examples have been published but those that
have been reviewed in detail by Stafford and Hussey [1] will not be included in this
section.
5.2 Electrodeposition of Al-containing Alloys from Chloroaluminate Ionic Liquids 127
5.2.1
Al–Ti
The electrodeposition of Al–Ti alloy has been investigated in the acidic aluminum
chloride-1-ethyl-3-methylimidazolium chloride ([EMIM]+ Cl− /AlCl3 ) ionic liquid
containing Ti(II) up to 0.17 mol l−1 at 353 K [16]. Such alloys are technically interesting due to their high temperature resistance. Ti(II) can be introduced into the
liquid by direct dissolution of TiCl2 or by the reduction of TiCl3 with Al metal in the
liquid. It was proposed that TiCl2 dissolves in the liquid by forming [Ti(AlCl4 )3 ]−
and its solubility increases with increasing IL acidity.
TiCl2 (s) + 2Al2 Cl7 − → [Ti(AlCl4 )3 ]− + AlCl4 −
(5.1)
Lowering the liquid acidity from 66.7–33.3% to 60.0–40.0% mole fraction results
in the disproportionation of [Ti(AlCl4 )3 ]− producing TiCl3 and Ti precipitates.
3[Ti(AlCl4 )3 ]− + 3AlCl4 − ↔ 2TiCl3 (s) + Ti + 6Al2 Cl7 −
(5.2)
Ti(II) tends to form polymers or aggregates upon increasing the Ti(II) concentration
or the liquid acidity. Electrochemical, either galvanostatic or potentiostatic, oxidation of Ti metal produces either passive TiCl3 film or volatile TiCl4 which escapes
from the liquid. The oxidation of metallic titanium to Ti(II) by direct anodization
of Ti metal in this liquid has not yet been described.
Cyclic voltammograms recorded on polycrystalline stationary and rotating Pt
disk electrodes in the acidic [EMIM]+ Cl− /AlCl3 ionic liquids demonstrated that
the reduction of Ti(II) to Ti(0) occurs at a potential where the deposition of Al
also occurs. As an Al–Ti alloy forms, both the deposition and stripping waves shift
to values more positive than the pure Al oxidation. The magnitude of the shifts
increases with increasing Ti(II) concentration. Bulk deposits of Al–Ti alloys were
prepared by using DC galvanostatic electrolysis on a copper rotating disk electrode
(Cu-RDE) at a current density of −10 mA cm−2 in a Ti(II)-saturated liquid. The Ti
metal content of bulk Al–Ti alloys prepared in this way decreased with increasing
applied current density, suggesting that the reduction potential of the Ti(II)/Ti couple would be positive of that for the Al(III)/Al couple. Increasing the total applied
reduction current densities or making the applied potential more negative probably leads to a mass-transport-limited value for Ti deposition, whereas the partial
current density for the deposition of Al still increases, resulting in alloy deposits
with decreased Ti content. The Ti content in the Al–Ti electrodeposited from this
ionic liquid is limited by the solubility of Ti(II) in the liquid and by the minimum
practical current density that can be applied. Scanning electron micrographs of the
electrodeposited Al–Ti alloys revealed that the deposits were compact dense nodules of single crystals. The nodule size decreases with decreasing current density
and increasing Ti content. X-ray powder diffraction (XRD) patterns of the electrodeposits containing 7.0 to 18.4 a/o Ti metal showed a disordered face-centered cubic
(fcc) structure, very similar to that of pure Al. The aluminum lattice parameter
decreases as the smaller Ti atoms substitute for Al. Furthermore, X-ray reflections
128 5 Electrodeposition of Alloys
of the Al–Ti alloys broaden with increasing Ti content, suggesting a decrease in
the grain size of the deposit. Potentiodynamic anodic polarization curves recorded
for Al–Ti alloys electrodeposited on copper electrodes in deaerated aqueous NaCl
solution revealed that, similar to what is found for Al–Ti alloys prepared by sputter
deposition, the electrodeposited Al–Ti alloys exhibit a significant increase in pitting
potential relative to pure Al.
The electrodeposition of Al–Ti alloys has also been examined at 298 K on
Au(111) in an acidic aluminum chloride-1-butyl-3-methylimidazolium chloride
([BMIM]+ Cl− /AlCl3 ) containing 10 mM TiCl4 [17]. Cyclic voltammograms showed
that Ti(IV) can be electrochemically reduced to Ti(III) in the form of hardly soluble
TiCl3 , which can be further reduced to Ti(II) at a potential close to the underpotential deposition (UPD) of Al on Au(111), followed by the co-deposition of Al–Ti
prior to the Al bulk deposition. The stripping of Al–Ti can be observed during
the anodic scan. Comparing the electrochemical scanning tunneling microscopy
(EC-STM) images of the deposits revealed that Al UPD clusters preferentially deposit along the Au step edges in the absence of Ti whereas the UPD of Al–Ti begins
with the formation of monoatomically high clusters on the Au terraces without any
site preference. The formation of an Al–Ti phase in the electrodeposits was further
confirmed by X-ray photoelectron spectra (XPS) analysis.
5.2.2
Al–Mo
The electrodeposition of Al–Mo high-temperature and corrosion resistant alloy was
investigated in a Lewis acidic [EMIM]+ Cl− /AlCl3 ionic liquid using the octahedral hexanuclear Mo(II) cluster compound, (Mo6 Cl8 )Cl4 [18]. A previous study [42]
showed that in basic [EMIM]+ Cl− /AlCl3 liquid, (Mo6 Cl8 )Cl4 picks up two excess
chloride ions from the liquid to form [(Mo6 Cl8 )Cl6 ]2− complex anion but the reduction of this species does not produce Mo metal. The (Mo6 Cl8 )Cl4 is soluble in the
acidic [EMIM]+ Cl− /AlCl3 and preserves its {Mo6 Cl8 }4+ core structure. However,
it cannot be oxidized within the anodic potential limit of this liquid and cannot be
reduced prior to the electrodeposition of Al.
Cyclic voltammograms recorded at Pt stationary and rotating disk electrodes
in 66.7 mol% [EMIM]+ Cl− /AlCl3 liquid show that the addition of (Mo6 Cl8 )Cl4
results in slightly negative shifts in the potential of the electrodeposition process.
Furthermore, the stripping wave of pure Al deposits is replaced by a new stripping
wave at a more positive potential. These results indicate that Al–Mo alloys are
formed through the following reduction process.
x{Mo6 Cl8 }4+ + 8(3 − 2x) Al2 Cl7 − + 6(3 − x) e−
= 6Al1−x Mox + 2(21 − 13x)AlCl4 −
(5.3)
Electrodeposits of Al–Mo alloys were prepared at Cu rotating disk electrode and rotating wire electrode substrates with galvanostatic electrolysis and examined with
5.2 Electrodeposition of Al-containing Alloys from Chloroaluminate Ionic Liquids 129
energy dispersive X-ray (EDX), SEM, and XRD for compositional and morphological analysis. As the concentration of {Mo6 Cl8 }4+ in the plating bath is much
smaller than that of Al2 Cl7 − , the partial current density for {Mo6 Cl8 }4+ reduction
is fixed and small, and increasing the total applied deposition current density simply leads to more Al deposition. As a result, the Mo content of the alloy decreases
with increasing applied current density. Increasing the {Mo6 Cl8 }4+ concentration
increases the partial current density for Mo deposition. Thus, at a fixed total applied
current density, the Mo content increases with increasing {Mo6 Cl8 }4+ concentration. The partial current for Mo deposition relative to that for Al deposition can
also be increased by increasing the deposition temperature, leading to higher Mo
content in the alloy. SEM images of the Al–Mo electrodeposits reveal a surface
morphology consisting mainly of spherical nodules and a nearly specular surface
could be obtained. EDX maps for Al and Mo in this electrodeposit indicate that both
elements are distributed more or less evenly over the surface of the deposit. XRD
analysis of the Al–Mo alloy deposits shows that those containing less than 5 atom%
Mo are single phase, supersaturated solid solutions having an fcc structure very
similar to that of pure Al. Broad reflection indicative of an amorphous phase appears in deposits containing more than 6.5 atom% Mo. As the Mo content of the
deposits is increased, the amount of fcc phase in the alloy decreases whereas that of
the amorphous phase increases. When the Mo content is more than 10 atom%, the
deposits are completely amorphous. As the Mo atom has a smaller lattice volume
than Al, the lattice parameter for the deposits decreases with increasing Mo content. Potentiodynamic anodic polarization experiments in deaerated aqueous NaCl
revealed that increasing the Mo content for the Al–Mo alloy increases the pitting
potential. It appears that the Al–Mo deposits show better corrosion resistance than
most other aluminum–transition metal alloys prepared from chloroaluminate ionic
liquids.
5.2.3
Al–Zr
The electrodeposition of Al–Zr alloys was examined in the 66.7–33.3 mol%
[EMIM]+ Cl− /AlCl3 liquid [19]. The reduction of Zr(IV), which was introduced as
ZrCl4 in the liquid, produces a small ill-defined cathodic wave and a small negative
shift to the Al deposition wave. Voltammetric data show that the small ill-defined
cathodic wave corresponds to the Zr(IV)/Zr(III) reaction. It is noted that a surface
passivating film is formed on the electrode surface after this reaction, indicating
that the Zr(III) is insoluble in the liquid.
Solutions of Zr(II) can be prepared by chemical reduction of Zr(IV) with Al or Zr
metals. The in situ formtion of Zr(II) is, however, more efficient with Al metal in
more acidic liquid. The maximum concentration of Zr(II) that could be produced
by the Al reduction of Zr(IV) was about 0.02 mol l−1 in 66.7–33.3 mol% liquid. The
measured diffusion coefficient for the Zr(II) species is much smaller than that
for the Zr(IV) species and decreases as the Zr(II) concentration increases. This
phenomenon suggests that polymerization of Zr(II) may possibly have occurred.
130 5 Electrodeposition of Alloys
The electrodeposition of Al–Zr alloys was investigated by using galvanostatic
electrolysis using Cu rotating wire electrodes at 353 K in the 66.7–33.3 mol% liquid
containing either Zr(IV) or Zr(II). As a limiting current is more rapidly obtained
for the reduction of Zr(II) than for the reduction of Al2 Cl7 − , the partial current for
the deposition of Zr reaches a constant value whereas that for the deposition of Al
continuously increases with increasing current density. As a result, the Zr content
of the electrodeposited alloys decreases with increasing current density. It was also
noted that because the diffusion coefficient of Zr(II) is much smaller than that of
Zr(IV), solutions of Zr(II) lead, in Al–Zr alloys, to smaller amounts of Zr- relative
to Zr(IV)-containing solutions of equal concentration.
Results from SEM and XRD examinations revealed that the structure of the
Al–Zr deposits varies with the alloy composition. Al–Zr deposits containing less
than 5 atom% Zr consist of nodules of fcc crystals similar to pure Al. The nodules
decrease in size with increasing Zr composition and decreasing current density,
suggesting that a certain grain refining is driven by the incorporation of Zr into
the alloy rather than by the deposition overpotential. In addition to fcc Al, an amorphous phase becomes apparent when the Zr content for the Al–Zr alloy is further
increased, and the alloy deposit containing 16.6 atom% Zr is completely amorphous. The corrosion resistance of the electrodeposited Al–Zr alloy was examined
by pitting potential measurements. It was found that the addition of 8 atom% or
more Zr increases the pitting potential of the alloy by about +0.3 V vs. pure Al.
5.2.4
Al–Pt
The electrodeposition of Al–Pt, which is an interesting material for catalysis, has
been studied in Lewis acidic ionic liquids formed from AlCl3 with benzyltrimethyl
ammonium chloride ([BTMA]+ Cl− /AlCl3 ) [20]. This ionic liquid is slightly less
water sensitive than [EMIM]+ Cl− /AlCl3 . Cyclic voltammograms recorded at an Fe
electrode in a 1:2 [BTMA]+ Cl− /AlCl3 liquid containing BTMA2 PtCl6 showed that
the reduction of Pt(II) is slightly less negative than the reduction of Al(III). This is
surprising at first glance but shows to what amount complexation can alter electrode
potentials. Constant potential electrolysis was employed to prepare Al–Pt deposits in
liquids containing different platinum complex ions such as tetraaminoplatinum(II)
and bis(acetylacetonato)platinum(IV). At a more negative potential (−1 V vs. an Al
quasi-reference electrode immersed in the same liquid) where bulk deposition of
Al would occur, poorly adherent powders were obtained which contained primarily
Al and some trapped chloride. At a less negative applied potential (−0.6 V) where
the deposition of Al may still be in the kinetic controlled region, bright, adherent
Al–Pt alloy deposits with dense nodules could be obtained with negligible amounts
of chloride. The Pt content for the Al–Pt deposits increased (from 5% to 13% by
weight) with increasing Pt(II) (or Pt(IV)) concentration, which depended on the
solubility of the Pt complex used in the solution. EDX analysis suggested that
the Al–Pt deposits were homogeneous. At even less positive deposition potential
(−0.2 V) bright, adherent pure Pt could be obtained although the deposition rate was
very slow. Information from XRD analysis was, however, not provided in this study.
5.2 Electrodeposition of Al-containing Alloys from Chloroaluminate Ionic Liquids 131
5.2.5
Al–Mg
Al-Mg alloys are widely used for chemical-processing and food-handling equipment. While electrodeposition may be an effective and cost efficient option for
the preparation of thin alloy coatings, the fact that the standard potential of the
Mg(II)/Mg couple is much more negative than that of the Al(III)/Al makes the
electrodeposition of Al–Mg alloys, at first glance, implausible. Nevertheless, induced codeposition of Al–Mg alloys from Lewis acidic [EMIM]+ Cl− /AlCl3 (mole
ratio 1:2) ionic liquid at 30 ◦ C has been examined by Morimitsu et al. [21]. Mg(II)
was introduced to the liquid by dissolution of 0.2 mol kg−1 MgCl2 . Cyclic voltammograms were recorded at a tungsten working electrode vs. an Al(III)/Al reference
electrode. They show that the deposition of pure Al at a W substrate requires a
large nucleation overpotential. Addition of MgCl2 reduces the overpotential and
shifts the deposition process to less negative electrode potentials. Multiple stripping waves appear in the presence of MgCl2 and the relative magnitude of these
stripping waves depends on the reversing potential, indicating the codeposition of
Mg with Al. The Mg atomic content of the deposits was found to increase with
increasing current density (or cathodic overpotential), supporting the fact that deposition of Mg starts at a potential more negative than that of pure Al. The alloy
deposits obtained in this study were single phase Al–Mg solid solutions. As the
Mg atomic content was very low, 2.2 atom%, the XRD patterns of the alloy deposits
were almost identical to that of pure Al. The codeposition of Mg with Al in this
study was classified as “induced codeposition” of which the alloy deposition occurs
at more positive potentials than the deposition of the more noble metal (Al in this
case) of the alloy components [43].
5.2.6
Al–Mo–Mn
Since both Al–Mo [18] and Al–Mn [13] alloys can be electrodeposited from Lewis
acidic [EMIM]+ Cl− /AlCl3 liquid, Tsuda et al. investigated the electrodeposition of
the Al–Mo–Mn ternary alloys in the 33.3–66.7% mole ratio [EMIM]+ Cl− /AlCl3
liquid containing Mo(II) as (Mo6 Cl8 )Cl4 and Mn(II) as MnCl2 at 55 ◦ C [22].
Cyclic voltammograms recorded at a Pt disk electrode for the solutions revealed
that the Al–Mo–Mn electrodeposition process varies with the concentration ratio
CMn(II) /CMo(II) and that the presence of Mn(II) in the solution inhibits the nucleation of Al. Controlled current techniques were employed to prepare Al–Mo–Mn
alloy samples at a copper rotating wire electrode. The relationships between the
total applied current density, jt , and the partial current densities of Mo, Mn, and
Al were expressed as jMo , jMn , and jAl . Plots of −jMo , −jMn , and −jAl vs. −jt show
that −jAl varies linearly with jt . It appears that the deposition of Mo is inhibited
by Mn(II) so that −jMo reaches a limiting value and becomes a smaller fraction of
the total current density as –jt is increased. Because of this, the Mo content of the
Al–Mo–Mn alloy decreases with increasing −jt . On the other hand, −jMn increases
with −jt so that the Mn content of the alloys increases with increasing −jt . Overall,
132 5 Electrodeposition of Alloys
Al–Mo–Mn alloys rich in Mo are obtained at small −jt , whereas alloys rich in Mn
are obtained at large −jt , provided that CMn(II) /CMo(II) ≧ 2. The inhibition of the
deposition of the more noble component, Mo, by the less noble component, Mn,
makes the codeposition of Al, Mo, and Mn an anomalous process. SEM images
of the Al–Mo–Mn alloy deposits reveal that the morphology varies from spherical
nodules to a shining surface, depending on the current density and the Mo and
Mn concentrations. XRD analysis revealed that the deposits containing less than
approximately 10 atom% Mo + Mn exhibited a face-centered cubic Al and an amorphous phase. When the concentration of Mo + Mn exceeded 10 atom%, only an
amorphous phase was observed. Pitting potential measurements of the electrodeposited Al–Mo–Mn alloys revealed that the addition of relatively modest amounts
of Mo and Mn to the alloy resulted in a significant increase in corrosion resistance
compared to pure Al and the comparable binary alloy containing only one of the
transition metal components.
5.2.7
Al–Cr–Ni
Alloys with Al, Cr and Ni are of technical importance due to their excellent temperature and corrosion resistance. The electrodeposition of an Al–Cr–Ni layer was
attempted in 1:2 [EMIM]+ Cl− /AlCl3 liquid containing 5 × 10−2 mol l−1 NiCl2 and
6 × 10−2 mol l−1 CrCl2 at 338 K [23]. Cyclic voltammetric experiments showed
that the reduction of Ni(II) occurs at a potential more positive than for Cr(II) and
Al(III) whereas the reduction potential of Cr(II) almost overlapped with that of
Al(III). Al–Cr–Ni alloy samples were prepared by constant potential electrolysis
at glassy carbon substrates and their compositions were analyzed by fluorescence
X-ray spectroscopy. The results showed that the atomic ratio of Al:Cr:Ni was 97:2:1
at a potential where bulk deposition of Al occurred and 90:1:9 at a less negative
potential. The low atomic ratio of Cr and Ni in the deposits is partly due to the
low concentration of Ni(II) and Cr(II) in comparison to that of Al(III). To increase
the Ni and Cr content in the deposit, pulse potential electrolysis was adopted. In
a typical pulse electrolysis cycle, the potential was first stepped to a sufficiently
negative value for the codeposition of Al, Cr, and Ni, and then the potential was
applied to a more positive value for the dissolution of Al. By changing pulse conditions, including negative and positive potentials and frequency, the concentration
of Ni and Cr in the deposits was enhanced to 20–27 and 15 atom%, respectively.
5.3
Electrodeposition of Zn-containing Alloys from Chlorozincate Ionic Liquids
Ionic liquids can be obtained by the combination of zinc halide with certain organic
halides (see Chapter 3.3). As the cathodic potential limit of Lewis acidic liquids
(with more than 33 mol% zinc halide) is due to the deposition of metallic zinc, the
5.3 Electrodeposition of Zn-containing Alloys from Chlorozincate Ionic Liquids 133
electrodeposition of zinc and its alloys from these ionic liquids is feasible. Some of
the studies are described below.
5.3.1
Alloys of Zn with Cu, Cd and Sn
The electrodeposition of Zn–Cu, Zn–Sn, and Zn–Cd has been investigated in Lewis
acidic [EMIM]+ Cl− /ZnCl2 liquid containing Cu(II) [24], Cd(II) [25] and Sn(II) [26],
respectively. Figure 5.1 illustrates the cyclic voltammograms of the 50.0–50.0 mol%
[EMIM]+ Cl− /ZnCl2 with and without Cu(I). A typical alloy formation was observed.
The deposition of Cu, Sn, and Cd occurs at a potential of 0.5, 0.3 and 0.1 V,
respectively, more positive than the deposition of Zn. In these studies samples
of the alloys were prepared on Ni substrates by constant potential electrolysis
and examined with EDX, SEM, and XRD. It was found that the Zn content in
the electrodeposits increased as the deposition potential became more negative
but decreased with increasing concentrations of Cu(II), Sn(II), and Cd(II) in the
solution. Increasing the deposition temperature increases the mass-transport rates
Fig. 5.1 Staircase cyclic voltammograms for the 50.0–50.0 mol%
[EMIM]+ Cl− /ZnCl2 liquid on tungsten and nickel electrodes at
80 ◦ C, (a) and (b) with 200 mM Cu(I), (c) without Cu(I). Scan rate
50 mV s−1 . [Ref. 24].
134 5 Electrodeposition of Alloys
of the metal ions in the plating bath and decreases the overpotential required for
the deposition. As a result, increasing the temperature increases the content of the
more noble metal (Cu, Cd, and Sn) in the deposit.
5.3.2
Zn–Co
The electrodeposition of Zn–Co and Zn–Fe alloys in an aqueous bath is classified
as an anomalous codeposition [44] because the less noble Zn is preferentially deposited with respect to the more noble metal. This anomaly was attributed to the
formation of Zn(OH)+ which adsorbs preferentially on the electrode surface and
inhibits the effective deposition of the more noble metal. This anomaly was circumvented by using zinc chloride-n-butylpyridinium chloride ([BP]+ Cl− / ZnCl2 ) [27] or
[EMIM]+ Cl− /ZnCl2 [28] ionic liquids containing Co(II). The Zn–Co deposits can
be varied from Co-rich to Zn-rich by decreasing the deposition potential or increasing the deposition current. XRD measurement reveals the presence of Co5 Zn21
in the deposited Zn–Co alloys and that the Co-rich alloys are amorphous and the
crystalline nature of the electrodeposits increases as the Zn content of the alloys
increases. Addition of propylene carbonate cosolvent to the ionic liquid decreases
the melting temperature of the solution and allows the electrodeposition to be performed at a lower temperature. The presence of CoZn alloy is evidenced by the
XRD patterns shown in Figure 5.2.
Fig. 5.2 (A) XRD patterns of electrodeposits produced from a 60.0–40.0 mol%
[EMIM]+ Cl− /ZnCl2 liquid containing
1.16 wt% of CoCl2 at 80 ◦ C at deposition
potential of (b) 0.13, (c) −0.17, (d) −0.22,
and (e) −0.26 V. For comparison, the XRD
pattern of a pure zinc deposit is given in
(a). (B) XRD patterns of electrodeposits
produced from 6.0 g of 60.0–40.0 mol%
[EMIM]+ Cl− /ZnCl2 liquid containing
1.16 wt% of CoCl2 and 6.0 g of propylene
carbonate at 40 ◦ C at a deposition potential
of (a) 0.20 and (b) −0.27 V. [Ref. 28].
5.3 Electrodeposition of Zn-containing Alloys from Chlorozincate Ionic Liquids 135
5.3.3
Zn–Fe
Similar to the electrodeposition of Zn–Co, the electrodeposition of corrosion resistant Zn–Fe alloy in an aqueous bath is an anomalous codeposition and the
Zn/Fe ratio in the deposit is higher than that in the electrolyte. However, nonanomalous deposition of Zn–Fe was achieved by conducting the deposition in
a 60.0–40.0 mol% [EMIM]+ Cl− /ZnCl2 ionic liquid containing Fe(II) [29]. Cyclic
voltammograms showed that the deposition of Fe occurs at a potential less negative than that of Zn. Underpotential deposition of Zn on Fe occurred through an
instantaneous two-dimensional nucleation process observed prior to the deposition
of bulk Zn. It is interesting to note that, as shown in Figure 5.3, the Zn adatoms
from UPD were able to diffuse into the bulk iron to form Zn–Fe alloy. Zn–Fe
alloy could also be prepared at potentials where bulk deposition of Zn occurred.
The Fe content in the deposit can be varied from 100 to 50 atom% by decreasing
the deposition potential or the Fe(II) concentration in the solution. SEM images
of the deposits revealed that they were dense and compact, and the morphology
varied from nodules to pyramidal and hexagonal as the iron content in the deposits
decreased.
Fig. 5.3 SEM analysis of Zn electrodeposits obtained at −0.12 V where only UPD
of Zn on Fe substrate occurs from pure
60.0–40.0 mol% [EMIM]+ Cl− /ZnCl2 ionic
liquid, on to a 0.25 cm2 iron foil: (a) sec-
ondary electron image of a polished crosssection; (b) EDS line scan of the polished
cross-section(scanned along the white line
in the micrograph). [Ref. 29].
136 5 Electrodeposition of Alloys
5.3.4
Zn–Ni
The electrodeposition of Zn–Ni alloy from the [EMIM]+ Cl− /ZnCl2 /NiCl2 ionic liquid was examined by Koura et al. [30]. The Ni content in the deposit decreases
from 98.6 to 12.3 mol% with increasing deposition current density. Due to the high
viscosity of the [EMIM]+ Cl− /ZnCl2 /NiCl2 liquid, the current efficiency was low,
even when the temperature was increased to 100 ◦ C. XRD analysis of the deposits
revealed both amorphous and crystalline ZnNi, but the formation of the Zn21 Ni5
compound was not observed. In order to reduce the viscosity and to enhance the current efficiency, ethanol (EtOH) was added to the [EMIM]+ Cl− /ZnCl2 /NiCl2 liquid
at 40 ◦ C. The current efficiency was improved to almost 100% in all the electrodepositions performed in the [EMIM]+ Cl− /ZnCl2 /NiCl2 /EtOH solution. The XRD patterns of the deposits obtained from the [EMIM]+ Cl− /ZnCl2 /NiCl2 /EtOH showed
both crystalline Zn21 Ni5 and amorphous ZnNi. Differential scanning calorimetry (DSC) of the deposit showed exothermic peaks that were attributed to the
amorphous-to-crystalline transformation together with crystal growth.
5.3.5
Zn–Mg
The electrodeposition of Zn–Mg alloy was examined in mixtures of 1-ethyl-3methylimidazolium bromide ([EMIM]+ Br− )/ZnBr2 /MgBr2 /ethylene glycol (EG)
at 120 ◦ C.[31] The total concentration of ZnBr and MgBr in the bath was kept at
10 mol% while the ZnBr2 /MgBr2 mole ratio was varied. Linear scan voltammetry revealed a single reduction wave resulting from the deposition of Zn in the
[EMIM]+ Br− /ZnBr2 /EG solution. The addition of MgBr2 shifted the reduction
wave to more negative electrode potentials due to the codeposition of Mg. Dense
Zn–Mg alloys could be electrodeposited by potentiostatic electrolysis at a Cu substrate. The XRD diffraction analysis showed that the alloys contained Zn11 Mg2 in
addition to Zn, and CuZn5 . EDX analysis showed that the Zn and Mg were distributed uniformly in the alloy. The composition of the alloy could be controlled by
the ionic liquid bath composition. A steel sample that was coated with Zn–Mg alloy
containing 2.5 mol% Mg showed significantly improved corrosion resistance.
5.3.6
Pt–Zn
The electrodeposition of Pt–Zn from a 60.0–40.0 mol% [EMIM]+ Cl− /ZnCl2 liquid
containing Pt(II) was investigated at 90 ◦ C [32]. Cyclic voltammograms showed
that the reduction of Pt(II) to Pt occurs at a potential slightly less negative than
the reduction of Zn(II). Multiple anodic stripping waves were observed for the
Pt–Zn electrodeposits, indicative of a multiphasic structure of the deposits. The
Zn component of the deposits was stripped at a less positive potential than the Pt
component of the deposits. Samples of Pt–Zn deposits containing 8 to 42 atom%
5.4 Fabrication of a Porous Metal Surface by Electrochemical Alloying and De-alloying 137
Fig. 5.4 The XRD pattern of the Pt–Zn coating (Pt a/o = 40.91%)
that was electrodeposited on a tungsten foil at a deposition potential
of −0.2 V in the 60.0–40.0 mol% [EMIM]+ Cl− /ZnCl2 ionic liquid containing 120 mM PtCl2 at 90 ◦ C. [Ref. 32].
Pt were prepared on tungsten by constant potential electrolysis. EDX analysis of
the deposits indicated that Pt and Zn were distributed uniformly in the deposits.
The Pt content in the deposit decreases as the deposition potential approaches the
value where bulk deposition of Zn occurs. Increasing the Pt(II) concentration in
the liquid increases the Pt content in the deposits. As shown in Figure 5.4, XRD
results indicated the presence of crystalline Zn and amorphous PtZn. If Zn is
electrodeposited on a Pt substrate, the deposited Zn atoms interact with the Pt to
form Pt–Zn surface alloys.
5.4
Fabrication of a Porous Metal Surface by Electrochemical Alloying and De-alloying
Porous metals are of interest due to their potential applications in catalysis, fuel
cells, chemical sensors and so on. The fact that the electrodeposition of Zn on certain metals, M, can lead to surface alloys Mx Zn1–x and that the Zn in the alloy can
be subsequently removed by anodic stripping makes it possible to prepare porous
metal surfaces by an electrochemical alloying/de-alloying process. Some examples
including Pt, Au, and Ag have been demonstrated [33–35]. The formation of a
porous metal surface during electrochemical de-alloying can be accounted for with
the model described by Erlebacher [45]. For example, Figure 5.5 illustrates the dealloying of a Ag–Zn surface alloy. The process starts with selective dissolution of
the zinc atoms from the outermost Ag–Zn alloy surface, leaving behind the more
noble Ag atoms, which agglomerate to islands, leading to the formation of tiny
pits. The more pits formed, the more the original alloy is exposed to the electrolyte.
The selective dissolution of zinc atoms from the newly exposed Ag–Zn releases
more Ag atoms to the surface. These atoms diffuse to the Ag clusters left over
from dissolution of previous layers, continuing to leave more Ag–Zn exposed to
138 5 Electrodeposition of Alloys
Fig. 5.5 Plan-view SEM images of silver wire samples that have been
electrodeposited with 12.73 C cm−2 of zinc followed by de-alloying at
0.6 V. The amounts of the zinc that was de-alloyed are: (a) 1.59, (b)
4.77, (c) 9.55, and (d) 12.73 C cm−2 . The temperature was 150 ◦ C.
[Ref. 35].
electrolyte and resulting in increased pore size. Such selective dissolution of zinc
(roughening) and surface diffusion of Ag (agglomeration or smoothing) continues
as the de-alloying proceeds and an interconnected porous structure is formed. The
structure and morphology of the porous metal surface are affected by electrochemical variation of the composition and the thickness of the M–Zn surface alloys
(M = Pt, Au, or Ag). Higher deposition temperature favors the effective formation
of the M–Zn alloy and increases the thickness of the alloy. Higher de-alloying temperature enhances the surface diffusion of the metals and results in larger pores.
As both the deposition and de-alloying steps are performed in a single bath of
[EMIM]+ Cl− /ZnCl2 ionic liquid, the Zn(II) species consumed in the deposition
step returns to the ionic liquid during the de-alloying step, the composition of the
ionic liquid is essentially unchanged and thus can be re-used. The fabricated nanostructured platinum electrode was tested for the electro-oxidation of methanol. A
much higher current density is observed on the nanostructured platinum electrode
than on the polished platinum electrode, indicating that the former has a much
higher surface area.
The prospective application of the fabricated porous Ag electrode was tested
for the electrochemical reduction of chloroform. The advantageous catalytic effect
of the Ag electrode over a glassy carbon electrode is illustrated in Figure 5.6 [35]
which shows that while no significant current due to chloroform could be observed
5.5 Nb–Sn 139
Fig. 5.6 Typical cyclic voltammograms showing the electroreduction
of 50 mM chloroform in acetonitrile containing 1 M H2 O and 0.1 M
tetraethyl ammonium perchlorate (TEAP) as the supporting electrolyte
recorded at (a) a bare glassy carbon, (b) a polished Ag, and (c) a
porous Ag electrode. [Ref. 35].
at the glassy carbon electrode, appreciable current was observed at a polished Ag
electrode at about −1.3 V. The current density was further enhanced at the porous
Ag electrode compared with that on the polished Ag electrode. It has been demonstrated that a nanoporous Au surface can be prepared by electrochemical alloying/
dealloying from the [EMIM]+ Cl− /ZnCl2 liquid. It should be mentioned here that
according to IUPAC there are no nanoporous structures: <2 nm: microporous,
2–50 nm: mesoporous, > 50 nm: macroporous. However, even in the peer-reviewed
literature the expression “nanoporous” is increasingly employed for materials with
pores in the nanometer regime. Porous Au can for example be successfully functionalized with self-assembled monolayers of L-cysteine. Such functionalization
greatly improves the utility of the nanoporous gold, as was demonstrated in the
sensitive and selective determination of Cu(II) [34].
5.5
Nb–Sn
Koura et al. investigated the electrodeposition of Nb–Sn alloy in the ionic liquid
formed from a mixture of [EMIM]+ Cl− /SnCl2 /NbCl5 [36]. For the [EMIM]+ Cl− /
SnCl2 liquid, only one redox couple due to the cathodic deposition and anodic stripping of Sn was observed in the cyclic voltammogram. When NbCl5 was introduced
into the [EMIM]+ Cl− /SnCl2 liquid, new waves appeared in the voltammogram,
suggesting the codeposition of Nb–Sn alloy. The potential of these waves shifted
negatively on increasing the mole fraction of EMIC in the liquid, indicating that the
140 5 Electrodeposition of Alloys
Sn(II) and Nb(V) species changed their coordinations with the liquid composition.
Nb–Sn alloy samples were prepared by the potentiostatic method and analyzed. The
results showed that the Nb content in the alloy could be increased by increasing the
bath temperature to 160 ◦ C and increasing the NbCl5 content in the bath. However,
increasing the NbCl5 mole fraction in the bath also increased the viscosity of the
bath. Pulse electrolysis was found to be effective in increasing the Nb content in
the alloy. The maximum Nb content in the alloy was 60.8 wt% from constant potential electrolysis and 69.1 wt% from pulse electrolysis. XRD diffraction patterns
showed that the electrodeposits contained crystalline Sn and Nb3 Sn which is a
superconductor material.
5.6
Air- and Water-stable Ionic Liquids
For the electrodeposition of metals or alloys from air- and water-stable ionic liquids,
it is necessary first to dissolve the corresponding metal ions in the ionic liquid. Such
a dissolution process is made possible by introducing excess amounts of halide ions
(such as Cl− ) to form soluble metal-halide complex anions. Alternatively, the metal
is electrochemically oxidized in the ionic liquid to form the soluble salt such as
Sn(Tf2 N) in the trimethyl-n-hexylammonium [bis(trifluoromethyl)sulfonyl]amide
([TMHA]+ Tf2 N− ) ionic liquid.
5.6.1
Pd–Au, Pd–Ag, Pd–In
The electrodeposition of Pd–Au [37] and Pd–Ag [38] was investigated in the temperature range from 30 to 120 ◦ C in the air- and water-stable ionic liquid 1-ethyl3-methylimidazolium chloride-tetrafluoroborate ([EMIM]+ BF4 − ) containing Pd(II)
and Au(I), or Pd(II) and Ag(I), as well as excess chloride ions. Cyclic voltammograms indicated that the reduction of Au(I) and Ag(I) occurs at potentials less
negative than that of the Pd(II). Pd–Au and Pd–Ag alloys could be prepared by
galvanostatic or potentiostatic deposition on nickel substrates. EDX analysis of the
deposited alloys indicated that the Pd content of the alloy increased with increasing
Pd concentration in the bath and with lowering the deposition potential (or increasing the current density). XRD measurements indicated that the alloys were solid
solutions of Pd–Au and Pd–Ag. SEM images showed that increasing the deposition
temperature made the deposited alloys more compact.
The electrodeposition of Pd–In was investigated in the [EMIM]+ BF4 − ionic liquid
containing Pd(II) and In(III), and excess chloride ions [39]. The cyclic voltammograms shown in Figure 5.7 indicate that the reduction of Pd(II) occurs at a potential
less negative than that for the bulk deposition of In. However, UPD of In on Pd
occurs at the same potential as the deposition of Pd. Pd–In alloys can be prepared
within the indium UPD regime or at more negative potentials where overpotential
deposition of In occurs. As Figure 5.8 shows, in the UPD regime the In content
5.6 Air- and Water-stable Ionic Liquids 141
Fig. 5.7 Cyclic voltammograms of (a) 10 mM Pd(II), (b) 20 mM In(III)
and (c) 10 mM Pd(II) + 20 mM In(III) in a [EMIM]+ Cl− /BF4 − ionic
liquid at a GC electrode at 120 ◦ C. Scan rate = 100 mV s−1 . [Ref. 39].
Fig. 5.8 Variation of the Pd–In electrodeposit composition with deposition potential. The deposits were prepared in a [EMIM]+ Cl− /BF4 −
ionic liquid at 120 ◦ C containing: (△) 10 mM Pd(II) and 10 mM In(III);
() 10 mM Pd(II) and 40 mM In(III); (⋄) 40 mM Pd(II) and 10 mM
In(III). [Ref. 39].
142 5 Electrodeposition of Alloys
of the alloys was less dependent on the In(III) concentration of the bath because
the UPD of In on Pd is a slow process. In the mass-transport-limited potential
regime the alloy composition corresponds to the Pd(II)/In(III) composition in the
plating bath. SEM images showed that the alloys prepared within the UPD regime
appeared to have more compact and smooth morphologies.
5.6.2
In–Sn
The electrodeposition of In–Sn alloys from the Lewis basic [EMIM]+ BF4 − ionic
liquid containing 0.1 mol kg−1 InCl3 and 0.1 mol kg−1 SnCl2 was investigated by
Morimitsu et al. using cyclic voltammetry and potentiostatic electrolysis [40]. The
cyclic voltammograms indicated that the reduction of Sn(II) occurred at a potential
less negative than the reduction of In(III). Furthermore, the deposition of Sn greatly
reduced the overpotential required for the deposition of In, making the codeposition of InSn more feasible. The formation of at least two different phases during
the deposition was indicated by the presence of multiple anodic stripping waves.
Indium–tin alloy samples were prepared at a Pt flag electrode by constant potential
electrolysis. The obtained samples were analyzed by inductively coupled plasma
(ICP) atomic emission spectrometry and XRD measurements. The results showed
that the indium content increased to 29 atom% as the applied potential became
more negative. The InSn alloys could not only be obtained by the bulk deposition of
both In and Sn but could also be prepared by UPD of indium on the predeposited
tin. XRD measurements showed that the electrodeposited alloys were mixtures of
Sn and InSn4 . The crystallinity of the InSn4 phase in the electrodeposits is significantly affected by electrolysis temperature. Characteristic diffraction patterns of
InSn4 were not observed at room temperature but became evident when the alloy
samples were electrodeposited at 80 ◦ C. The fact that the maximum In content in
the deposits did not exceed 29 atom% indicates that the In–Sn alloy deposition is
not simply controlled by the mass-transport of individual Sn(II) and In(III) species.
Otherwise, the In content would be close to 50 atom% considering that the In(III)
and Sn(II) concentration ratio is 1:1 in the plating solution.
5.6.3
Cu–Sn
The formation of Cu–Sn alloy by galvanic contact deposition in the trimethyl-nhexylammonium [bis(trifluoromethyl)sulfonyl]amide ([TMHA]+ Tf2 N− ) ionic liquid at a temperature above 100 ◦ C has been demonstrated by Katase et al. [41]
Sn(II) was introduced into the liquid by dissolution of the Sn(Tf2 N) salt which has
a solubility of 0.2 mol dm−3 . In the plating cell, a copper sheet was used as the
cathodic substrate, a Sn sheet was used as the anode, and a Sn rod immersed in
the same solution was used as a quasi-reference electrode. On short-circuiting, the
Sn anode was oxidized to Sn(II) giving two electrons through external circuit to
5.6 Air- and Water-stable Ionic Liquids 143
the Cu cathode where Sn(II) was cathodically deposited. During deposition, the
current density decreased gradually to a steady-state current due to the depletion of
Sn(II) concentration in the vicinity of the cathode. The cathode potential remained
positive against Sn(II)/Sn, ensuring that the activity of the deposited Sn atoms is
lower than unity. After the conclusion of the deposition, silver-gray pinholes and
crack-free coatings were obtained. The XRD pattern of the Cu–Sn deposits obtained at 120, 130, and 140 ◦ C showed that the deposits are composed of crystalline
Cu6 Sn5 , Cu3 Sn, Cu10 Sn3 intermetallic phases. The amounts of Cu-rich phases and
the alloy thickness increased with increasing temperature due to the enhanced diffusion of deposited Sn atoms into the Cu substrate. The temperature dependence
of the thickness (x) of the deposits obtained by deposition for 72 h was studied and
Arrhenius behavior was observed in the plot of log x vs. T −1 . From the slope of this
plot, the apparent activation energy of the growth was estimated to be 58 kJ mol−1 .
5.6.4
Zn–Mn
The electrodeposition of Zn–Mn was investigated at 80 ◦ C in the hydrophobic tri-1butylmethylammonium bis((trifluoromethyl)sulfonyl)amide ([TBMA]+ Tf2 N− ) [46]
ionic liquid containing Zn(II) and Mn(II) species that were introduced into the
ionic liquid by anodic dissolution of the respective metal electrodes. Cyclic voltammograms indicated that the reduction of Zn(II) occurs at a potential less negative
than that of the Mn(II). Due to some kinetic limitations, which is a common
phenomenon in air- and water-stable ionic liquids, incomplete oxidation of Mn
electrodeposits was observed in this system. The current efficiency of Mn electrodeposition in this ionic liquid approaches 100%, which is a great improvement
compared to the results obtained in aqueous solution (20–70%). Electrodeposition
of Zn–Mn alloy coatings has never been carried out in chloroaluminate ionic liquid
because of the unavoidable codeposition of Mn and Al.
Coatings containing Zn, Mn or Zn–Mn were obtained by controlled-potential
electrolysis and analyzed by SEM, EDX and XRD. It is very interesting that the
reduction wave of Zn(II) disappeared when the ionic liquid contained both Zn(II)
and Mn(II) species; this is illustrated by the CVs shown in Figure 5.9. The reason
is still not clear but compact and adherent Zn–Mn alloy deposits of various compositions can be obtained and the Mn/Zn ratio of these alloys depended almost
completely on the Mn(II)/Zn(II) concentration ratio in the ionic liquid. The SEM
images shown in Figure 5.10 demonstrate that the Zn–Mn alloy deposits were very
smooth and the grain size increased with increasing concentration of Mn. The XRD
results indicate that the Zn–Mn alloy deposits obtained from the [TBMA]+ Tf2 N−
were metallic glasses or amorphous. Potentiodynamic anodic polarization experiments in deaerated aqueous NaCl revealed that the addition of Mn up to 50 atom%
improves the corrosion resistance of Zn. However, the addition of Mn beyond this
amount decreases the corrosion resistance of the Zn–Mn alloy.
144 5 Electrodeposition of Alloys
Fig. 5.9 Staircase cyclic voltammograms of (—) 0.3 M Zn(II), (. . .)
0.2 M Mn(II) and (—) a mixture of 0.16 M Zn(II) + 0.1 M Mn(II)
recorded at a W electrode in [TBMA]+ Tf2 N− ionic liquid. Temperature,
80 ◦ C. Scan rate, 50 mV s−1 . [Ref. 46].
Fig. 5.10 SEM micrographs of pure Mn, pure Zn and various compositions of Zn–Mn alloys. The composition of the alloy coatings are
shown as atomic ratios on each plot. The magnification of each micrograph is 5000×. [Ref. 46].
References
5.7
Summary
In this chapter some results on the electrodeposition of alloys from ionic liquids are
summarized. Many fundamental studies have been performed in chloroaluminate
first generation ionic liquids but the number of studies employing air- and waterstable ionic liquids rather than the chloroaluminates is increasing. Currently, new
ionic liquids with better electrochemical properties are being developed. For example, Abbott et al. [47] have prepared a series of ionic liquids by mixing commercially
available low-cost choline chloride and MCl2 (M = Zn, Sn) or urea and demonstrated that these ILs are good media for electrodeposition for pure metals (see
Chapter 4.3). It can be expected that in the near future, the electrodeposition of alloys from ILs may become available for industrial applications. Furthermore, due to
their variety, their wide electrochemical and thermal windows air- and water-stable
ionic liquids have unprecedented prospects for electrodeposition.
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6
Electrodeposition of Semiconductors in Ionic Liquids
Natalia Borisenko, Sherif Zein El Abedin, and Frank Endres
In this chapter we report on the electrodeposition of semiconductors in ionic liquids. It is shown that ionic liquids are, due to their extraordinary physicochemical
properties, well suited as a solvent medium for the electrodeposition of elemental
semiconductors (like Si and Ge), their mixtures (Six Ge1−x ) and compound semiconductors (GaAs, AlSb, InSb, ZnTe, CdTe, CuInSe2 , etc.).
6.1
Introduction
There is a wide variety of applications for elemental and compound semiconductors. Compound semiconductor thin films, for example, are used in many optoelectronic devices like photon detectors, light emitting diodes (LED), photovoltaics
and lasers. Cadmium based II–VI semiconducting thin films, such as CdTe, CdSe,
CdS, and CdHgTe, display a variety of band gaps and lattice constants, which
make them interesting for optoelectronic applications [1, 2]. The family of III–V
compound semiconductors, in particular antimony-based semiconductors (AlSb,
GaSb, and InSb), are of great interest as barrier materials in high-speed electronics and long-wavelength optoelectronic devices [3, 4]. Au–Cd alloys are employed
as ohmic contacts with semiconducting films and may provide additional doping
in these materials [5]. Ternary compound semiconductors, for example CuInSe2
(CIS), are promising materials for thin film photovoltaic applications due to their
stability, direct energy band gap and high absorption coefficient [6]. Elemental
semiconductors, such as Si and Ge, are widely used as wafer material for different
electronic applications, and junctions of n- and p-doped Si are still interesting for
photovoltaic applications. Ge quantum dots made by molecular beam techniques
under ultrahigh vacuum (UHV) conditions have interesting optical properties. For
example, Ge quantum dots on Si(111) show a photoluminescence at about 1 eV
[7]. Silicon nanocrystals embedded in a SiO2 matrix have been discussed for the
development of nanoscale silicon-based lasers [8].
Electrochemical deposition is one of the main fields in electrochemistry, both
in industrial processes and in fundamental research. It has been applied to make
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
148 6 Electrodeposition of Semiconductors in Ionic Liquids
semiconductors by electrochemical means for over 30 years and a good general
overview can be found in Ref. [9]. However, standard industrial procedures for
semiconductor electrodeposition are rare. Many studies on the deposition of semiconductors and their characterization have been performed in different solutions
such as aqueous media, organic solutions and molten salts. In fundamental research, most of the investigations have been performed by molecular beam epitaxy
(MBE) under UHV conditions. In industrial processes, chemical or physical vapor deposition methods are preferred. The obtained layers are of a high quality,
but such processes are cost-intensive thus making the deposits quite expensive.
Therefore electrodeposition of semiconductors would be technically interesting as
electrodeposition is, in contrast to UHV techniques, a comparatively cheap process;
only the imagination of the user limits the size of the objects onto which the semiconductor can be deposited. Recently Stickney et al. demonstrated that compound
semiconductors like CdTe, CdSe, CdS or HgSe can be electrodeposited in aqueous media by the electrochemical atomic layer epitaxy (ECALE) method [10–15].
The desired semiconductor is made by the subsequent layer-by-layer growth of the
respective elements. The direct electrodeposition of compound semiconductors in
one step is often difficult for kinetic reasons. At room temperature the elements
are codeposited in varying amounts together with the desired semiconductor and
variation in temperature can strongly affect the quality of the films [16]. The electrodeposition of III–V compound semiconductors, like InSb and InAs, has also
been investigated in aqueous solutions [17–21]. Unfortunately, elemental semiconductors like silicon, germanium and their mixtures (Six Ge1−x ) cannot be obtained
in aqueous solutions as the deposition is strongly disturbed by hydrogen evolution.
Instead of Si deposition there would only be hydrogen evolution in aqueous media. Macroscopically thick and amorphous germanium films can be obtained from
GeI4 dissolved in propylene glycol at elevated temperature under galvanostatic conditions [22]. Szekely et al. showed that a roughly 130 µm thick germanium layer
was electrodeposited from GeCl4 dissolved in propylene glycol at 59 ◦ C and a constant current density of 0.4 mA cm−2 . Unfortunately, the current efficiency in these
systems is only about 1% [23]. Ge can also be electrodeposited in high-temperature
molten salts [24]. There have been several attempts to electrodeposit silicon in organic solvents [25–27], and smooth and uniform silicon deposits up to 0.25 µm thick
were described. However, Auger electron spectroscopy analysis of such deposits
evidenced oxygen content and it was not clear whether the deposit was oxidized
during the deposition process or if it was a consequence of an open porosity in the
film. Of course, silicon can also be electrodeposited in high-temperature molten
salts [28]. Recently the electrodeposition of silicon from its halides in non-aqueous
solutions was investigated [29]. These authors also reported strong oxidation of the
electrochemically made silicon. The electrodeposition of semiconductors in ionic
liquids is a comparably young research area. Ionic liquids are, due to their extraordinary physical properties, (see Chapter 3) very interesting as electrolytic media.
They are good solvents for a variety of both organic and inorganic materials (depending on the liquid), they are immiscible with a variety of organic solvents and, in
part, immiscible with water, they are nonvolatile (in most cases) and, hence, can be
6.3 Indium Antimonide 149
used even in ultra-high-vacuum systems. Ionic liquids exhibit wide electrochemical
windows of up to 6 V, and wide thermal windows of 300–400 ◦ C. Depending on
the cation/anion combination, the temperature can be varied over several hundred
degrees, so that kinetic barriers in semiconductor compound formation could be
overcome. The unique properties of ionic liquids open the opportunity to apply
them as solvents for electrodeposition of semiconductors, which are hardly (or
even not at all) accessible from aqueous solutions, such as Si, Ge, GaAs, etc. The
electrodeposition of GaSb, InP, InSb, InSe and ternary semiconductors in ionic
liquids is also interesting, especially at elevated temperatures as kinetic barriers
in compound formation might be more easily overcome at temperatures around
200–300 ◦ C. In this chapter we summarize recently published results on the electrodeposition of semiconductors in different ionic liquids.
6.2
Gallium Arsenide
GaAs is a well-known III–V compound semiconductor with a direct band gap
of 1.43 eV at room temperature. Due to the high mobility of the charge carriers,
GaAs-based electronic devices can operate at higher frequencies than equivalent Si
devices, resulting in faster electronics, that makes these semiconductor interesting
for many optoelectronic applications including semiconductor lasers, LEDs and
solar cells. The direct electrodeposition of GaAs in ionic liquids has been studied
by two groups. In 1986 Wicelinski and Gale showed that GaAs can, in principle, be
electrodeposited from GaCl3 and AsCl3 at 40 ◦ C in the Lewis acid chloroaluminate
ionic liquid composed of AlCl3 and 1-butylpyridinium chloride [30]. The authors report that Al codeposition occurs in the underpotential deposition regime on the Ga
surface. In order to minimize Al contamination Carpenter and Verbrugge employed
a chlorogallate ionic liquid [31, 32]. It was shown that GaAs film can be obtained at
room temperature in the Lewis basic GaCl3 /1-methyl-3-ethylimidazolium chloride
ionic liquid, to which AsCl3 was added. The quality of the deposit can be improved by
thermal annealing, which makes this method promising, in principle, for the electrodeposition of GaAs-based compound semiconductors. However, GaCl3 -based
ionic liquids are extremely aggressive and AsCl3 is extremely poisonous so that
such liquids would involve enormous security issues.
6.3
Indium Antimonide
InSb is an important compound semiconductor of the III–V family for optoelectronic purposes. At room temperature the semiconductor has a direct band gap
of 0.17 eV and a high mobility of charge carriers. Similar to GaAs, it was reported
that InSb can be directly electrodeposited at 45 ◦ C in the Lewis basic chloroindate ionic liquid InCl3 /1-methyl-3-ethylimidazolium chloride, to which SbCl3 was
150 6 Electrodeposition of Semiconductors in Ionic Liquids
added before the deposition [33]. The In/Sb ratio in the deposit is strongly dependent on the applied electrode potential. Consequently elemental Sb and In can
also be found in the films. Recently Sun et al. employed the Lewis basic 1-ethyl3-methylimidazolium chloride/tetrafluoroborate ionic liquid containing InCl3 and
SbCl3 for InSb deposition [34]. The composition of the InSb films also depends
strongly on the deposition potential. However, the crystallinity of the deposits is
strongly improved by increasing the deposition temperature, and polycrystalline
InSb can be directly electrodeposited at 120 ◦ C without additional annealing. The
band gap was determined by absorption spectroscopy to be 0.2 eV. Although the
quality of the deposits depends on the absolute concentrations of In(III) and Sb(III)
species and although individual indium and antimony crystals can be found in the
films this result proves that the wide thermal windows of ionic liquids and the wide
electrochemical windows allow one to find parameters under which the compound
semiconductors can be made directly.
6.4
Aluminum Antimonide
The binary semiconductor AlSb is, like GaAs and InSb, one of the III–V semiconductors. In particular, AlSb is a highly efficient solar cell material. It exhibits
a direct band gap of 2.5 eV and an indirect band gap of 1.2 eV at room temperature. The electrodeposition of AlSb was investigated at room temperature in the
Lewis neutral ionic liquid AlCl3 /1-butyl-3-methyl-imidazolium chloride [35, 36].
A liquid containing Sb(III) was prepared by addition of SbCl3 /1-butyl-3-methylimidazolium chloride to the chloroaluminate ionic liquid. The electrodeposition of
AlSb was investigated by in situ scanning tunneling microscopy (STM) and in situ
scanning tunneling spectroscopy (STS). A band gap of about 2.0 eV was obtained.
As in the case of GaAs and InSb, codeposition of the elements occurs, furthermore
strong doping effects by the elements occur if the deposition is performed at electrode potentials away from the compound deposition potential. In future studies
it should be investigated whether deposition at elevated temperatures (∼ 200 ◦ C)
allows better control of AlSb-stoichiometry. Furthermore the use of air- and water
stable ionic liquids might lead to more reproducible results.
6.5
Zinc Telluride
ZnTe is usually applied in switching devices and in solar cells. It is one of the II–VI
compound semiconductors with a direct band gap of 2.3 eV at room temperature.
The electrodeposition of ZnTe was investigated by Sun et al. in the Lewis basic
ZnCl2 /1-ethyl-3-methylimidazolium ionic liquid containing propylene carbonate
as a cosolvent at 40 ◦ C [37]. 8-Quinolinol was added to the solution to shift the
reduction of Te(IV) to more negative potential, thus facilitating the codeposition.
The composition of the ZnTe deposits is dependent on the deposition potential and
6.7 Germanium
on the concentration of Te(IV) in the solution. After thermal annealing the band
gap was determined by UV/vis absorption spectroscopy to be 2.3 eV, which is in
good agreement with ZnTe films made by others methods.
6.6
Cadmium Telluride
CdTe, a II–VI compound semiconductor with a direct band gap of 1.44 eV at room
temperature, is, from its physical properties, a promising photovoltaic material.
The electrodeposition of CdTe in ionic liquid was published recently by Sun et al.
[38]. They were able to show that the semiconductor can be electrodeposited at elevated temperature (above 120 ◦ C) in the Lewis basic 1-ethyl-3-methylimidazolium
chloride/tetrafluoroborate ionic liquid containing CdCl2 and TeCl4 . CdTe films
were obtained by the underpotential deposition (UPD) of Cd on the deposited
Te. The deposit composition was independent of the deposition potential within
the Cd UPD regime. The crystallinity of the deposits is improved by increasing the
deposition temperature, which again demonstrates the high potential of the wide
thermal windows of ionic liquids for compound electrodeposition.
6.7
Germanium
Germanium is an elemental semiconductor with an indirect band gap of 0.67 eV at
room temperature in the microcrystalline phase. Its crystal structure is determined
by the tetrahedral symmetry of Ge in the crystalline phase. An interesting aspect
is that Ge nanoparticles with diameters of only a few nanometers exhibit a sizedependent photoluminescence. Nanocrystalline Ge is a direct semiconductor and
it is regarded today as a promising candidate for optical sensors. However, almost
all studies on the production and characterization of Ge nanocrystals or quantum
dots hitherto have been performed under UHV conditions, which require a high
instrumental effort for a possible nanotechnological process. An electrochemical
process would be quite interesting as there is, in principle, no limit on surface area
and geometry, furthermore electrochemical experiments are comparatively easy to
perform.
The electrodeposition of germanium in ionic liquids was primarily investigated
in our group. The original aim was to find out if germanium can be made by electrochemical means at all. In situ STM and in situ STS techniques were employed
for this purpose. These techniques allow one to investigate the initial stages of
the semiconductor electrodeposition and to understand the deposition process on
the nanometer scale. The electrodeposition of Ge was studied at room temperature from GeX4 (where X = I, Br, Cl) on Au(111) in the ionic liquid 1-butyl-3methylimidazolium hexafluorophospate ([BMIM]PF6 ) [39–42]. At the time of the
experiments this ionic liquid was one of a few which could easily be prepared
151
152 6 Electrodeposition of Semiconductors in Ionic Liquids
Fig. 6.1 CV of pure [BMIM]PF6 on Au(111). The scan rate is
1 mV s−1 . Mainly capacitive currents flow between the cathodic and
the anodic limits. The electrochemical window is a little more than
4 V.
with water levels below 20 ppm and was, therefore, the best choice for such investigations. All experiments have to be performed under inert gas conditions as
the germanium halides react rapidly with water. For comparison purposes the
cyclic voltammograms were calibrated versus the overpotential deposition (OPD)
of germanium.
The electrochemical window of dry [BMIM]PF6 on Au(111) is a little more than
4 V, as can be seen in the cyclic voltammogram (CV) of Figure 6.1. In dry and
high purity liquids only capacitive currents flow between the cathodic and anodic
limits. At the cathodic limit the organic cation is irreversibly reduced and STM
pictures in this potential regime show that a less defined deposit is formed on
the electrode surface. This decomposition product is dissolved in the liquid when
the potentiostat is switched off, turning the originally colorless liquid to purple.
At the anodic limit gold oxidation occurs which, in the initial state, can also be
probed by in situ STM.
In the following the focus is on the electrodeposition of Ge from GeCl4 . The
electrodeposition is quite similar from all the halides but the oxidation of the gold
substrate was least severe in the case of GeCl4 .
If [BMIM]PF6 ionic liquid is saturated with GeCl4 (Figure 6.2), two main reduction processes (P1 and P2 ) are observed in the cathodic regime [42]. The first
reduction peak, with a minimum at +500 mV vs. Ge (P1 ) is attributed to the reduction of Ge(IV) to Ge(II). At potentials below 0 mV (P2 ) the bulk deposition of Ge
from Ge(II) sets in, as can be seen with the naked eye. The rising cathodic current
at about −1000 mV vs. Ge is attributed to the irreversible reduction of the organic
cation. If only P1 is passed, an oxidation process is not observed. If Ge deposition
is performed an oxidation peak at ∼1000 mV is observed, which means that this
peak must be correlated to Ge electrooxidation. A series of oxidation peaks above
+1500 mV is also observed if the electrode potential is cycled between +1000 and
6.7 Germanium
Fig. 6.2 CV of GeCl4 saturated in dry
[BMIM]PF6 on Au(111). The scan rate is
1 mV s−1 . Upon addition of GeCl4 two
mainly irreversible reduction peaks (P1
and P2 ) are observed. The process P1 is
correlated with the reduction of Ge(IV) to
Ge(II). At P2 the bulk deposition of Ge is
observed. The irreversible reduction of the
imidazolium ion starts at −1000 mV (P3 ).
A strong oxidation peak at +1000 mV is
attributed to partial Ge dissolution. At electrode potentials above +1500 mV gold oxidation sets in.
+3000 mV vs. Ge, thus avoiding GeCl4 reduction. Therefore these processes are
due to the oxidation of the gold substrate which is difficult to probe with in situ
STM.
At the open circuit potential, OCP, (+1200 mV vs. Ge) a typical Au(111) surface
with step heights of 250 pm is probed (Figure 6.3(a)). When the electrode potential
is reduced to +1000 mV, the step edges become quite obviously decorated, and this
decoration is not observed in the pure ionic liquid. First, small two-dimensional
islands with heights between 100 and 150 pm appear at about +950 mV (Figure
6.3(b)). These islands can be reversibly stripped from the surface. When the electrode potential is further reduced to +750 mV, islands with an average height of
250 pm form. If the potential is set back to +1200 mV the islands dissolve and tiny
holes with a depth of about 100 pm appear. The holes completely heal in a few
minutes and a flat terrace-like gold structure is again obtained. Between +300 mV
and 0 mV a rough but completely closed layer with a maximum thickness of 300 pm
forms (Figure 6.3(c)). The in situ I/U tunnelling spectrum clearly reveals the metallic behavior of this layer but the tunnelling barrier is much higher than for a pure
gold substrate at OCP (Figure 6.3(d)), where even a linear increase in tunneling
current can occur. It is therefore likely that, in the UPD regime, a surface alloying
between Au and Ge takes place so that the UPD germanium gets a more noble
metallic character. The deposit does not grow further if the electrode potential of
the STM tip is set to sufficiently high values. However, if the tip potential is kept
close to the bulk deposition potential of Ge, cluster agglomerates composed of
small clusters of only some nanometers in diameter grow on the surface, probably
due to a jump from the tip to the sample (Figure. 6.4(a)). In situ I/U tunnelling
153
154 6 Electrodeposition of Semiconductors in Ionic Liquids
Fig. 6.3 The series of STM pictures
shows the UPD of Ge on Au(111) in
GeCl4 /[BMIM]PF6 . At +1200 mV vs. Ge
(OCP) a typical Au(111) structure with step
height 250 pm is observed (a). When the
electrode potential is reduced to +950 mV
small two-dimensional islands with heights
between 100 and 150 pm appear (b). Be-
tween +300 and 0 mV a completely closed
but rough layer with an average height of
300 pm forms (c). The layer shows metallic
behavior but the tunneling barrier is higher
than that for the pure Au substrate at OCP.
Most likely an alloying between Ge and Au
occurs (d).
spectroscopy on different sites of these clusters shows a bias range of about 500 mV
with almost zero tunneling current (Figure 6.4(b)). At −50 mV vs. Ge (slightly in the
OPD regime) islands/crystallites with diameters of 50 nm and heights of 5–10 nm
are observed (Figure 6.5(a)). The band gap of these individual crystals is 0.7 ±
0.1 eV, which is typical for intrinsic bulk germanium at room temperature. If the
electrode potential is further reduced to −200 mV vs. Ge about 100 nm thick deposits with a band gap of 0.7 ± 0.1 eV form (Figure 6.5(b, c)). The film is composed
of nanocrystals and tongue-like germanium islands. An XPS study of a micrometer
6.8 Silicon 155
Fig. 6.4 If the bias is only +200 mV small
cluster agglomerates grow on the surface
at +100 mV vs. Ge (a). A higher resolution
image (inset in (a)) shows that the cluster
agglomerates are composed of small clus-
ters with diameters of only a few nm. The
in situ I/U tunneling spectrum reveals that
these islands exhibit a bias range of about
500 mV with almoust zero tunneling current
(b).
thick Ge film made from GeX4 shows that indeed elemental Ge was obtained [43].
The electrochemically made Ge is, however, subject to some attack by environmental oxygen (Figure 6.6(a, b)). We would like to mention a further interesting effect
which we observed when GeI4 microcrystals were directly reduced in the ionic
liquid, instead of microcrystals Ge crystals with sizes around 100–200 nm were
obtained (Figure 6.7). Most likely the confined diffusion space between crystal and
electrode surface led to comparatively small crystals.
These results show that not only Ge layers but also Ge nanocrystals can be made
electrochemically in ionic liquids by adjusting the experimental parameters.
6.8
Silicon
Currently silicon is still one of the most important semiconductors as it is the basis
of any computer chip. It exhibits an indirect band gap of 1.1 eV at room temperature
in the microcrystalline phase. Similar to Ge, silicon nanoparticles show a sizedependent photoluminescence. It was reported by Katayama et al. that a thin Si
layer can be electrodeposited in 1-ethyl-3-methylimidazolium hexafluorosilicate at
90 ◦ C [44]. However, upon exposure to air the deposit reacted completely to SiO2 ,
which makes it difficult to decide whether the deposit was semiconducting or not.
Recently, we showed for the first time that silicon can be well electrodeposited
from SiCl4 in the air and water stable ionic liquid 1-butyl-1-methylpyrrolidinium
bis(trifluoromethylsulfonyl)amide ([BMP]Tf2 N) [45, 46]. This ionic liquid can be
156 6 Electrodeposition of Semiconductors in Ionic Liquids
Fig. 6.5 The STM images represent the OPD of Ge on Au(111) in
GeCl4 /[BMIM]PF6 . At −50 mV vs. Ge clusters with diameters of 50 nm and heights of
5–10 nm are observed (a). When the electrode potential is set to −200 mV an about
100 nm thick Ge film, composed of small
nanoclusters and tongue-like islands, is
probed (b). The in situ tunneling spectrum
also shows semiconducting behavior with a
symmetrical band gap of 0.7 ± 0.1 eV, typical for elemental Ge (c).
dried to water contents below 1 ppm. Similar to Ge, the experiments were performed
under inert gas conditions. The reversible ferrocene/ferrocinium (Fc/Fc+ ) redox
couple was employed as a reference electrode.
The liquid itself exhibits on Au(111) an electrochemical window slightly more
than 5 V (Figure 6.8). At the cathodic limit a series of peaks (C1 –C3 ) is observed
prior to the irreversible reduction of organic cation at −3200 mV vs. Fc/Fc+ . At the
anodic limit at +2000 mV vs. Fc/Fc+ gold oxidation sets in. The oxidation processes
A3 and A4 are only observed if the reduction processes C3 and C4 have been passed.
For the peaks C1 and C2 the respective oxidation processes are missing. It was
shown that the peaks C1 to C3 are correlated to the irreversible breakdown of the
Tf2 N ion [47, 48].
If SiCl4 (0.1 mol l−1 ) is dissolved in [BMP]Tf2 N two main reduction processes
preceding the reduction of the organic cation at C3 are observed (Figure 6.9). As
6.8 Silicon 157
Fig. 6.6 XPS spectra of germanium deposits made potentiostatically at −500 mV
vs. Ge for 6 h from GeBr4 (a) and from
GeCl4 (b) on Au(111). For both deposits
Auger peaks with binding energies of ∼105
and ∼109 eV are found, which can be attributed to elemental Ge. A transition at
114–115 eV is likely due to a chemical shift
as a higher oxidation state of Ge leads to
transitions at higer energies. For Ge from
GeBr4 , the 3p transitions with binding energies of ∼122 and ∼126 eV are found, as expected for elemental Ge. For Ge from GeCl4
besides the 3p transitions for elemental Ge,
two more transitions shifted by 4–5 eV with
respect to 3p are obtained, indicating the
presence of Ge(IV) in the deposit.
C2 only appears when SiCl4 is in the solution, this peak must be correlated to the
bulk deposition of silicon. It is interesting that the reduction of the organic cation
on a silicon surface is strongly hindered. Furthermore, the decomposition of the
ionic liquid might passivate the Si surface as the deposition of Si can also start in
the anodic branch at C* . Since the peak C1 was not observed on HOPG [45] and
as there is no deposition at this potential regime, the process might be correlated
with ionic liquid breakdown and surface restructuring/reconstruction. A further
explanation might be the formation of low valent silicon species in the solution.
158 6 Electrodeposition of Semiconductors in Ionic Liquids
Fig. 6.7 SEM image of a germanium deposit, obtained upon electrodeposition under a GeI4 crystal at −1000 mV vs. Ge. The surface
is completely covered by a thin germanium layer. A collection of nanoclusters with a grain size of about 100 mn is obtained.
The broad oxidation processes at E > 0 mV are partly due to Si oxidation, gold
oxidation and oxidation of cation reduction products.
STM images in Figure 6.10 show the growth of silicon in SiCl4
(0.1 mol l−1 )/[BMP]Tf2 N on Au(111). Between −300 mV (OCP) and −600 mV vs.
Fc/Fc+ , instead of flat terraces, which would be typical for Au(111), the terraces
show a worm-like restructuring (Figure 6.10(a)), which is less clear if SiCl4 is in
Fig. 6.8 CV of dry ultrapure [BMP]Tf2 N
ionic liquid on Au(111) at a scan rate of
10 mV s−1 . The electrochemical window is
about 5 V, limited by the reduction of the
organic cation at C4 and by gold oxidation
at +2000 mV vs. Fc/Fc+ . Within the electrochemical window a series of reduction
peaks C1 –C3 is observed, due to an irreversible breakdown of the Tf2 N anion.
6.8 Silicon 159
Fig. 6.9 CV of SiCl4 (0.1 M) in [BMP]Tf2 N
ionic liquid on Au(111) at a scan rate of
10 mV s−1 . The bulk deposition of silicon
starts at C2 . The reduction of the organic
cation (C3 ) is hindered on the silicon surface. The first reduction process C1 is not
correlated with a definite surface process. In
the reverse scan further silicon deposition is
observed (C* ). Strong oxidation processes
at a potential of more positive than 0 mV
are correlated to the oxidation of silicon,
gold and cation reduction products.
the solution. The step height between the terraces is still 250 ± 30 pm. The defects/vacancy islands are one monolayer deep and the width is about 10–20 nm.
When the electrode potential is reduced to more negative values the number of
these vacancy islands is strongly decreased. Therefore the allocation of the reduction process C1 in Figure 6.9 to a definite surface process is difficult. Between
−1100 mV and −1600 mV a typical terrace-like Au(111) surface with a step height
of 250 pm is probed (Figure 6.10(b)). At −1700 mV small silicon islands with diameters of less than 50 nm and heights between 150 and 450 pm start to grow (Figure
6.10(c)). With time, the number of these islands very slowly increases and they grow
slightly in height. After 1 h a completely closed thin layer of silicon whose height
is in the nanometer regime forms (Figure 6.10(d)). Some small islands rise above
this thin layer. The in situ I/U spectrum reveals typical semiconducting behavior
with a band gap of 1.1 ± 0.2 eV, which is in excellent agreement with literature data
for microcrystalline silicon in the bulk phase at room temperature (1.1 eV). When
the electrode potential is further reduced the islands grow above the surface and
merge laterally, resulting in silicon agglomerates (Figure 6.11(a)). These structures
can be as high as 10 nm, their width can reach 30 nm and they also exhibit a band
gap of 1.1 eV (Figure 6.11(b)).
SEM pictures show that a very thin layer of small clusters/crystallites forms first
on the gold surface, followed by crystalline agglomerates that finally leads to a
500–1000 nm thick silicon layer (Figure 6.12).
In order to get more ex situ information on the electrochemically made silicon
we performed a detailed XPS study. For this purpose the Si was made inside the
inert gas glove box. The deposit was subsequently purified by rinsing in isopropanol
inside the glove box and transferred via a transport chamber to the XPS device. Thus
160 6 Electrodeposition of Semiconductors in Ionic Liquids
Fig. 6.10 The sequence of STM pictures
shows the nanoscale growth of Si in the
SiCl4 (0.1 M)/[BMP]Tf2 N on Au(111) probed
by in situ STM. Instead of a typical flat
Au(111) surface, a worm-like structure is
probed at −300 mV vs. Fc/Fc+ (OCP) (a).
When the electrode potential is reduced a
typical Au(111) terrace-like surface with a
step height of 250 pm is observed (b). If
the electrode potential is set to −1700 mV,
small silicon islands with a width of less
than 60 nm and 150–450 pm in height start
to grow (c). After 1 h an about 5 nm high
close silicon layer forms (d) close to the
equilibrium potential.
the sample was never in air. As an example Figure 6.13 shows the XPS spectrum of
the Si 2p peak. Besides the elemental peak at 101.3 eV there is strong evidence for
SiOX at 104.4 eV. As discussed elsewhere [49] silicon is deposited electrochemically,
but even in an inert gas glove box with an oxygen concentration as low as 1 ppm it is
attacked by oxygen at the surface. The XPS study shows undoubtedly that elemental
silicon can be electrodeposited in ionic liquids.
6.9
Grey Selenium
Grey selenium exhibits both photovoltaic and photoconductive properties, which
make it useful in the production of photocells and solar cells. Moreover,
6.9 Grey Selenium
Fig. 6.11 When the electrode potential is reduced to −1800 mV vs.
Fc/Fc+ silicon islands grow above the surface and merge laterally
leading to agglomerates (a). The in situ I/U tunneling spectrum shows
that both the layer and the islands exhibit a band gap of 1.1 ± 0.2 eV,
typical for mycrocrystaline semiconducting Si (b).
Se-containing compound semiconductors, such as InSe, CdSe or CuInSe2 (CIS)
have many optoelectronic applications, including advanced solar cells, IR detectors
and solid-state lasers. The CIS solar cells, especially, are very promising as a highly
efficient power supply. The electrodeposition of selenium has been intensively investigated in aqueous solutions [50–53]. However, the exclusive electrodeposition
of grey selenium in aqueous solutions is pretty difficult as at temperatures below
100 ◦ C red and black selenium, which are both insulators, grow in certain amounts
together with the desired grey phase. Consequently, in the technical process selenium is applied by gas phase condensation. Thermodynamically a phase transition
Fig. 6.12 SEM micrographs of electrodeposited Si, made potentiostatically at
−2700 mV vs. Fc/Fc+ . The surface is completely covered by a thin silicon layer composed of individual clusters. Globular
50–150 nm wide crystallites consisting of
many tiny crystals grow above this layer (a).
A 500 nm thick silicon layer consists of coherent spherical crystallites (b).
161
162 6 Electrodeposition of Semiconductors in Ionic Liquids
Fig. 6.13 High resolution XPS spectrum of
a nanoscale silicon deposit made potentiostatically at −2200 mV vs. Pt quasi-reference
for 2 h on a stainless steel substrate. (Dots:
original data obtained from the measurement. Solid lines: fitted data. Dotted line:
sum of both fitted contributions.) A Si 2p
peak consists of two different contributions
with binding energies of 101.3 and 104.4 eV,
which can be attributed to elemental Si and
SiOX , respectivelly.
from amorphous red to crystalline grey selenium occurs at about 80 ◦ C. But even
at 100 ◦ C in aqueous solutions the deposit does not contain solely the grey phase
of Se. Obviously grey selenium can only be electrodeposited at elevated temperatures of more than 100 ◦ C, which cannot be achieved in aqueous solutions. The
direct electrodeposition of grey selenium can be performed in ionic liquids as the
deposition process can be realized at elevated temperatures due to the high thermal
stability and low vapor pressure of most of the ionic liquids. In a recent paper we
reported that grey selenium can be well electrodeposited from SeCl4 in [BMP]Tf2 N
ionic liquid [54]. All experiments were performed under an inert gas atmosphere.
Unfortunately, we could not employ ferrocene as an internal reference, due to
the complexity of voltammograms and the need to exclude any interference with
ferrocene. Therefore, a Pt-wire had to be used as a quasi-reference electrode. In
our experience Pt has a sufficiently stable electrode potential for a while under the
applied conditions.
The cyclic voltammogram of [BMP]Tf2 N containing 0.1 mol l−1 SeCl4 on a platinum substrate at 25 ◦ C is presented in Figure 6.14(a). At a potential of −750 mV
vs. Pt a dark red deposit forms on the electrode surface, obviously passivating it.
It is likely that this peak is correlated to the reduction of Se(IV) to the red phase
of elemental Se. It cannot be excluded that the black phase is also formed. If the
6.9 Grey Selenium
Fig. 6.14 CVs of SeCl4 (0.1 M) in
[BMP]Tf2 N ionic liquid on Pt substrate
at 25 ◦ C (a) and 150 ◦ C (b). Scan rate is
10 mV s−1 . At about −750 mV the deposition of red amorphous Se occurs. The red
colour of the deposit turns to grey at C4 .
The deposited Se is partly dissolved at A3 .
The shoulders C1 and C2 and their anodic
counterparts A1 and A2 might be related to
two different UPD processes. The pair C5
and A5 are likely to be due to the reduction
of the deposited Se to Se2− .
temperature is increased to 150 ◦ C (Figure 6.14(b)) five cathodic processes C1 –C5
and their corresponding anodic counterparts A1 –A5 are observed. C3 and C4 are
correlated with the electrodeposition of selenium. Visually, first a red deposit forms
at C3 , which turns to a grey colour at C4 . The peaks C1 and C2 are presumably correlated with different UPD processes. The peaks C5 and A5 are likely to be associated
with the further reduction of the deposited selenium to Se2− as the selenium film
can disappear completely at this electrode potential.
As one example the SEM picture in Figure 6.15 shows an electrodeposited Se
layer made potentiostatically at −1100 mV vs. Pt at 150 ◦ C. The XRD pattern of
Fig. 6.15 SEM image of Se layer, made potentiostatically at
−1100 mV vs. Pt at 150 ◦ C. The XRD pattern of this layer shows that
crystalline grey Se is electrodeposited.
163
164 6 Electrodeposition of Semiconductors in Ionic Liquids
the electrodeposit reveals the characteristic peaks of the crystalline grey selenium.
From XRD and SEM alone it cannot be excluded that some red or black Se also
form, thermodynamically, however, it is unlikely at these temperatures.
In our opinion the electrodeposition of selenium is quite promising for a variety
of applications. For example, the possibility to deposit grey selenium, indium,
and copper in one ionic liquid at variable temperatures might be regarded as the
first step in making selenium-containing compound semiconductors like CIS by
electrochemical means.
6.10
Conclusions
In this chapter we have summarized selected literature data on the electrodeposition
of semiconductors in ionic liquids. It has been demonstrated that elemental silicon,
germanium, and selenium can be elecrodeposited in ionic liquids. Furthermore,
it is shown that compound semiconductors like InSb, AlSb, CdTe and others can
be made, especially at elevated temperatures where kinetic barriers are easier to
overcome, even allowing the exclusive electrodeposition of grey selenium. In this
context ionic liquids are very promising for semiconductor electrodeposition. Both
wide electrochemical and thermal windows allow processes which are impossible
in aqueous or organic solvents.
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165
167
7
Conducting Polymers
Jennifer M. Pringle, Maria Forsyth, and Douglas R. MacFarlane
7.1
Introduction
The utilization of ionic liquids for the synthesis and use of conducting polymers
brings together two of the most exciting and promising areas of research from
recent years.
Conducting polymers are organic materials that can display electronic, magnetic
and optical properties similar to metals, but that also have the mechanical properties and low density of a polymer. They have the potential to allow the design and
fabrication of a vast number of electrochemical devices including photovoltaics,
batteries, chemical sensors, supercapacitors, conducting textiles, electrochromics
and electromechanical actuators [1–4]. In addition, these materials have the potential to impact in a major way on new biomedical processes such as controlled
neural growth, which has significant application in spinal regeneration [5]. The use
of electroactive polymers for the fabrication of electromechanical actuators, where
the polymer can be made to bend and straighten on application of a small potential,
has particular significance in the medical field, where they are being investigated
as artificial muscles for a range of prosthetic and therapeutic uses. The use of
conducting polymers in photovoltaic devices is another, extremely important area
of research and potential applications range from simple solar cells to sunlight
harvesting paints and fabrics, for the production of electricity from sunlight.
Research into conducting polymers has been increasingly intense for the last
25 years, since MacDiarmid, Heeger and Shirakawa published their seminal work
on polyacetylene, which demonstrated that the conductivity of these materials can
be increased by several orders of magnitude by doping with anions [6, 7]. The
importance of these materials and the progress made in this field is reflected in the
award of the Nobel Prize for Chemistry in 2000 to these founding researchers in
this area.
However, to allow the widespread use of conducting polymers, more research is
needed to improve their general performance, and one of their present limitations
is the rapid degradation of key properties such as conductivity and electrochemical cyclability. This limitation is primarily a result of the electrolyte used in the
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
168 7 Conducting Polymers
preparation and cycling of the electroactive polymer, which is required as a source of
dopants when the polymer is oxidized. These dopants have a significant influence
on all the properties of the polymer, including conductivity, mechanical properties, electrochemical efficiency and stability, most likely through the control of the
structure and morphology of the material. Ionic liquids offer a unique combination
of chemical and physical properties that make them interesting as electrolyte and
solvent in one. Interestingly, but not entirely of surprise at the fundamental level,
many of the anions that are effective in producing high conductivities in conducting polymers are also the anions that commonly occur in ionic liquid compounds.
Thus appropriate ionic liquids provide a superb source of the dopant anions.
Conducting polymers, such as poly(aniline), poly(pyrrole) or poly(thiophene)
(Figure 7.1) have a conjugated system of delocalized π-orbitals, which allows conduction to occur in the oxidized, or “doped” polymer.
The technological interest in these materials lies in their redox behavior. When
the films are oxidised in an appropriate electrolytic medium, positive charges are
generated along the backbone and solvated counterions enter the polymer from
the solution to effect charge balance. This results in an opening of the polymeric
structure and an increase in volume. The opposite process occurs on reduction,
when the incorporated anions are expelled back into solution and the film recovers
its original volume. The size and nature of the dopant counterion incorporated
during synthesis can have a dramatic effect on the ion movement occurring during
redox processes. There is a competitive reaction between anion expulsion and cation
incorporation (from the electrolyte) during the reduction cycle. These two reactions
compete to achieve charge neutrality caused by the loss of charge on the polymer
backbone (Figure 7.2) [4]. For example, polyelectrolyte dopants, being relatively
immobile, tend to remain in the polymer and with such films cation incorporation
processes dominate.
The potential benefits of using ionic liquids as electrolytes in conducting polymer
devices have been investigated by a number of authors in recent years, for applications such as actuators [8–17], supercapacitors [18–20], electrochromic devices
[12, 21] and solar cells [22], with significant improvements in lifetimes and device
performance reported.
For example, in 2002, Lu et al. [12] reported significant improvements in
device performance when the ionic liquids 1-ethyl-3-methylimidazolium hexafluorophosphate, [C2 mim][PF6 ], and 1-ethyl-3-methylimidazolium tetrafluoroborate, [C2 mim][BF4 ], were used as supporting electrolytes for poly(pyrrole) and
poly(aniline) actuators and for polyethylenedioxythiophene (PEDOT) in electrochromic devices, respectively. For the PEDOT study, the ionic liquid was
also used as the growth medium for the electropolymerization of PEDOT. Notably, in both the poly(aniline) in the [C2 mim][BF4 ] and the poly(pyrrole) in the
[C2 mim][PF6 ] actuator systems, cation incorporation and expulsion was the predominant strain-generating mechanism during actuation, whereas in the propylene carbonate/tetrabutylammonium hexafluorophosphate (PC/[Bu4 N]PF6 ) system
it was anion movement that was observed [12]. Thus, the linear displacement of this
actuator is in the opposite direction, illustrating the importance of the nature of ion
7.1 Introduction 169
Fig. 7.1 The chemical structure of common conducting polymers, in
their undoped, neutral form [4].
170 7 Conducting Polymers
Fig. 7.2 The redox cycling of poly(pyrrole) involving intercalation and
expulsion of (a) the anion or (b) the cation from the electrolyte to
effect charge balance.
incorporation in such ionic liquid systems. The poly(aniline) actuator maintained
electromechanical actuation and electroactivity for >10 000 cycles, without any
significant decrease in stress or strain (<1%), clearly demonstrating the potential
benefits of the use of ionic liquids in such systems.
Thus, when an ionic liquid is used as the solvent/electrolyte for electrochemical
cycling of conducting polymers, both the cation and the anion of the ionic liquid
may be intricately involved in the redox processes and, therefore, the nature of each
must be considered. The nature of the electrolyte is also critical in dictating the
stability of conducting polymers at extreme potentials, which is often limited by
degradation of the solvent or electrolyte and, therefore, the use of an ionic liquid
with a wide electrochemical window of stability, particularly in the anodic region,
will be of particular benefit in this respect.
In addition, the use of ionic liquids is often prompted by safety and environmental
considerations, where their negligible volatility and nonflammability makes them
ideal replacements for more toxic molecular solvents and, importantly, overcomes
the problem of solvent evaporation that exists with the long-term use of volatile
solvents in electrochemical applications. The wide liquid range and good thermal
stability are also extremely advantageous for device applications.
Further to their role as supporting electrolytes, the conductivity and electrochemical stability of ionic liquids clearly also allows them to be used as solvents for the
electrochemical synthesis of conducting polymers, thereby impacting on the properties and performance of the polymers from the outset. Parameters such as the
ionic liquid viscosity and conductivity, the high ionic concentration compared to
conventional solvent/electrolyte systems, as well as the nature of the cation and
7.2 Electropolymerization – General Experimental Techniques 171
anion themselves, may all influence integral polymer properties such as structure,
doping level, growth rate, growth mechanism, morphology, conductivity etc.
Exploration of these concepts began a number of years ago with the use of
chloroaluminate ionic liquids and has more recently focused on the use of air and
moisture stable species. These are discussed separately in Section 7.3. First we
discuss general experimental techniques in Section 7.2. Characterization methods
are then surveyed in Section 7.4.
7.2
Electropolymerization – General Experimental Techniques
The generally accepted polymerization mechanism for heterocyclic polymers such
as poly(pyrrole) is shown in Figure 7.3 [23].
Fig. 7.3 The polymerization mechanism for heterocyclic polymers [23].
172 7 Conducting Polymers
However, this polymerization process is far from simple and there are still aspects that are the subject of some debate, some of which may also be influenced
by the use of ionic liquids as the polymerization medium. For example, the point
at which the polymer begins to deposit on the electrode must be considered; it is
probable that there is a certain degree of polymerization that occurs in solution,
followed by the precipitation of oligomers onto the electrode surface, rather than
the adsorption of the monomer onto the electrode followed by its polymerization.
This oligomer deposition is influenced by the solubility of the intermediates and
the extent to which they diffuse away from the electrode rather than deposit onto
it; these factors may be influenced by the use of ionic liquids as a result of their
different viscosities, conductivities, solubilising properties and even their potential
for stabilizing radical or charged species. The color change that we have observed
during the use of ionic liquids for the polymerization of pyrrole suggests the presence of significant quantities of oligomers dissolved in the ionic liquid. They way in
which the polymer subsequently grows is also important; it may be via nucleation
and growth processes, where the electropolymerization and precipitation continues, or by chain growth of the polymer already on the electrode [23]. These different
mechanisms will influence the morphology of the final polymer film, as discussed
in Section 7.4.2, and again may be influenced by the use of an ionic liquid.
There are also a number of other variables to consider when planning the electrochemical synthesis of conducting polymers in ionic liquids. While most of these
variables also exist for the synthesis of the polymers in molecular/solvent systems
and have been investigated in detail, it is worth considering that the influence of
any of these factors may be different when utilizing ionic liquids as the growth
medium because of their distinctly different properties. These are discussed in
more detail below.
7.2.1
Temperature
When ionic liquids are used, this will have a significant effect on the viscosity and
hence the conductivity and rate of ion diffusion within the ionic liquids. Growth of
conducting polymers at reduced temperatures (as low as −28 ◦ C) [4, 24] in molecular
solvent systems is generally accepted to result in smoother, more conductive films,
but we have found that in ionic liquids the significant increase in the viscosity
can be problematic. In addition, the temperature used for the conducting polymer
synthesis may be limited by the melting point of the ionic liquid [25].
7.2.2
Electrochemical Technique
Conducting polymer films are generally grown galvanostatically, potentiostatically
or potentiodynamically, although pulsed current and pulsed potentials can also be
used. All three of these techniques have been reported for the growth of various
conducting polymers in ionic liquids, but there has been no comprehensive comparison of the relative benefits of each. In molecular solvents, constant potential
7.2 Electropolymerization – General Experimental Techniques 173
growth may result in less homogeneous films than those from constant current
growth due to inhomogeneities on the electrode surface, and this technique may
be more affected by iR drop effects [4], which may be considerably larger in the
more viscous ionic liquid medium. During galvanostatic growth, the potential must
be monitored closely to ensure that overoxidation of the polymer does not occur.
Growth of the polymer film potentiodynamically (by cyclic voltammetry) is the
most time-consuming technique and is not generally employed for the synthesis
of significant quantities of film, but it is an advantageous technique for detailed
electrochemical analysis as the redox characteristics of the polymer can be monitored during film growth. It should be noted, however, that during potentiodynamic
growth the auxiliary electrode can be exposed to extreme positive potentials. In a
molecular solvent system this can result in oxidation of the solvent or any water
present, but this problem may be reduced in a more electrochemically stable ionic
liquid medium.
7.2.3
Growth Potential
The oxidation potential of various monomers may be different in ionic liquids compared to molecular solvent/electrolyte systems. An increase in monomer oxidation
potential in ionic liquids has been attributed to a decrease in the stability of the
monomer cation (3-(4-fluorophenyl)thiophene) [26], and also to the viscosity of the
ionic liquid, which causes significant iR drop within the electrochemical cell [27].
However, it should also be noted that determination of this potential is somewhat
dependent on the nature of the reference electrode used and also any variation that
this electrode might display on moving from molecular solvent systems to ionic
liquids [28]. Thus, direct comparison of monomer electrode potentials, as reported
by different researchers, should be undertaken with considerable caution. Upon
the initial investigation of any new ionic liquid or monomer, or even when using
new experimental conditions, it is prudent to record the cyclic voltammograms of
the monomer in the ionic liquid, starting with a modest potential range and slowly
increasing until monomer oxidation is observed, thereby determining the optimal
conditions for polymer growth.
It is interesting to note at this point that, where previously the potential used to
grow the electroactive polymers has been limited by the solvent window, one of
the many benefits of using ionic liquids is their greater electrochemical stability,
which allows access to much higher potentials for film growth. This should allow
access to the oxidation potentials of monomers previously unreachable and thus increase the range of conducting polymers that can be synthesized by electrochemical
methods.
7.2.4
Electrodes
The choice of a suitable reference electrode for use in ionic liquids is a complex
issue [28], which will be discussed in more detail elsewhere. There is also a range
174 7 Conducting Polymers
of different working electrodes available; existing studies of conducting polymer
synthesis in ionic liquids almost entirely use platinum working electrodes, or indium titanium oxide (ITO) glass to allow spectroelectrochemical analysis of the
polymer. The synthesis of poly(pyrrole) onto iron has also been studied [29], and
a glassy carbon electrode has been used for the deposition of PEDOT [25]. As for
all electrochemical techniques, the condition and cleanliness of the electrodes is
of great importance; it has been reported that the morphology of poly(pyrrole)
films can be greatly altered by inadequate polishing of the working electrode [30].
The working electrode can also be used as a template to allow the synthesis of
nanostructured polymers, although this technique has not been extensively utilized in ionic liquids [31]. The cell design and the arrangement and size of the
electrodes are also very important, influencing the hydrodynamics and potentials
within the cell and thus the rate of ion transport and the quantity of unwanted
side-reactions.
7.2.5
Atmosphere and Water Content
The presence of oxygen in a molecular solvent/electrolyte system during the electrochemical synthesis of a conducting polymer can be problematic as it can react
with radical intermediates and be reduced at the auxiliary electrode to form hydroxide [4]. The vast majority of reports detailing the synthesis of conducting polymers
in ionic liquids perform this procedure under anaerobic conditions, using either a
nitrogen or an argon atmosphere. In our laboratory we have found that even a basic
dry-nitrogen blanket can yield significant improvements in the conductivities of
synthesized poly(pyrrole) films, probably by decreasing over-oxidation of the polymer. Similarly, the redox responses of poly(pyrrole) in ionic liquids can be much
more defined after nitrogen purging [32].
The water content of the ionic liquid is another, perhaps more complex, factor for
consideration. The presence of small amounts (1–4%) of water may be beneficial
to the properties of some polymers [33], but it should be noted that some ionic
liquids utilizing fluorinated anions, particularly hexafluorophosphate, can slowly
hydrolyse in the presence of water to form HF and other species [34]. It is important
to note that even if an ionic liquid is immiscible with water, it may still contain a
significant quantity of water before drying, which can alter its physical properties,
such as viscosity [34–36].
7.2.6
Choice of Ionic Liquid
There is a plethora of different ionic liquids readily available, either commercially
or through straightforward laboratory synthesis; investigations into their use for
the synthesis of conducting polymers has, so far, focused on a relatively small
number (Figure 7.4). Thus, the avenues for future investigation in this area are vast.
7.2 Electropolymerization – General Experimental Techniques 175
Fig. 7.4 The cations and anions utilized to date for the electrochemical synthesis of conducting polymers in ionic liquids, and their abbreviations.
176 7 Conducting Polymers
However, when considering the choice of ionic liquid it is worth noting that a
number of the anions that are utilized in ionic liquids have already been investigated as dopants for conducting polymers using a conventional molecular solvent/electrolyte system. For example, the relative merits of the trifluoromethanesulfonate [OTf]− [37], hexafluorophosphate [PF6 ]− [38, 39], sulfonated aromatics
[40, 41] and, particularly, bis(trifluoromethanesulfonyl)amide [NTf2 ]− [37, 42–45]
anions have been well studied and it may be pertinent to consider this research
when selecting an ionic liquid for investigation.
There are also a number of other considerations when selecting an ionic liquid.
Of course, both the cation and anion must be chemically and electrochemically
stable. High viscosity and low conductivity ionic liquids may be problematic, particularly for polymer deposition onto large working electrodes. The viscosity of the
ionic liquid affects the conductivity and rate of ion diffusion within the ionic liquid;
hence it is easy to envision that it will also have a significant effect on the electrochemical polymerization process. The viscosity is easily tailored by changing the
cation and anion; the dicyanamide anion imparts very low viscosities, but we have
found synthesis of conducting polymers in ionic liquids utilizing this anion very
problematic. It is also primarily the anion that dictates the hydrophobicity of the
ionic liquid, but there has been no investigation into how this parameter affects the
synthesis of conducting polymers in ionic liquids.
The monomer must be soluble in the ionic liquid at adequate concentrations.
Indeed, the solubility of some monomers may be improved using an ionic liquid; synthesis of poly(terthiophene) is often hampered by the poor solubility of
the monomer, but terthiophene can be dissolved in 1-ethyl-3-methylimidazolium
bis(trifluoromethanesulfonyl)amide, [C2 mim][NTf2 ] or N,N-butylmethylpyrrolidinium bis(trifluoromethanesulfonyl)amide, [C4 mpyr] [NTf2 ], at concentrations up
to 0.05 M [27].
The size and nature of the ionic liquid ions are expected to influence the extent to
which they are incorporated into the polymer during growth or electrochemical cycling. The size of the cation can be easily tailored by modifying the length of the alkyl
substituent, and choice of a planar aromatic cation, such as imidazolium, rather
than non-planar aliphatic species such as pyrrolidinium, may also enhance cation
intercalation. These are also particularly important considerations with respect to
their use in electromechanical actuators, where the magnitude of displacement is
a direct result of ion movement into and out of the film.
Considerations of ionic liquid recyclability, cost, toxicity etc. are also important.
These are less pertinent for small-scale laboratory investigations but will become
increasingly important as the field progresses towards larger scale synthesis for
device applications.
Sekiguchi et al. [46] have reported the recycling and re-use of 1-ethyl3-methylimidazolium trifluoromethanesulfonate, [C2 mim][OTf], after poly
(pyrrole) synthesis by extraction of the unreacted monomer with chloroform. The
ionic liquid was reused five times with little change in the growth CVs of the
polymer.
7.3 Synthesis of Conducting Polymers 177
7.3
Synthesis of Conducting Polymers
7.3.1
Synthesis in Chloroaluminate Ionic Liquids
The first investigations into the use of ionic liquids, or molten salts as they were
formerly known, for the synthesis of conducting polymers utilized those composed
of a mixture of AlCl3 and organic chlorides such as N-butylpyridinium chloride,
[C4 py][Cl] cetylpyridinium chloride, [C16 py][Cl] and 1-ethyl-3-methylimidazolium
chloride, [C2 mim][Cl]. In these chloroaluminate systems, if the organic chloride is
present in excess then the melt contains Cl− and AlCl4 − anions and is considered
to be basic. If the AlCl3 is in excess then the melt is acidic and only AlCl4 − and
Al2 Cl7 − anions are present. These can also be made superacidic by the addition of
protons from, for example, 1-ethyl-3-methylimidazolium hydrogen dichloride [47].
A neutral chloroaluminate ionic liquid, which contains Al2 Cl7 − and Cl− anions,
is obtained by using an exactly equimolar amount of the organic chloride and the
AlCl3 . The melts can also be buffered to neutrality using alkali halides.
Although the use of these molten salts is hampered by their instability in air and
water, and this instability may also be reflected in the resultant polymer films, it is
important to note that much of this earlier work clearly identifies a number of the
potential benefits of using ionic liquids for the synthesis of conducting polymers.
7.3.1.1 Poly(pyrrole)
Pickup and Osteryoung investigated the polymerization of pyrrole in both the
AlCl3 /[C4 py][Cl] molten salt [48] and the more conductive AlCl3 /[C2 mim][Cl] [49].
Synthesis of poly(pyrrole) is only possible in neutral melts. In basic melts oxidation
of Cl− to Cl2 occurs before the monomer oxidation and in acidic melts no films
are produced due to the formation of a 1:1 AlCl3 –pyrrole adduct, as determined by
NMR spectroscopy [50].
The electrochemical behavior of poly(pyrrole) films prepared and cycled in an
AlCl3 : [C2 mim][Cl] melt was investigated in detail and improvements in reproducibility and the rate of oxidation and reduction of these films were observed
compared to films prepared under similar conditions in acetonitrile [49]. This was
postulated to be a result of an increase in the porosity of poly(pyrrole) films deposited from the melt compared to those from acetonitrile, although attempts to
describe this porosity using porous electrode models were not totally conclusive.
This appears to be contradictory to the smoother poly(pyrrole) films that are formed
in air- and water-stable ionic liquids [46, 51].
Most recently, Geetha and Trivedi have revisited the synthesis of poly(pyrrole)
in such media using a [C16 py][Cl]/AlCl3 melt [52]. They report that use of this
melt avoids the unwanted cathodic deposition of aluminum that prevents the
AlCl4 − /Al2 Cl7 − from being available as the dopant or the polymerization initiator
in electrochemical Friedel Crafts reactions. Using [C16 py][Cl] prevents aluminum
178 7 Conducting Polymers
deposition onto the electrode as the cetyl groups are preferentially deposited. For
similar reasons, the authors have also recently investigated this media for the electropolymerization of benzene [53].
7.3.1.2 Poly(p-phenylene)
The polymerization of benzene in chloroaluminate salts has attracted attention
from a number of authors. The electrosynthesis of this polymer is highly desirable
but challenging, with one of the primary considerations being the strict elimination of water from the reaction medium, which is most commonly effected using
difficult systems such as liquid sulfur dioxide, concentrated sulfuric acid or HF.
The polymerization of benzene and biphenyl in organic solvents yields only low
conductivity films with low degrees of polymerization.
The earliest investigations into the use of chloroaluminate molten salts overcame
problems associated with the relatively high melting points of the then-known
chloroaluminates (containing [C2 py][Br] or [C16 py][Cl]) by mixing with equivalent
volumes of benzene. This also results in a decrease in viscosity and marked increase
in conductivity [54, 55]. Later, Goldenberg and Osteryoung used the lower-melting
system of [C2 mim][Cl]/AlCl3 and reported that the polymerization of benzene was
facilitated in this melt, compared to solvents like acetonitrile or nitromethane,
because it is drier and the nucleophilicity is low [56]. Unlike pyrrole, aniline or
thiophene, benzene does not form adducts with AlCl3 and is polymerized using an
acidic chloroaluminate melt. Further, the oxidation potential of the monomer does
not appear to vary with melt composition. The resultant poly(p-phenylene) films
demonstrated good electrochemical stability when cycled in the melts; films were
cycled in a 1.5:1 acidic melt in excess of 100 times with no significant change in
activity [56].
Kobryanskii and Arnautov reported a significant increase in the relative molecular mass of poly(p-phenylene) synthesised in [C4 py][Cl]/AlCl3 compared to the
highest values previously reported using alternative media [57], and Goldenber et
al. [58] reported the formation of highly conductive films from this melt, although
this conductivity decreased very rapidly with time, even in an inert atmosphere.
Substitution of the chloroaluminate anions in the film with BF4 − anions reportedly
increased the stability, but at the expense of the conductivity. The film instability
was attributed to the presence of the highly moisture sensitive chloroaluminate
anions and possible destruction of the polymer by anion oxidation and polymer
chlorination. There is also some question as to the quality (degree of polymerization and crosslinking, conjugation length and so on) of the poly(p-phenylene)
films made from such melts [59]. Most recently, Geetha and Trivedi have used a
[C16 py][Cl]/AlCl3 system dissolved in benzene for the synthesis of poly(p-phenylene)
[53], producing films with good conductivity, but again this decreased markedly on
exposure to air.
7.3.1.3 Poly(thiophene)s and Poly(fluorene)
Osteryoung and coworkers have also used chloroaluminate molten salts, utilizing the [C2 mim] cation, for the electrosynthesis of poly(thiophene) and poly
7.3 Synthesis of Conducting Polymers 179
(bithiophene) [60] and poly(fluorine) [61]. The electrosynthesis of poly(fluorine)
can be achieved in both neutral and acidic melts but in the basic composition
the chloride ions are more easily oxidized. The resultant poly(fluorene) films are
reportedly more stable than those synthesized in acetonitrile and the ability to observe a ferrocene/ferrocenium couple through a thin deposit of poly(fluorine) on
the working electrode was concluded to indicate a porous film morphology [61].
Similar observations of stability and porosity were made for poly(thiophene) and
poly(bithiophene) films prepared in this ionic liquid [60]. Poly(bithiophene) films
prepared in the neutral melt were found to be unstable in acetonitrile, although
films grown in acetonitrile and then electrochemically cycled in the melt exhibited
excellent stability (>2000 cycles).
7.3.1.4 Poly(aniline)
Osteryoung and coworkers have also investigated the use of chloroaluminate ionic
liquids for the synthesis of polyaniline [62, 63]. Unlike pyrrole and thiophene,
aniline was successfully polymerized in acidic, neutral and basic chloroaluminate
melts, although the best results were obtained using the neutral composition.
The oxidation potential of aniline is significantly affected by the composition of
the melt, probably as a result of the formation of an adduct between aniline and
AlCl3 , similar to that observed for pyrrole. The oxidation potential of the aniline
was also found to be influenced by the nature of the electrode used (platinum or
glassy carbon), and further shifted upon deposition of polymer onto the working
electrode, in a direction that again depended on the composition of the molten
salt used [62]. Further, using molten salts of different acidity is reported to result
in changes to the backbone structure of the resultant poly(aniline). In comparison
to poly(aniline) films formed from aqueous and organic solvents, those prepared
and cycled in basic chloroaluminate melts were reportedly very stable, retaining
more than 90% of their electrochemical activity after 30 000 cycles at 100 mV s−1
[63]. They also observed what was believed to be the influence of the viscosity
of the molten salt medium on the electrochemical behavior of the conducting
polymer; when the poly(aniline) electrode, in its insulating state, was placed in
a basic or neutral melt, it required around 30 potentiodynamic cycles before the
maximum electroactivity was obtained. This kinetic limitation is most likely related
to the ability of the anions to permeate the film (solvent swelling), which can be
expected to be significantly different in this, relatively viscous, molten salt media
compared to molecular solvent systems. It is also dependent on the size of the anion
present in the molten salt. This observation has subsequently been corroborated
by researchers using air and water stable ionic liquids [64].
7.3.2
Synthesis in Air- and Water-stable Ionic Liquids
7.3.2.1 Poly(pyrrole)
Poly(pyrrole) is one of the most popular conducting polymers as it can be highly conducting, quite environmentally stable and relatively easy to synthesize. Sekiguchi
180 7 Conducting Polymers
et al. [46] have studied synthesis of this polymer in ionic liquids utilizing the [C2 mim]
cation and the [BF4 ]− , [PF6 ]− and [OTf]− anions. Comparison of the growth CVs of
poly(pyrrole) in these ionic liquids showed the highest polymer oxidation and reduction currents were obtained during growth in [C2 mim][OTf], and the lowest from
using the [C2 mim][PF6 ]. The authors thus suggest that the higher viscosity of the
former species results in a faster polymerization rate, because the polymerization
reaction products are accumulated near the electrode surface and thus can undergo
further radical coupling, oxidation and deposition onto the electrode rather than
diffusing away into the bulk. However, this is not always seen and therefore other
factors also influence these processes. The authors concluded that the [OTf]− anion
was the superior choice for the electropolymerization of pyrrole. When comparing
the performance of this neat ionic liquid to that of solutions of the ionic liquid
in water or acetonitrile it was noted that the undiluted ionic liquid gave greater
polymer redox peak currents and smaller redox peak separation. The film grown
in the neat ionic liquid was reportedly thinner than that from the dilute solutions
but more electrochemically active and more highly doped. The films grown from
the aqueous and acetonitrile solutions displayed a granular morphology (larger
from the water solution) whereas the film from the ionic liquid appeared to be
very smooth. The origin of the cauliflower-like features observed in poly(pyrrole)
from molecular solvent systems has been investigated by a number of researchers,
who suggest that their appearance can be significantly altered depending on the
nature of the anion used, the electrode material and polishing techniques used,
the synthesis temperature and so on [4]. Sekiguchi et al. [65] also report the use of
poly(pyrrole) films as a matrix for hosting palladium particles and show that these
are considerably more dispersed when deposited from the ionic liquid than from
aqueous solutions of the salt.
Significantly smoother film morphologies have also been observed for
poly(pyrrole) grown from the ionic liquids [C2 mim][NTf2 ] and [C4 mpyr] [NTf2 ] compared to those grown under the same experimental conditions from PC/Bu4 NPF6
(Figure 7.5) [51]. However, lower polymer redox currents were observed in the
more viscous, less conductive pyrrolidinium ionic liquid compared to the imidazolium species (85 cP and 2.2 mS cm−1 , respectively, at 25 ◦ C for [C4 mpyr] [NTf2 ]
[66] compared to 34 cP and 8.8 mS cm−1 at 20 ◦ C for [C2 mim][NTf2 ] [67]).
Fig. 7.5 Poly(pyrrole) films grown in [C4 mpyr][NTf2 ] (a),
[C2 mim][NTf2 ] (b) and PC/Bu4 NPF6 (c), constant potential onto Pt.
7.3 Synthesis of Conducting Polymers 181
The electrochemical synthesis of poly(pyrrole) from [C4 mim][PF6 ] has also been
studied in some detail [29, 32, 68]. Fenelon and Breslin used this ionic liquid to
deposit poly(pyrrole) onto an iron electrode [29], an example of the electrochemical deposition of a conducting polymer from an ionic liquid onto a corrosionsusceptible electrode rather than the inert species used in the other studies. The
polymer was deposited onto iron using a relatively high potential (1.3 V vs. a Ag
wire quasi-reference electrode) but was electroactive and conducting, indicating
negligible polymer overoxidation problems compared to those associated with using aqueous systems or even those observed using a platinum working electrode
at these potentials in this ionic liquid. The authors determined a critical growth
concentration of 0.2 mol dm−3 , above which the rate of electropolymerization does
not significantly increase, and reported that the poly(pyrrole) layers grown onto
this substrate were also smoother than those grown from aqueous systems. The
poly(pyrrole)/iron electrode was stable for periods greater that 16 h in this ionic
liquid, maintaining its low charge-transfer resistance and with no dissolution of
iron through the pores on the polymer coating detected, as has been observed using
aqueous systems. In this investigation the authors performed the polymerization in
a dry nitrogen atmosphere, with an ionic liquid water content of ca. 10 ppm, below
which it is reported that the polymerization rate decreases. Mazurkiewicz et al. [32]
have also reported an influence of water content on the polymerization of pyrrole
in this ionic liquid. They demonstrated that the redox response of the polymer
growth CV was much more defined when the polymer was grown in [C4 mim][PF6 ]
that had been purged with dry nitrogen (Figure 7.6(a)) compared to the response
using ionic liquid that had been equilibrated in air (Figure 7.6(b)), and also showed
the difference in growth CVs in [C4 mim][PF6 ] compared to that in deoxygenated
0.25 M Bu4 NPF6 in PC (Figure 7.6(c)) [32].
Figure 7.6 demonstrates the typical appearance of growth CVs of conducting
polymers; the potential is repeatedly cycled in a positive and negative direction and
at potentials above the oxidation potential of the monomer, polymer deposition
onto the electrode occurs. The polymer oxidation and reduction peaks show an
increase in current with successive cycles (the arrows show the direction of peak
progression) indicating the deposition of increasing amounts of electroactive polymer. The peak position may also shift as the film becomes thicker, which may be
attributed to factors such as heterogeneous electron-transfer kinetics or a decrease
in conductivity, counter-ion mobility or conjugation length.
A significant improvement in cycle life was also demonstrated for the
poly(pyrrole) in the [C4 mim][PF6 ] (>900 cycles) compared to cycling in 0.25 M
PC/Bu4 NPF6 (300 cycles) [32].
Finally, during the synthesis of poly(pyrrole) in [C2 mim][NTf2 ] an unusual mechanism of growth was observed, with the polymer growing along the surface of the
ionic liquid [69]. When the working electrode is a thin platinum wire, and the
reaction is performed in air but using nitrogen-purged ionic liquid, the polymer
grows along the surface of the ionic liquid after forming an initial thin layer on the
submerged body of the electrode (Figure 7.7(a)). We believe that the presence of
some water is necessary for this “solution-surface electropolymerization”, to react
182 7 Conducting Polymers
Fig. 7.6 Growth of poly(pyrrole) in (a) air-equilibrated [C4 mim][PF6 ], (b) [C4 mim][PF6 ] after N2(g) purging and (c) deoxygenated 0.25 M Bu4 NPF6 in
PC, 100 mV s−1 , 30 cycles [32].
7.3 Synthesis of Conducting Polymers 183
Fig. 7.7 Poly(pyrrole) grown along the surface of [C2 mim][NTf2 ]: (a)
constant potential, (b) with a circular auxiliary electrode, (c) using
voltage pulses [69].
with the H+ produced in the reaction (water being a stronger base than the [NTf2 ]−
anion) and when a dry ionic liquid is used this is provided by absorption from the
atmosphere. This phenomenon can be encouraged using an auxiliary electrode that
circles the working electrode, to give directional growth (Figure 7.7(b)). Further,
using a pulsed voltage the polymer forms first as a series of fibrils that can extend
over a significant portion of the film. (Figure 7.7(c)). This fine structure imparts a
larger surface area to the polymer than would be present in a solid, homogeneous
film.
7.3.2.2 Thiophenes
Shi et al. [70] were the first to demonstrate the use of an air and moisture stable
ionic liquid, [C4 mim][PF6 ], for the electrochemical synthesis of poly(thiophene),
grown onto a platinum working electrode by potentiodynamic, constant potential
or constant current techniques. The use of growth potentials between 1.7 and
1.9 V (vs. Ag/AgCl) reportedly gave smooth, blue–green electroactive films, whereas
potentials above 2 V resulted in film destruction by overoxidation.
Sekiguchi et al. [65] used [C2 mim][OTf] for the polymerization of thiophene and
demonstrated larger polymer redox currents during potentiodynamic growth in this
ionic liquid than were observed in a 0.1 M solution of the ionic liquid in acetonitrile.
Thus, as for poly(pyrrole), the authors concluded that this ionic liquid gave a higher
growth rate, as well as smoother films and improved electrochemical capacity.
The growth of poly(thiophene), poly(bithiophene) and poly(terthiophene) in
[C2 mim][NTf2 ] and [C4 mpyr][NTf2 ] has also been studied [27]. The oxidation potential of these monomers decreases with increasing chain length, consistent with
their behavior in conventional electrolyte/solvent systems. This is the primary advantage of using such materials – the high potential required to oxidise thiophene
can result in side-reactions and overoxidation of the poly(thiophene) polymer film;
in other words poly(thiophene) is not stable at the potentials required for its synthesis [71]. The use of dimers or oligomers of thiophene is one way of overcoming this
“polythiophene paradox” [72], and increasing the stereoregularity of the polymer
by reducing the number of β,β or α,β mis-linkages. Ideally the conjugation length
184 7 Conducting Polymers
Fig. 7.8 Cyclic voltammograms of thiophene polymerization
(0.2 M, 50 mV s−1 ) onto a Pt working electrode: (a) growth and (b)
post-growth in [C2 mim][NTf2 ], (c) growth and (d) post-growth in
[C4 mpyr][NTf2 ], vs. a Ag pseudo-reference electrode. Arrows indicate
the peak development with successive scans [27].
of the polymer would also be increased, although in reality the opposite may occur
[73, 74].
The oxidation potential for all of the monomers appeared to be higher in the
pyrrolidinium ionic liquid by approximately 0.1 V, which may be due to the higher
viscosity and lower conductivity of this ionic liquid compared to the imidazolium
species. There are also significant differences in the growth CVs of the polymers
in the two different ionic liquids (Figure 7.8).
For each monomer and ionic liquid, measurement of the total cathodic charge
passed during reduction of the polymers in the final post-polymerization CVs,
compared to the peak polymer oxidation currents from the final growth cycles,
allows comparison of the film electrochemical activities while taking into account
the relative amounts of the polymer. The former value is often used as an indication
of the amount of polymer grown, but this assumes that the electrochemical activities
of the films are identical.
7.3 Synthesis of Conducting Polymers 185
The peak oxidation current during the final growth cycle of poly(thiophene) is
slightly higher in the imidazolium ionic liquid (Figure 7.8(a)), whereas the film
from the pyrrolidinium species exhibits a larger total reduction charge in the postpolymerization CV, suggesting better electrochemical activity, possibly as a result
of slower, more ordered film growth. Alternatively, this may indicate the superiority
of the pyrrolidinium ionic liquid as a cycling solvent, but this is probably less likely
given its high viscosity; differences in the post-polymerization CVs of such polymer
films in the ionic liquids probably reflect both an influence of the nature of the ionic
liquids on the polymer growth and the effect of using the different ionic liquids
as the solvent for the post-polymerization cycling. Thus, for a direct assessment of
the influence of the ionic liquid on polymer growth, post-polymerization cycling
may be best performed in the same solvent, possibly a molecular solvent/electrolyte
system. However, the proven benefits of using ionic liquids as the supporting media
for the electrochemical cycling of conducting polymers, such as improved stability,
suggest that assessment of post-polymerization cycling of the polymers in the ionic
liquids is of more interest to researchers considering utilization of these materials
in electrochemical devices.
For the electropolymerization of bithiophene [27], which is adequately soluble
in both ionic liquids, under the same conditions the growth CVs (Figure 7.9)
suggest a stronger influence of the nature of the ionic liquid than was observed for
the thiophene monomer, with the polymerization of bithiophene appearing to be
four times faster in the [C2 mim][NTf2 ] than in the [C4 mpyr][NTf2 ], and the redox
currents in the post-polymerization CVs also proportionally larger. Further, there
is only one distinct reduction peak evident during growth and cycling of the film in
the imidazolium ionic liquid but there are two distinct peaks evident during growth
in the pyrrolidinium species, although in the post-polymerization cycles these are
considerably broadened.
Terthiophene is less soluble than thiophene or bithiophene, and is insoluble
in most organic solvents, but concentrations of 0.05 M are attainable in both
[C2 mim][NTf2 ] and [C4 mpyr][NTf2 ] with gentle (50 ◦ C) heating. This is an adequate concentration for the electrosynthesis of coherent poly(terthiophene) films
onto either small Pt working electrodes, for electrochemical analysis, or onto ITO
glass of a few square centimeters for analysis by UV–Vis and any photovoltaic testing required [75]. In the growth CVs of poly(terthiophene) in these ionic liquids
(Figure 7.10) the influence of the nature of the ionic liquid on the growth rate of
the polymer films is again evident, with significantly faster growth in the imidazolium species, and proportionally larger redox currents in the post-polymerization
CVs (Figure 7.10 (b,d)). The films display multiple redox peaks during postpolymerization cycling in the ionic liquids, more clearly defined in the imidazolium
ionic liquid. The poly(terthiophene) films exhibit good reversibility and the redox
currents appear relatively stable over the 15 cycles recorded.
Murray et al. [76] have demonstrated that an ionic liquid can be used as both
the growth medium for poly(terthiophene) and also as a route to incorporation of
anionic dyes into the polymer, for use in photovoltaic devices. Again, the solubility
of terthiophene in [NTf2 ]-based ionic liquids is demonstrated; 1.6 × 10−2 M in
186 7 Conducting Polymers
Fig. 7.9 Cyclic voltammograms of bithiophene polymerization (0.1 M,
50 mV s−1 ): (a) growth and (b) post-growth in [C2 mim][NTf2 ], (c)
growth and (d) post-growth in [C4 mpyr][NTf2 ], vs. a Ag pseudoreference electrode. Arrows indicate the peak development with successive scans [27].
[C4 mim][NTf2 ]. Electropolymerization of terthiophene, by potentiodynamic cycling
with an ITO glass working electrode, from a solution of the ionic liquid containing
terthiophene and the anionic dye Erioglaucine, resulted in the formation of a thick,
mechanically strong polymer film with the dye incorporated as a dopant. The films
produced were more robust than those obtained using dimethylformamide as a
solvent, and these could then be reduced in an acetonitrile solution of a cationic
dye (brilliant green) to yield polymer films containing both anionic and cationic
dopants, resulting in a significantly improved photovoltaic performance.
Naudin et al. [26] have studied the electrochemical synthesis of poly(3(4-fluorophenyl)thiophene) in two alternative [NTf2 ]− ionic liquids, utilizing
the 1-ethyl-2,3-dimethylimidazolium and 1,3-diethyl-5-methylimidazolium cations.
These ionic liquids have melting points of 20 and −22 ◦ C, respectively, and viscosities of 88 and 36 cP, respectively, at 20 ◦ C [67]. The authors report that the oxidation
potential of the monomer is higher in these ionic liquids (1.16 and 1.22 V, respectively, for galvanostatic growth at 12.7 mA cm−2 ) compared to growth of this
polymer in propylene carbonate or acetonitrile (0.98 and 1.1 V, respectively), which
7.3 Synthesis of Conducting Polymers 187
Fig. 7.10 Cyclic voltammograms of terthiophene polymerization
(0.01 M, 50 mV s−1 ): (a) growth and (b) post-growth in [C2 mim][NTf2 ],
(c) growth and (d) post-growth in [C4 mpyr][NTf2 ], vs. a Ag pseudoreference electrode [27].
they attribute to the lower stability of the monomer cation radical in the ionic liquids
(a lower stability of the cation that is formed during the electrochemical polymerization would make its formation less energetically favorable and therefore require
higher potentials). Electrochemical analysis of the films by cycling in the ionic liquids indicated slower redox processes than those observed for the films grown and
cycled in 1 M acetonitrile/[Et4 N]BF4 , as evidenced by a larger separation between the
anodic and cathodic peaks. This peak separation is only partly attributed to the lower
ionic conductivity of the ionic liquids (3.2 and 6.6 mS cm−1 , respectively) compared
to the acetonitrile solution (43 mS cm−1 ); it also reflects differences in the polymer
films. The slower redox processes are also consistent with the smooth morphology
of the films. The doping levels, determined electrochemically, appeared to be little
influenced by the growth media and unfortunately the n-dedoping wave of this
188 7 Conducting Polymers
polymer overlaps with the negative limit of electrochemical stability of these ionic
liquids (−2.1 V), which somewhat limits the redox switching. In general, the films
exhibited poorer electrochemical activity in the ionic liquids, which was attributed
to poorer swelling and slower ion transport kinetics, although the kinetics could be
improved by growth of a thinner film. The doping of the polymers was also studied
using X-ray photoelectron spectroscopy (see Section 7.4.3), which also indicated the
presence of residual ionic liquid in the films that was hard to remove by washing.
The electrochemical activity of the films was assessed by scanning at 100 mV s−1
over the complete p-doping and n-doping range for the polymer and a rapid decrease
in activity was observed (75% after 50 cycles). The authors suggest that this may be
a result of gradual loss of ionic liquid from the polymer (deswelling) during cycling.
An alternative explanation is that ions are trapped in the polymer, as evidenced by
a significant charge imbalance between the n-doping and n-dedoping charges in
the CV. Interestingly, the authors note that a polymer film cycled in the n-doping
region in the ionic liquid could be reactivated by cycling in the p-doping region in
the same ionic liquid, or by transferal to the acetonitrile solution to re-swell the
polymer.
7.3.2.3 Poly(3,4-ethylenedioxythiophene)
Poly(3,4-ethylenedioxythiophene) (PEDOT) is a particularly popular conducting
polymer as it can have good conductivity and stability and has a low band gap,
which is pertinent to its use in photovoltaic devices. A number of authors have now
studied the electrochemical synthesis of this polymer in different ionic liquids. Lu
et al. [77] first demonstrated the use of [C4 mim][BF4 ] to electrodeposit PEDOT onto
ITO, and its application in an electrochromic numeric display.
Randriamahazaka et al. [64, 78, 79] have studied the synthesis and behavior
of PEDOT in [C2 mim][NTf2 ] in detail. In their primary report, the authors electrodeposited a PEDOT film from an acetonitrile/LiClO4 solution and studied its
electrochemical behavior when cycled in the ionic liquid [79]. In their subsequent
paper [64], they reported the electrochemical response of PEDOT that was grown
in the ionic liquid, and cycled in the ionic liquid that also contained lithium
bis(trifluoromethanesulfonyl)amide (LiNTf2 ), and contrasted this with the behavior
of a PEDOT film prepared in acetonitrile. PEDOT grown from acetonitrile and cycled in the ionic liquid displayed two oxidation and reduction peaks, the less anodic
of which decreased in peak potential but increased in current upon increasing the
concentration of LiNTf2 in the ionic liquid. In contrast, PEDOT that was prepared
in the ionic liquid itself displayed only one anodic and one cathodic peak (scan rate
100–500 mV s−1 ), at the same position as the second oxidation and reduction peak
that was observed in the CV of PEDOT grown from acetonitrile, and the presence
of LiNTf2 in the ionic liquid had little effect on the electrochemical behavior of
the film. In both cases it was concluded that the imidazolium cation of the ionic
liquid was the primary intercalating/de-intercalating species. It is also interesting
to note that when the PEDOT film from acetonitrile was cycled in the ionic liquid,
the authors observed a continuous change in the shape of the CV and increases
in the redox current (up to 20 cycles), which was attributed to the uptake of the
7.3 Synthesis of Conducting Polymers 189
ionic liquid into the film. This effect has also been observed during the cycling of
PEDOT in the pyrrolidinium analogue [80].
Damlin et al. [81] have reported the synthesis and p-doping and n-doping of
PEDOT in [C4 mim][BF4 ] and [C4 mim][PF6 ] and characterized the resultant films
by CV, in situ UV–Vis spectroelectrochemistry and ATR-FTIR (see also Section
7.4). Here, two oxidation peaks were observed in the first few growth cycles in the
[C4 mim][BF4 ] (at 50 mV s−1 ), which merged into one as the film became thicker, and
two reduction peaks were also seen. In the [C4 mim][PF6 ] three oxidation peaks were
observed at first, again merging into one with successive cycles, thus indicating an
influence of the anion. In this ionic liquid, two reduction peaks are again evident.
The authors report that the shapes of the CVs, and the oxidation potential of the
monomer, are similar in the ionic liquids to those in organic solvents using Bu4 N
PF6 or Et4 N BF4 electrolytes.
The synthesis of PEDOT in [C2 mim][NTf2 ] and [C4 mpyr][NTf2 ] has also been
studied, and the multiple redox peaks observed were influenced by the choice of
ionic liquid [80]. The current increase during potentiodynamic growth of the film in
the pyrrolidinium species was less than in the imidazolium analogue, suggesting
a slower film growth due to the higher viscosity and lower conductivity of this
medium that limits ion/molecule transport kinetics. This is as observed for growth
of poly(pyrrole) and poly(thiophene)s. Post-polymerization CVs of these films were
recorded in both acetonitrile/Bu4 NClO4 and the ionic liquid, and compared to
those of PEDOT grown from an acetonitrile solution. For both films grown from
ionic liquid there was an increase in the electrochemical activity upon cycling
in the acetonitrile solution, suggesting better swelling of the polymer and thus
faster transport of ionic species into and out of the polymer during cycling. Thus,
the observed activity reflects the electrochemical accessibility of the polymer to
the electrolyte, which may suggest that the electrochemistry of the polymer is
a surface-dominated phenomenon. On return to the [C4 mpyr][NTf2 ] there was a
rapid return to the lower charge capacity regime, which was ascribed to structural
changes. However, there was a progressive increase in current with cycling in the
[C4 mpyr][NTf2 ], as observed by Randriamahazaka et al. [64] and Wagner et al. [80]
for the cycling of PEDOT in [C2 mim][NTf2 ] after growth or cycling in acetonitrile,
which is likely to be linked to the slow dissolution of entrained acetonitrile out of
the film and/or the slow uptake of ionic liquid into the film. In agreement with
other authors [79], no memory effect upon cycling the films in these different
solvents was observed. The growth and post-polymerization CVs of PEDOT from
[C2 mim][NTf2 ] and from an acetonitrile solution are shown in Figure 7.11. There
is a decrease in the electrochemical activity of the film grown in the acetonitrile
solution when cycled in the ionic liquid, whereas the activity of the film grown in
the ionic liquid was less affected by the nature of the cycling solvent.
Comparison of the chronoamperograms recorded during EDOT electropolymerization in the two different ionic liquids and two conventional acetonitrile-based
electrolytes allows some conclusions to be drawn about the mechanism of polymer deposition of PEDOT from these different media (Figure 7.12) [80]. The current transients suggest that the process is initially much slower in the solution
190 7 Conducting Polymers
Fig. 7.11 Cyclic voltammograms of PEDOT (0.1 M, 20 cycles, every
third shown, 100 mV s−1 ): (a) Film I growth in acetonitrile/[Bu4 NClO4 ],
(b) post- growth of films I and II in acetonitrile/[Bu4 NClO4 ], (c) Film
II growth in [C2 mim][NTf2 ] and (d) post-growth of films I and II in
[C2 mim][NTf2 ], vs. a Ag pseudo-reference electrode [80].
containing Bu4 NClO4 as the electrolyte than in the other cases. Moreover, the
different shape of the curve suggests a different mechanism of deposition; the
current transient in acetonitrile/Bu4 NClO4 is indicative of progressive nucleation,
with a slower growth rate and thus lower currents, whereas the current transients
in the ionic liquids and the acetonitrile/LiNTf2 solution (Figure 7.12(b–d)) suggest
Fig. 7.12 Current–time responses to potential step from 0 to 1.4 V for
the electropolymerization of 0.1 M EDOT onto ITO electrodes in different media: (a) 0.1 M Bu4 NClO4 in acetonitrile, (b) [C4 mpyr][NTf2 ], (c)
0.1 M LiNTf2 in acetonitrile and (d) [C2 mim][NTf2 ] [80].
7.4 Characterization
instantaneous nucleation, thus indicating a strong influence of the anion on polymer growth. The spectroelectrochemistry of the films from these solutions was also
studied – see Section 7.4.3.
Danielsson et al. [25] have studied the synthesis of PEDOT in ionic liquids
that utilize bulky organic anions, 1-butyl-3-methylimidazolium diethylene glycol
monomethyl ether sulfate and 1-butyl-3-methylimidazolium octyl sulfate, the latter
of which is a solid at room temperature and thus requires the addition of either
monomer or solvent (in this case water) to form a liquid at room temperature.
Polymerization in a water-free ionic liquid was only possible in the octyl sulfate
species, but the polymerization of EDOT was successful in aqueous solutions of
both the ionic liquids (0.1 M). The ionic liquid anions appear to be mobile within
the polymer, exchangeable with chloride ions at a polymer/KCl(aq) interface, but
it is interesting that when the PEDOT is in aqueous solutions of the ionic liquid,
at higher concentrations (0.01–0.1 M) the imidazolium cation can suppress this
anion response. The ion mobility in both the ionic liquid and in the polymer film
in contact with the solution is significantly increased by addition of water.
7.3.2.4 Poly(para-phenylene)
Endres et al. [82] have demonstrated the suitability of an air- and water-stable
ionic liquid for the electropolymerization of benzene. This synthesis is normally
restricted to media such as concentrated sulfuric acid, liquid SO2 or liquid HF
as the solution must be completely anhydrous. The ionic liquid used, 1-hexyl-3methylimidazolium tris(pentafluoroethyl)trifluorophosphate, can be dried to below 3 ppm water, and this ionic liquid is also exceptionally stable, particularly in
the anodic regime. Using this ionic liquid, poly(para-phenylene) was successfully
deposited onto platinum as a coherent, electroactive film. Electrochemical quartz
crystal microbalance techniques were also used to study the deposition and redox
behavior of the polymer from this ionic liquid (Section 7.4.1) [83].
7.4
Characterization
It is particularly important to fully characterize the nature of the materials produced
when different growth media and methodologies are being investigated, and this is
perhaps one area of conducting polymer research that has generally lacked sufficient
attention. This is partly due to the insolubility of these materials, but there is still
a wide range of analytical techniques suitable, and a number of these have been
utilized for the analysis of conducting polymers synthesized in ionic liquids.
7.4.1
Electrochemical Characterization
Cyclic voltammetry of the conducting polymers allows assessment of the electrochemical activity of the films and the redox processes involved. The electrochemical
191
192 7 Conducting Polymers
response of the polymers can be strongly dependent on the solvent/electrolyte system used as the cycling medium. This is related to the extent to which the solvent
swells the polymer, which affects the size of response and capacitance measured.
Thus, post-polymerization CVs of the polymers reflect not only the properties of
the polymer film and any influence of the growth solvent used, but are also affected
by the nature of the solvent used for the post-polymerization cycling. It has been
noted that the post-polymerization CVs of polymers grown in ionic liquids can
be significantly different depending on whether they are cycled either in the ionic
liquid or in a molecular solvent/electrolyte system, or that the films may take time
to equilibrate on transferal to a new medium [64, 80]. This effect has primarily been
attributed to the different degrees of polymer swelling in the different media.
Electrochemical analysis can provide valuable information relating to the structure of the polymer. For example, Randriamahazaka et al. [64, 78, 79] have studied
the behavior of PEDOT in [C2 mim][NTf2 ], using a PEDOT film electrodeposited
from an acetonitrile/LiClO4 solution and cycled in the ionic liquid [79]. Using potential step experiments, the authors showed that the redox switching dynamics
of the polymer consisted of two different processes, with different kinetics. Both
a fast and a slow process were identified for the polymer oxidation and reduction,
and the authors note that this is consistent with the proposal that the structure of
the PEDOT films is composed of two coexisting zones, one rigid and compact zone
containing polymer chains that are long and highly conjugated, and one zone with
a more open polymer configuration with chains of a shorter conjugation length
[84], as also proposed by other authors [85, 86]. The presence of these two zones is
a suggested explanation for the presence of the two oxidation and reduction peaks
that are observed in the cyclic voltammogram of this polymer, with the lower potential peak arising from oxidation of the compact, highly conjugated polymer and
the more anodic peak ascribed to oxidation of the more open polymer with shorter
conjugation.
There is extensive and somewhat inconclusive discussion in the literature regarding the possible origin of the multiple redox peaks sometimes observed for
poly(thiophene) species. It has been proposed that these peaks are due to the transitions between the neutral, polaron, bipolaron and metallic states of the polymer
[87], which may also be influenced by the rate of counterion transport [88], the effect of “charge-trapping” [84], conformational changes accompanying radical cation
formation [89], consideration of the mechanical strain on the polymer that results
from the forced intrusion of anions into the film [71], as well as the aforementioned
reduction of different areas of the polymer film or of polymer chains of significantly
different lengths [86, 90, 91]. We have observed multiple redox peaks in a number
of conducting polymer films synthesized in ionic liquids and an additional possible
explanation for this, when an ionic liquid is used as the supporting electrolyte, is
the presence of additional redox processes involving the intercalation and expulsion
of the ionic liquid cation.
For further electrochemical characterization of the conducting polymer films,
the charge passed during the electrochemical synthesis can also be used to estimate the film thickness and level of doping, and thus the capacitance of the
7.4 Characterization
film if required. This may be more accurately achieved using an electrochemical quartz crystal microbalance (EQCM), to study both polymer growth and the
intercalation/expulsion of anions/cations from the film during cycling. In this
technique, the polymer is deposited onto a metal-coated quartz crystal (as the
working electrode, typically gold- or platinum-coated) whose oscillating resonance
frequency changes according to the mass of the polymer. Thus, both growth of
the polymer and its redox cycling (with the associated mass changes due to ion
incorporation/expulsion) may be studied. Endres et al. [83] have demonstrated
the applicability of this technique for the analysis of a conducting polymer in
ionic liquids, through the synthesis and cycling of poly(p-phenylene) in 1-hexyl3-methylimidazolium tris(pentafluoroethyl)trifluorophosphate. It should be noted
that when using this technique in combination with a viscous medium such as an
ionic liquid, the damping effect that the ionic liquid has on the frequency of the crystal must be considered [83]. However, if this damping does not change significantly
during the growth and cycling of the polymer, whereas the resonance frequency
does (by at least one order of magnitude more than the change in damping) then
under these conditions it is valid to convert the frequency change to changes in the
mass of the polymer using the Sauerbrey equation. Using this technique, Schneider
et al. [83] showed that, as a result of the large size of the anion used, cycling of the
polymer in the ionic liquid involved a significant amount of cation exchange, which
was particularly prominent at higher scan rates. At low scan rates there is sufficient time for movement of the large anion and this dominates, especially in the
cathodic peak region, whereas in faster scans anion removal or insertion into the
polymer film becomes more difficult. Outside the cathodic peak region, at low scan
rates anion and cation movement are approximately equal but, as the scan rate is
increased, movement of the ionic liquid cation becomes the dominant process over
the whole potential range.
Cyclic voltammetry is also an ideal analytical tool for assessing the electrochemical stability of the polymer films. This is a fundamental requirement for any
conducting polymer to be considered for long-term use in electrochemical devices.
The use of ionic liquids for the electrochemical cycling of poly(aniline) has been
reported to enhance lifetimes to over a million cycles [12], and significant improvements in the cycling stability of poly(pyrrole) have also been reported [32].
Electrochemical impedance spectroscopy is a powerful tool for elucidating the
diffusion and conduction parameters of the film, particularly when used in combination with computer modeling. Danielsson et al. [25] utilized this technique to
study parameters such as solution resistance, charge transfer resistance, double
layer capacitance and bulk redox capacitance for PEDOT grown in ionic liquids
composed of imidazolium cations and bulky anions, and assessed how these
were influenced by dilution of the ionic liquid to different concentrations in
water. Naudin et al. [26] also used this technique for the analysis of poly(3-(4fluorophenyl)thiophene) films grown in [1-ethyl-2,3-dimethylimidazolium][NTf2 ]
and [1,3-diethyl-5-methylimidazolium][NTf2 ]. In both these studies, the polymers
in the neat ionic liquids displayed slower ion transport compared to those in
molecular solvents or diluted ionic liquids. Fenelon and Breslin [29] also used
193
194 7 Conducting Polymers
electrochemical impedance measurements to demonstrate the conductivity and
stability of poly(pyrrole) deposited onto iron from [C4 mim][PF6 ].
The switching potential of the polymer – the potential at which there is a transition between conducting and insulating states – can be significantly influenced
by the nature of the solvent and electrolyte used for growth and cycling in molecular solvents [4], and thus this will also be influenced by the use of ionic liquids,
and the nature of the cations and anions used. Boxall and Oseryoung [68] determined the switching potentials of poly(pyrrole) and poly(N-methylpyrrole) from
[C4 mim][PF6 ], using rotating ring-disk voltammetry, to be 0.63 ± 0.04 and 1.07 ±
0.03 V, respectively, vs. the cobaltocenium/cobaltocene couple. This technique was
also used to study the potential-dependant conductivity of the polymers.
The DC conductivity of the polymer films is clearly an important property of
conducting polymer films and this can be accurately measured using a four-point
probe conductivity apparatus, or may be extrapolated from EIS measurements.
There are few reports in the literature documenting the conductivity of polymers
synthesized in ionic liquids. Sekiguchi et al. [65] measured the conductivities of
poly(pyrrole) and poly(thiophene) synthesized in [C2 mim][OTf], and compared this
to polymers synthesized in a dilute solution of this ionic liquid in acetonitrile
or water. The conductivities were measured using a two-probe method, with the
polymers collected as powders by scraping off the electrode. This technique gives
relatively low values but nonetheless showed that the conductivities of the polymers
from the neat ionic liquids were significantly higher than those from the acetonitrile or water solutions. Consistent with this, the doping level of the poly(pyrrole)
from the ionic liquid was significantly higher. In our preliminary investigations
into the synthesis of poly(pyrrole) from [NTf2 ]-based ionic liquids, without any optimization of the synthesis technique, we have measured conductivities up to ca.
100 S cm−1 .
7.4.2
Morphological Characterization
Very striking differences between films grown in conventional solvents and those
grown in ionic liquids have been observed using scanning electron microscopy
(SEM) [51, 65, 80, 92, 93]. Generally, the films grown from ionic liquids appear to
be considerably smoother, which may also result in improved conductivities.
SEM analysis of poly(thiophene) grown from [C2 mim][NTf2 ] and [C4 mpyr][NTf2 ]
reveals a slightly smoother morphology for the poly(thiophene) films from the
pyrrolidinium ionic liquid [27]. The influence of the nature of the ionic liquid on the film morphology is consistent for poly(thiophene), poly(pyrrole), poly
(bithiophene) and poly(terthiophene), although the difference is less marked in the
latter two.
The polythiophene film grown in the [C2 mim][NTf2 ] (Figure 7.13) displays a
‘packed grain’ structure commonly observed in polythiophene films, and is very
similar to that reported by Sekiguchi et al. [65], who noted that the grain size
was smaller than that obtained when acetonitrile was used as the solvent. This
7.4 Characterization
Fig. 7.13 SEMs of poly(thiophene) films grown from [C2 mim][NTf2 ]
(a)–(c) and [C4 mpyr][NTf2 ] (d)–(f), viewed from above (a), (b), (d)
and (e) and edge-view (c) and (f) [27].
grain morphology is typical of 3D nucleation and growth [94]. The film grown
from the [C4 mpyr][NTf2 ] (Figure 7.13) was slightly smoother, suggesting a more
ordered film, which is consistent with slightly slower film growth (Figure 7.8).
The edge-on views of the poly(thiophene) films (Figure 7.13) also show a dense
film on the ITO electrode below the granular polymer, clearly suggesting that a
different nucleation and growth mechanism operates at the onset of film growth,
as has been proposed by Schrebler et al. [95]. Thus, initial film growth is consistent
with an instantaneous 2D mechanism, involving soluble oligomer growth at the
ITO electrode and subsequent deposition at some critical chain length, to form a
compact film. This is then followed by progressive 3D nucleation and growth giving
the granular morphology, which is probably a result of the formation of a more
branched poly(thiophene). This is not unique to growth of polythiophene in ionic
liquids – similar behavior is reported in other solvents and electrolytes [95–97] and
for poly(pyrrole) grown by solution-surface electropolymerization from this ionic
liquid [69].
Poly(bithiophene) films from these two ionic liquids are morphologically similar
(Figure 7.14), even though the redox behavior (Figure 7.9) is markedly different,
suggesting that the dominant differences in the films produced are on an atomic or
sub-micron rather than macroscopic level. The morphology of the poly(bithiophene)
films appears to be similar to that described by Roncali et al. [74] who reported a thin
film on the surface of the electrode, covered by a thick brittle powdery deposit, from
the galvanostatic polymerization of bithiophene in acetonitrile. The nodular structures are smaller in the poly(bithiophene) films than in the poly(thiophene), which
is consistent with the formation of shorter chain polymers [73], but this does not
195
196 7 Conducting Polymers
Fig. 7.14 SEM images of poly(bithiophene) ((a) and (b), 10 :m,
(c) and (d) 2 :m edge-view) and poly(terthiophene) ((e) and (f),
10 :m) grown from [C2 mim][NTf2 ] (a), (c), (e) and [C4 mpyr][NTf2 ]
(b), (d), (f).
appear to result in inferior electrochemical activity (Figure 7.9). The edge-on views
(Figure 7.14) again shed some light on the polymer growth mechanism. In contrast
to the polythiophene films, the initial growth layer of both poly(bithiophene) films
is granular, suggestive of a 3D nucleation and growth mechanism, and this appears
to be followed by a second granular 3D nucleation and growth phase [96], reflecting
an influence of the starting monomer.
Poly(terthiophene) films from these ionic liquids are very smooth, with a spongy
morphology at the micron level (Figure 7.14(e, f)) similar to that described by Sezai
7.4 Characterization
Fig. 7.15 SEM images of reduced PEDOT films on ITO, prepared
by cyclic voltammetry (20 cycles, 100 mV s−1 ) from (a) 0.1 M LiNTf2
in acetonitrile, (b) [C4 mpyr][NTf2 ], (c) [C2 mim][NTf2 ] and (d) 0.1 M
Bu4 NClO4 in acetonitrile [80].
Sarac et al. [98] for poly(terthiophene) grown in acetonitrile, but without the large
amount of powdery deposit observed in the poly(bithiophene) films (Figure 7.14(a,
b)). Again, the film from the pyrrolidinium ionic liquid appears to be slightly more
compact.
SEM analysis of PEDOT films grown from these ionic liquids compared to films
grown from acetonitrile/ Bu4 NClO4 (Figure 7.15) [80] suggests that the effect of
changing the dopant anion can be a more significant influence than the use of the
ionic liquids as the film from acetonitrile/ Bu4 NClO4 is markedly different from
those grown in either the ionic liquids or the acetonitrile/LiNTf2 .
Reduction of a conducting polymer, with the simultaneous expulsion of anions,
is generally expected to result in films becoming more compact, and this can be
studied by SEM. For these PEDOT films, the morphology of the polymer grown in
acetonitrile/ Bu4 NClO4 became more compact upon reduction, but this effect was
not clearly observed in the other films [99].
197
198 7 Conducting Polymers
7.4.3
Spectroscopic Characterization
Raman and FTIR Spectroscopy are important techniques for analysing the level
of doping of polymer films, both across the entire surface and by depth profiling
using a cross-section of film [100]. Raman spectroscopy measures the vibrational
energies of molecules by irradiating the sample with monochromatic laser light and
measuring the scattered radiation, which differs in frequency from the incident light
by an amount determined by the Stokes shift of the molecule. Raman spectra of
conducting polymers can provide information about the identity of the polymer (for
example Figure 7.16), the degree of oxidation of the polymer, and the identification
and quantification of any Raman-active dopants within the film.
Similar but complementary information can be obtained using infrared spectroscopy. In situ Fourier transform infrared spectroscopy (FTIR) with attenuated
total reflection (ATR) allows changes in the structural and electronic properties
of polymers in the near-IR region to be studied during electrochemical reactions.
Damlin et al. [81] demonstrated the use of this technique to study the p-doping and
n-doping of PEDOT synthesized and cycled in [C4 mim][BF4 ] and [C4 mim][PF6 ].
Few conducting polymers can be both p-doped (by oxidation) and n-doped (by reduction). The poor stability of the reduced form of PEDOT in organic electrolytes
under atmospheric conditions has not allowed the investigation of the n-doping
of this polymer in organic solvents. However, the PEDOT films were successfully
n-doped (−0.9 to −1.95 V) and p-doped (−0.9 to +0.8 V) in the ionic liquids, as evidenced by the doping-induced infrared active vibrations. In the tetrafluoroborate
ionic liquid the current response of the n-doping cycles appeared to stabilize after
about five scans and remained stable for the next nine cycles. At a low degree of
doping (< 0.1 V) the changes in the doping induced infrared active vibrations of the
polymers were the same, irrespective of growth solvent, whereas during p-doping
at more positive potentials the polymers from the ionic liquids showed further increases in the intensities of these bands that were not observed in the polymer from
Fig. 7.16 Examples of Raman spectra of (a) poly(pyrrole) and (b)
poly(terthiophene), from [C2 mim][NTf2 ].
7.4 Characterization
acetonitrile. The in situ FTIR-ATR spectra of the polymer from the ionic liquids
were similar to those from 0.1 M Bu4 NClO4 in acetonitrile, with only slight shifts
in peak position. However, the relative intensities of the bands of the n-doped and
p-doped polymer are higher in the polymer from the [C4 mim][BF4 ], suggesting that
it is easier to create charge carriers in this film compared with those grown from
the [C4 mim][PF6 ] or the acetonitrile solution [81].
The p-doping and n-doping of conducting polymers, and the resultant incorporation of cations/anion, can also be studied by X-ray photoelectron spectroscopy
(XPS). This is a relatively surface-specific technique, where the photoelectrons from
the solid (which have quite a short range) give rise to spectra with peaks of binding
energies specific to the elements present. For each element, the binding energies
are influenced by chemical bonding or oxidation state (allowing, for example, the
nitrogen of the poly(pyrrole) to be identified separately from the negative nitrogen
of the [NTf2 ]− anion, as shown in Figure 7.17), thus deconvolution of the peaks
allows structural analysis of the polymer, and the relative intensities of the peaks
allows complete compositional analysis of the sample.
However, for analysis of conducting polymers grown in ionic liquids the spectra
may be complicated by the presence of elements common to both polymer and
growth medium. For example, poly(pyrrole) grown in an imidazolium NTf2 ionic
liquid will give peaks in the N 1s core level spectra from the imidazolium cation
Fig. 7.17 Example of the deconvolution of peaks from XPS analysis of
poly(pyrrole) containing the [NTf2 ] anion and the [C2 mim] cation.
199
200 7 Conducting Polymers
and the anion as well as the polymer itself. Further, there may be residual ionic
liquid on the polymer surface not removed by washing, as observed by Naudin et
al. [26], in addition to ionic liquid cations and anions that are actually intercalated
into the polymer film. However, some of this complication may be avoided by
judicious choice of ionic liquid, such as the use of the [BF4 ]− or [PF6 ]− anion or a
phosphonium cation. Naudin et al. [26] demonstrated the use of XPS for analysis
of poly(3-(4-fluorophenyl)thiophene) grown and cycled in ionic liquids and its use
for determining the nature and quantity of the dopant species in the p-doped and
n-doped polymer.
UV–Vis spectroscopy is a cheap and readily available technique that is ideal for
studying conducting polymer films at each stage of growth and reduction/oxidation
(when deposited onto ITO glass). As a result of electronic transitions between the
fundamental levels and polaronic or bipolaronic levels, which involve photons with
energy in the visible region of the spectrum, the UV–Vis spectra of conducting
polymer films are significantly altered on oxidation (films change color). The wavelength maxima position changes as a function of the degree of oxidation/reduction,
hence this is a valuable analytical technique for probing the electronic structure of
the films.
Damlin et al. [81] utilized this technique to study the p-doping and n-doping of
PEDOT synthesized and cycled in [C4 mim][BF4 ] and [C4 mim][PF6 ]. They reported
that the n-doping of PEDOT in the ionic liquids has little effect on the UV–Vis
spectrum, which can also be the case when organic solvents are used, whereas the
p-doping results in significant changes, suggesting the existence of different types
of charge carriers in the different doping regimes (polarons vs. bipolarons). The
UV–Vis spectra also suggested little effect on the average conjugation length of
the polymers from using the ionic liquid, which is indicated by the position of the
π–π * transition that occurs around 560 nm.
Fenelon and Breslin [29] demonstrated the use of spectroelectrochemistry to
monitor the growth of poly(pyrrole) from [C4 mim][PF6 ] onto an iron working electrode by plotting the natural logarithm of the absorbance of the 450 nm transition
(attributed to the electronic transition from the valence to the antipolaron band of
the polymer) against time. The authors thereby demonstrated that the growth of the
poly(pyrrole) was a two-step process, starting with a faster nucleation and growth
step (for approximately 8 min) followed by a steady growth phase (approximately
30 min), both of which obeyed first-order kinetics.
Using this technique a difficulty in fully reducing conducting polymer films in
ionic liquids has been observed [80], as indicated by the presence of a significant
free carrier electron band in the spectrum at high wavelengths, which may be
due to poor solvent swelling or a result of cation incorporation rather than anion
expulsion during reduction (Figure 7.18). Spectroelectrochemistry (Figure 7.18(c))
of the PEDOT film electrodeposited from [C4 mpyr][NTf2 ] also shows incomplete
reduction, and the reduced spectra do not change, even with reduction potentials
down to −1.6 V. The appearance of the shoulder at around 900 nm, in both Figure
7.18 (b) and (c), is consistent with an incompletely reduced material [101].
7.4 Characterization
Fig. 7.18 UV–Vis spectra of PEDOT films
(a) deposited in [C2 mim][NTf2 ], oxidized
20 min at 0.8 V and reduced 20 min at
−1 V in the ionic liquid, (b) grown in
[C2 mim][NTf2 ], oxidized 20 min at 0.8 V and
reduced 20 min at −1 V in acetonitrile/0.1 M
Bu4 NClO4 solution, (c) spectroelectrochemistry of PEDOT in 0.1 M LiNTf2 in acetonitrile after deposition from [C4 mpyr][NTf2 ]
from −1 to 0.9 V [80].
7.4.4
Solid-state NMR
One of the primary hindrances to improving understanding of the influence of the
ionic liquid on conducting polymers, and studying the incorporation of ions from
these media into polymers, is the difficulty in analyzing such insoluble materials.
Understanding the ion incorporation processes is paramount to the development
of these systems in, for example, actuator devices. The use of nuclear magnetic
resonance (NMR) spectroscopy to study the nature of the polymers produced is a
valuable but under-utilized technique. This can not only provide information on
the polymer backbone but can also allow the study of intercalated ions within the
film [102–104]. Judicious choice of ionic liquid cation and anion can allow them to
be detected independently within the polymer film. For example, utilization of an
ionic liquid comprised of a phosphonium cation, to allow its detection by 31 P NMR,
and the NTf2 anion, for study by 19 F NMR spectroscopy [105]. The intercalation of
201
202 7 Conducting Polymers
the bulky phosphonium cation is potentially of interest for use in actuator devices.
In addition, the polymer backbone can be studied by 13 C NMR [103, 104].
Poly(pyrrole) films grown by constant potential in the ionic liquid tri(hexyl)
(tetradecyl) phosphonium bis(trifluoromethanesulfonyl)amide, [P6,6,6,14 ][NTf2 ], and
subsequently cycled 100 times and reduced in the monomer-free ionic liquid,
showed the presence of both ionic liquid cations and anions within the film. Interestingly, poly(pyrrole) films grown by constant potential in the ionic liquid but
with no subsequent electrochemical cycling also contained both ionic liquid cation
and anion within the film (Figure 7.19).
The 31 P NMR (Figure 7.19(a)) indicates incorporation of the phosphonium cation
into the film. The dominant signal appears at −12 ppm, with a minor peak at
32 ppm. This latter peak is assigned to the phosphonium cation of residual ionic
liquid on the surface of the film, which was not removed by washing but is not
actually intercalated into the film. This shift is consistent with that reported by
Bradaric et al. [106] and with our own analysis of the neat ionic liquid. The large
peak at −12 ppm is assigned to phosphonium cations intercalated into the polypyrrole film, and this shift is consistent with a P–N interaction. The dramatic change
in the chemical shift of the phosphorus compared to the neat ionic liquid indicates a significant change in environment on incorporation into the polymer film.
Fig. 7.19 (a) 31 P (b) 19 F and (c) 13 C NMR spectra of polypyrrole
grown at constant potential in [P6,6,6,14 ][NTf2 ].
7.5 Future Directions 203
However, the DC conductivity of the pyrrole film was sufficiently high to indicate that the polymer chain was intact and this precludes the possibility of any
chemical reaction between the cation and the polymer. The successful expulsion
of cations from the film (see below) also eliminates this possibility. The 19 F NMR
spectrum (Figure 7.19(b)) shows the presence of intercalated anions within the film,
as indicated by the peak at −82 ppm that is consistent with [NTf2 ]− . The 13 C NMR
spectrum of the film (Figure 7.19(c)) shows a broad resonance from the poly(pyrrole)
centered at ca. 120 ppm, and also significant intensity between 0 and 30 ppm from
the alkyl chains of phosphonium cations within the film [106].
NMR analysis of a poly(pyrrole) grown by constant potential from a 0.1 M solution of the [P6,6,6,14 ][NTf2 ] in acetonitrile also showed the presence of both the
ionic liquid cation and anion within the film. However, although the phosphonium
cations appear to be easily incorporated into the poly(pyrrole) during synthesis, initial results suggest that they are not easily incorporated during cycling. Poly(pyrrole)
films grown at constant potential from a 0.1 M solution of LiNTf2 in acetonitrile
and then cycled 100 times in the neat [P6,6,6,14 ][NTf2 ] contained no detectable phosphonium cations. This may be related to a lack of solvent swelling of the film in the
ionic liquid, which restricts ion movement.
This technique can also be used to investigate the ease of expulsion of anions from
the polymer. It was initially thought that given the large size of the phosphonium
cations they would remain incorporated in the poly(pyrrole), as is observed for bulky
anions such as polyelectrolyte dopants, but films grown in the ionic liquid and then
oxidized overnight in the ionic liquid showed no detectable phosphonium cations.
However, poly(pyrrole) oxidized for only 1 h still contained significant amounts of
phosphonium cations, thus it takes considerable time for the films to fully oxidize
through incorporation of NTf2 anions and expulsion of the phosphonium cations
(the current becomes negligible after about 5 h). However, poly(pyrrole) films grown
in the ionic liquid then oxidized in a 0.1 M solution of LiNTf2 in acetonitrile for 4 h
resulted in expulsion of all of the phosphonium cations from the film (the current
was reduced to a background level in less than 1 h). This was attributed to the effect
of solvent swelling of the polymer, which enables increased ion movement in and
out of the film. This may also explain the observed lower electrochemical activity of
the films in the ionic liquid compared with their activity when cycled in molecular
solvent systems. It is important to note, however, that the importance of solvent
swelling would be significantly reduced for thinner polymer films.
7.5
Future Directions
7.5.1
Chiral Ionic Liquids
There is significant interest in the formation of chiral conducting polymers, such
as chiral poly(aniline), as a result of their potential applications in chiral sensors,
204 7 Conducting Polymers
for chiral separations and so on. One route to these materials is by incorporating
a chiral dopant anion during the electrochemical polymerization [107]. Thus, if
the polymer was synthesized in a chiral ionic liquid, of which there are a growing
number [108], then formation of an optically active conducting polymer might
result.
7.5.2
Protic Ionic Liquids
One family of ionic liquids that has to date only been sparsely investigated for
use with conducting polymers is the protic ionic liquids, where the cation has
one or more mobile hydrogen atoms. A recent manuscript by Biçak detailed the
synthesis of 2-hydroxy ethylammonium formate [109], which melts at −82 ◦ C and
has a room-temperature ionic conductivity of 3.3 mS cm−1 , and reported the ability
of this protic ionic liquid to dissolve poly(aniline) (17 g mL−1 ) and poly(pyrrole) (no
concentration specified). The dissolution of conducting polymers into any solvent is
of significant interest for a variety of reasons, such as improving their processability
and ease of incorporation into different devices.
Li et al. [93] have used 1-ethylimidazolium trifluoroacetate, which is a Brønsted
acidic ionic liquid, as a medium for the electropolymerization of aniline. They
report that in this ionic liquid the oxidation potential of aniline is lower (0.58 V
compared to 0.83 V in 0.5 M H2 SO4 ) and that the growth rate of the polymer is increased. Further, the resultant films are smooth, strongly adhered to the Pt working
electrode and are very electrochemically stable. Similar results have been reported
by Liu et al. [92], who found that this was the best ionic liquid for the polymerization of aniline, compared to the unsatisfactory results observed in other protic
ionic liquids 1-butylimidazolium tetrafluoroborate, 1-butylimidazolium nitrate and
1-butylimidazolium p-toluenesulfonate, as well as the 1-butyl-3-methylimidazolium
hydrogen sulfate and 1-butyl-3-methyimidazolium dihydrogen phosphate.
However, this family of ionic liquids holds an additional attraction, which is the
potential to create “distillable ionic liquids.” Earle et al. [110] have recently reported
that a range of ionic liquids that are commonly perceived as nonvolatile, including
various imidazolium NTf2 − salts, can actually be distilled at low pressure and temperatures of 200–300 ◦ C without significant decomposition. In such a process the
ionic liquids are transferred into the gas phase as ionic species. However, in the
case of protic ionic liquids, hydrogen transfer between the cation and anion can
allow distillation of the neutral components and then subsequent recombination
to reform the ionic liquid as the distillate [111]. The primary amine reported by
Bicak [112], 2-hydroxy ethylammonium formate, decomposes on heating to give a
formamide, but the combination of the formate anion with tertiary amine cations
such as N-methylpyrrolidinium can give an ionic liquid that is distillable at modest
temperatures and pressures [111]. The use of distillable ionic liquids for the synthesis and use of conducting polymers may provide additional advantages in terms
of easy removal of the ionic liquid and isolation of any oligomeric species from
the growth solutions. Further, as this area of ionic liquid research is relatively new,
7.5 Future Directions 205
there is a considerable range of protic ionic liquids as yet undiscovered or still to
be investigated for their use for the growth and synthesis of different conducting
polymers, and the different acidities of these ionic liquids may be of particularly
interest for the synthesis of poly(aniline).
7.5.3
Nano-dimensional Polymers
The increasing interest in the use of conducting polymers [2] is fuelling a continued
need for materials with improved physical and chemical properties. In particular,
there is a recent drive towards nanostructured and reduced dimensionality materials, such as thin films, nanotubes, wires, particles and so on, which can exhibit
markedly different properties from those of the bulk materials [113, 114]. There is
already a small number of reports of the use of ionic liquids for the electrochemical synthesis of nanostructured conducting polymers and research in this area is
predicted to increase significantly in the coming years. Koo et al. [31] have reported
the polymerization of pyrrole using a nanoporous aluminum oxide template in
[C4 mim][BF4 ], which yields poly(pyrrole) nanotubes and nanowires. Poly(aniline)
nanotubules have been made by electrochemical polymerization onto an ITO glass
electrode from [C4 mim][PF6 ] containing 1 M trifluoroacetic acid [115].
7.5.4
Chemical Polymerization
The choice of synthetic route for the production of conducting polymers, either
through electrochemical or chemical oxidation of the monomer or, in rare cases,
photopolymerization or enzyme-catalysed polymerization, is primarily dictated by
the final application of the polymer. While electrochemical polymerization is widely
used for the controlled synthesis of polymer films, chemical polymerization results
in the formation of powders or colloidal dispersions, is much more amenable to
scale-up and can be used as a means to coat nonconducting substrates with conducting polymers. Although outside the scope of this discussion, it is interesting to note
that despite the increased interest in the use of ionic liquids for the electrochemical
synthesis of conducting polymers, there is a significant dearth of investigations into
the potential benefits of using ionic liquids for the synthesis of these materials via
alternative routes. The good electrochemical stability and solubilising properties of
ionic liquids should allow access to a plethora of monomers and chemical oxidants
at significant concentrations, some of which are either insoluble in, or outside
the electrochemical stability of, molecular solvents. Gao et al. [116] have reported
the chemical synthesis of poly(aniline) at a water/ionic liquid interface, in which the
polymer grows into the water phase as nanoparticles, and Biçak et al. [109] have
also reported the use of their aforementioned protic ionic liquid, 2-hydroxyethyl
ammonium formate, for the chemical synthesis of organo-soluble poly(aniline).
An initial investigation into the use of [C2 mim][NTf2 ] for the chemical synthesis of
poly(pyrrole), poly(terthiophene) and nano-dimensional poly(thiophene) has also
206 7 Conducting Polymers
been reported [117]. However, the range of different monomers, oxidants and ionic
liquids available, as well as different experimental techniques that may be employed, for example with the ionic liquid as one phase or in a biphasic system in
combination with a second solvent, suggests that this may be a particularly fruitful
area of future investigation for conducting polymer researchers.
7.5.5
Remaining Challenges
The potential improvements that ionic liquids may impart to conducting polymers
have been widely discussed – increased doping levels, smoother films, increased
conductivity, decreased over-oxidation and improved electrochemical stability and
so on. However, the research to date in this area has only just begun to investigate
these hypotheses and demonstrate any material advantages in the use of ionic
liquids; future directions in this area must focus on some of these issues in addition
to simply demonstrating the use of new ionic liquids for conducting polymer
synthesis.
The influence of the ionic liquid on the polymerization process itself is yet to be
clarified. For example, does the ionic nature of these media stabilize the radicals/
cations that are formed during the polymerization and, if so, how does this impact
on the mechanism of polymer growth, the resultant chain length, conjugation
length and so on? The solubility of the oligomers that are formed during the
polymerization in the ionic liquid, and the extent to which they diffuse away from
the electrode (which may be less in viscous ionic liquids) influences the length at
which they precipitate onto the electrode. This will impact on polymer properties
such as conjugation length, morphology, conductivity and so on, and therefore
warrants some investigation. The solubility of the oligomers formed during the
polymerization is also of interest not just from a mechanistic point of view but
because the solubilization of conducting polymers, even if they are relatively short
chained, is desirable for a number of applications.
It would also be interesting to measure the degree of solvent swelling of the
polymer films in ionic liquids compared to molecular solvents, as this will impact
significantly on the ion mobility and thus the electrochemical activity of the polymer. It is expected that this will be lower in ionic liquids compared to molecular
solvents – although the extent of this difference may also depend on the thickness
of the film – but this is yet to be quantified.
The extent of the ionic liquid cation and anion intercalation into the film during
growth and cycling, and the structural features of the ions that influence this,
requires more investigation. It has been postulated that the higher ion concentration
of ionic liquids compared to a traditional electrolyte/molecular solvent system
will result in a higher polymer doping level, but this has not been extensively
demonstrated. Anion and cation intercalation not only influences the properties of
the polymer but is also important for understanding and developing conducting
polymer–ionic liquid actuator devices.
7.6 Conclusions 207
The smoother film morphology imparted by use of ionic liquids has been attributed to slower film growth in this medium but film growth rate has not been
quantified or compared to growth in molecular solvent systems. EQCM would be
an ideal technique for such a study. The high viscosity of ionic liquids could, if desired, be reduced by using higher temperatures, mixtures of ionic liquids or ionic
liquids diluted with either a molecular solvent or the monomer itself (if liquid),
and this will, in turn, influence the film morphology. The addition of a second salt
component, such as Bu4 N PF6 , to the ionic liquid may also be beneficial for film
growth.
7.6
Conclusions
It would be most interesting at this point to be able to compile a list of the benefits
of using ionic liquids for the synthesis and use of conducting polymers, and weigh
these against the disadvantages of these new media. However, such a clear idea of
the pros and cons of using ionic liquids is still some time away. The benefits of
using ionic liquids as the supporting electrolyte in conducting polymer devices has
been clearly demonstrated and predominantly concerns extended lifetimes and the
potential to reduce problems such as solvent evaporation that are associated with
the use of molecular solvents in these devices. However, at this point even this
research is predominantly at a lab-based scale rather than a larger or commercial
scale. Concerns such as ionic liquid cost, toxicity and large-scale availability will
no doubt come to the fore as ionic liquid-containing conducting polymer devices
are developed on a larger scale, whereas it may take longer for benefits such as
improved lifetimes or performance to be fully realized.
When the ionic liquid is used as the growth medium for these materials, financial
concerns may be minimized by efficient recycling after use, and toxicity concerns
will be confined to their use in the production process rather than their widespread
use in devices for public use, for example in solar cells, batteries or sensors. Some
properties of ionic liquids, such as their higher viscosity and lower conductivity
compared to molecular solvents, may be a disadvantage for larger-scale conducting
polymer synthesis, but the potential benefits of using ionic liquids lie not in the
production process but in their ability to improve the conducting polymer itself, for
example by improved conductivity, electrochemical activity, morphology and so on.
This will, in turn, give better device performance. Thus, it is paramount that the
efficacy of ionic liquids as a route to conducting polymers with improved material
properties is demonstrated.
At this point, the application of ionic liquids for the electrosynthesis of conducting
polymers has been demonstrated by a number of authors and some differences
between this and the use of molecular solvents reported. In a number of these
cases an improvement in properties was reported (most extensively a smoother
polymer morphology and increased cycle life) but the full range of benefits of using
ionic liquids is yet to be fully realized or amply demonstrated. There is clearly
208 7 Conducting Polymers
massive scope for future investigation, utilizing the vast range of ionic liquids
presently available either commercially or through reported synthetic routes. Thus,
while the field of ionic liquids is immense, their potential use for the synthesis of
the different types of conducting polymers is even more extensive. A number of
common ionic liquid anions have already been proven to be beneficial dopants for
conducting polymers but there are also an extensive number that are still to be tested
and maybe one of these new ionic liquids will prove to be the key to synthesizing
conducting polymers with all of the physical and electrochemical properties that
we desire.
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211
213
8
Nanostructured Metals and Alloys Deposited from
Ionic Liquids
Rolf Hempelmann, and Harald Natter
8.1
Introduction
Nanomaterials are composed of structural entities – isotropic grains or particles,
rods, wires, platelets, layers – of size, at least in one dimension, between 1 and
100 nm [1–3]. Larger particles are called submicron particles, smaller ones are
known as clusters. Some physical properties of nanomaterials [4–6] differ from
those of coarse-grained materials of the same chemical composition due to two
essential features:
1. Large specific surface area and concomitantly large specific surface energy; hence
surface sensitive properties (like catalytic activity) are enhanced [7–10] and processes where the surface energy is the driving force (sintering, grain growth) are
facilitated [11–13].
2. Quantum size effects; famous examples are the color shift upon size reduction of semiconductor nanoparticles like CdSe [14–16], and the surface plasmon
resonance of metallic nanoparticles like gold [17, 18].
The term nanomaterial includes materials consisting of or containing individual nanoparticles, nanorods, nanowires or nanoplatelets (for instance as composites), materials in the form of thin layers or coatings, and compact polycrystalline
materials with grain sizes below 100 nm; the latter contain an appreciable volume fraction of interphase or grain boundary regions (analogously to the surface
region of nanoparticles) and correspondingly a large specific interphase energy,
such that these bulk nanomaterials are said to be dominated by their interphases
[19]. Magnetic and mechanical properties are examples of properties which, even
for coarse-grained materials, depend on grain boundaries and other lattice defects,
in spite of their tiny volume fraction [20–22]. In bulk nanomaterials, with their
large volume fraction of grain boundaries and, particularly with respect to the
above-mentioned properties, pronounced nanoeffects occur which is the reason
for the academic and industrial interest in bulk nanostructured materials.
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
214 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
Often the starting point of the route to bulk nanomaterials is a powder of nanoparticles which can be prepared either “top-down” by milling or “bottom-up” by controlled chemical synthesis [23]. This powder has to be compacted/densified in order
to get bulk nanomaterial. An example of such a route is the famous inert gas condensation (IGC) introduced by Gleiter and Birringer [24]: a metal, for instance
Pd, is evaporated under He gas, in the gas atmosphere condensation starts with
the formation of clusters and small nanoparticles; these are deposited on a cold
finger (77 K), scraped off, collected, and uniaxially cold-pressed into the form of a
tablet. Isostatic hot-pressing of metal oxide nanoparticles is a route to nanoceramics
[25–28]. In all these routes the densification step is crucial. Efficient densification
(sintering) involves diffusion processes and requires elevated temperatures. However, at elevated temperatures grain growth takes place and after a short time the
sample is no longer nanocrystalline. At lower temperatures the sample remains
nanocrystalline but the densification is incomplete; the resulting porosity is the
main disadvantage of all routes to bulk nanomaterials which start from powders.
This critical compaction step is avoided in the case of the electrochemical route of
pulsed electrodeposition (PED) [29] which transforms cations, i.e. atomic species,
directly into nanomaterials without the detour via nanoparticles. In this way densities up to 99% of the theoretical value can be achieved, such that these materials
exhibit, for instance, intrinsic mechanical properties and not those dominated by
voids.
Furthermore, this technique allows variation of the grain size [30–33]; this is
important because many chemical and physical properties of nanostructured materials depend on the grain size. Only by variation of the crystallite size – this is a
novel aspect in materials science and technology [34] – is it possible to tune and
hopefully improve certain physical properties of one and the same material: for example, the enhanced hardness of nano-Au, the toughness of nano-Ni/P alloys [35],
the soft magnetic properties of nano-Ni [36] and the resistance of nanostructured
materials [37, 38] promise industrial applications [39–41].
The production of such “tailor-made” nanomaterials by electrochemical procedures is advantageous because the two crucial steps in nanocrystal formation –
nucleation and growth of nuclei – can be controlled by physical (current, voltage,
time, temperature) and chemical (grain refiners, complex formers) parameters
during the deposition process [42, 43].
In Section 8.2 the basics of pulsed electrodeposition (PED) will be described
for the case of aqueous electrolytes which allow the deposition of comparatively
noble metals like Cu, Ni, Pd, or Au; less noble metals like Fe or Zn can still be
electrodeposited from aqueous electrolytes because they exhibit a comparatively
large overpotential for hydrogen evolution. However, the main limitation of aqueous electrolytes, of course, is their narrow electrochemical window which adversely
affects the electrodeposition of metals like Al or Ta. Therefore, recently, the PED
technique has been extended to ionic liquids as electrolytes. General electrochemical aspects of ionic liquids can be found in Ref. [44]; here, in Section 8.3, we will
only address the technical aspects with respect to PED. Examples of nanometals and
nanoalloys electrodeposited from chloroaluminate-based ionic liquids are given in
8.2 Pulsed Electrodeposition from Aqueous Electrolytes 215
Section 8.4. The main disadvantage of these electrolytes is their extreme sensitivity
to moisture [45]; so-called ionic liquids of the third generation are water stable
[46], and examples of electrodeposition from these water-stable ionic liquids are
presented in Section 8.5. Finally, a short summary and outlook for possible future
developments is given in Section 8.6.
8.2
Pulsed Electrodeposition from Aqueous Electrolytes
The pulsed electrodeposition technique (PED) is a versatile method for the preparation of nanostructured metals and alloys [47]. In the last two decades PED has
received much attention worldwide because it allows the preparation of large bulk
samples with high purity, low porosity and enhanced thermal stability.
8.2.1
Fundamental Aspects
The electrochemical deposition of nanostructured metals and alloys is a two-step
process:
1. The formation of a large number of nuclei
2. The controlled growth of the deposited nuclei.
These two conditions can be realized by the proper choice of the chemical and
physical process parameters. The size and the number of nuclei can be controlled
by the overvoltage (η):
r =
2σ V
z e 0 |η|
(8.1)
In this electrochemical version of the Kelvin equation [48] r is the critical nucleation
radius, σ the specific surface energy, V the atomic volume in the crystal and z the
number of elementary charges e0 . The message of Eq. (8.1) is: the higher the
overvoltage the smaller the formed nuclei. A large overvoltage brings about a large
current density and thus a large rate of formation of nuclei. So PED is a deposition
process at large overpotential, a rare type of study in electrocrystallization because,
in most cases, underpotential deposition is studied [49–54]. This high overpotential
and the concomitant high deposition rate, however, can be maintained only for a
few milliseconds (ton -time) because the metal ion concentration in the vicinity of
the cathode decreases drastically and the process threatens to become diffusioncontrolled. In order to avoid this and avoid losing control over the process via Eq.
(8.1), the current has to be switched off, one has to wait for 20–100 ms (toff -time) in
order that the metal ions can diffuse from the bulk electrolyte to the cathode and
216 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
compensate the metal ion depletion there. Then the next current and voltage pulse
is applied, and so on: this is the reason for the pulsing.
In the electroless break between two pulses exchange currents flow, an undesired
electrochemical phenomenon because it induces the so-called Ostwald ripening
[55]: the larger crystallites (with the lower surface energy) grow at the expense of the
smaller ones (with the higher surface energy), and the crystallite size distribution
broadens and shifts to larger size values, i.e., deteriorates. Another origin of Ostwald
ripening could be surface diffusion of metal adatoms.
In view of this deposition mechanism several measures have to be taken to
control the crystallite size:
1. The voltage/current density in the peak has to be large; a reasonable value is
1 A m−2 in the more common galvanostatic mode.
2. The ton -time has to be as short as is necessary to avoid the diffusion control
regime and as long as possible in order to reach sufficient deposition efficiency;
a reasonable value is 3 ms.
3. The toff -time has to be as long as necessary for the material transport in the
electrolyte to take place, but not longer in order to minimize Ostwald ripening
and reach sufficient deposition efficiency.
4. The use of organic additives (grain refiners): they have to adsorb on the fresh nuclei, with a Gibbs enthalpy of adsorption just sufficient to suppress the exchange
current in the toff -time but insufficient to impede the metal ion deposition in the
ton -time, i.e., in the next pulse. This enables one to control the crystallization process during the toff -time because these molecules are adsorbed reversibly on the
electrode surface and hinder the surface diffusion of the adatoms. For different
metals different additives are most suited, but a typical and widespread example
is thiourea. The use of additives is common practice in the galvanic industry, the
selection of additive is a matter of experience, i.e., purely empirical [41].
5. Temperature influences all diffusion processes (cation diffusion in the electrolyte, surface diffusion of the adatoms). If small crystallite sizes are desired the
deposition should be performed at ambient temperatures or below in order to
slow down the kinetics of recrystallisation of the nuclei.
The bath composition, the pH-value, the hydrodynamic conditions and also the
use of special current pulse shapes are further possibilities to influence the deposition process. It is advantageous to perform the pulsed electrodeposition in
the galvanostatic mode because the average deposition rate can be simply derived
from
Iav =
Ipulse ton
ton + toff
(8.2)
In addition, there is better control of the current efficiency and the alloy composition. The potentiostatic mode would be, on the basis of Eq. (8.1), more desirable
but is experimentally more difficult to realize because a three-electrode set-up is
8.2 Pulsed Electrodeposition from Aqueous Electrolytes 217
necessary and, upon the voltage jump, the current should start theoretically from
an infinite value, which is not feasible due to electronic limitations. Also, in the
potentiostatic mode a short reverse pulse would be necessary to fix the initial potential, which is not generally desirable because passivation may occur during this
reverse potential pulse. Suitable substrates for the electrodeposition are stainless
steel or titanium electrodes (20 × 20 mm, distance between the electrodes 25 mm);
because of the poor adhesion the resulting deposits can be removed mechanically
from the electrode.
Tuning the grain size must be accompanied by measuring the grain size. The
most convenient method is the line width or line shape analysis of X-ray diffraction peaks according to Scherrer [56] (estimate of the grain size), or according to
the Williamson and Hall procedure [57] which allows one to separate grain size
and microstrain effects on the line width. Fourier transform techniques like the
Warren–Averbach technique [58–60] or certain full profile fit routines [61] even
allow determination of the crystallite size distribution. Details can be found in
Ref. [62].
8.2.2
Nanometal Deposition with Nano-Gold as an Example
For n-gold, deposited from a commercial gold(I)sulfite bath, the influence of the
deposition parameters on the nanostructure has been demonstrated [63]. For a
better comparison of the different results the experiments were performed at an
average current density (Iav ) of 3 mA cm−2 , see Eq. (8.2). If ton and toff are kept at
constant values (for details see Figure 8.1), the smallest crystallite size of 12 nm
is found for Ipulse = 0.5 A cm−2 (see Figure 8.1); for higher current values powder
Fig. 8.1 The crystallite size dependence of the pulsed current density
for gold deposits deposited from a commercial sulfite bath without
any additives [29].
218 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
Fig. 8.2 The effect of the toff time on the nanostructure of gold deposits. An average current density of 3 mA cm−2 was used for all experiments [29].
formation on the electrode surface is observed. The effect of the toff -time on the
nanostructure of the deposits is shown in Figure 8.2. As expected the smallest
crystallites can be found for the shortest toff -times, in accordance with the deposition mechanism outlined above (recrystallization process of the nuclei during
the toff time). To study the effects of organic additives on the nanostructure of the
deposit, butanediamine (a molecule with free amino groups), diammonium-EDTA
(a chelating complex former) and saccharin (a molecule with a sulfur group) were
compared. The electrolytes free of additives yield crystallite sizes in the range of
100 nm. With very small amounts of additives a strong reduction (by a factor of
about 2) in the crystallite size is observed, see Figure 8.3. Increasing the concentration of the grain refiners does not cause a substantial further decrease in the
crystallite size. The three additives show the same effect, but the degree of grain
refining is substance-specific: as is well known, gold reacts strongly with sulfur
groups. Similar results were obtained for the system nano-copper/citric acid [64],
nano-nickel/saccharin [61, 65] and nano-Pd/Na2 EDTA [42]. A detailed description
of the influence of organic additives on the microstructure of metal deposits is given
by Fischer [66]. Further nanometals with crystallite sizes between 10 and 100 nm
prepared by pulsed electrodeposition from aqueous electrolytes are: Pd [42, 67],
Fe [61], Co [68], and Cr [69]. Detailed information on the preparation and physical
properties are given in the cited references.
8.2.3
Nanoalloy Deposition with Fex Ni1–x Alloys as an Example
The deposition of Fex Ni1−x alloys is of industrial interest because these materials
find applications in electronic devices (e.g. computes hard disk). The most popular
8.2 Pulsed Electrodeposition from Aqueous Electrolytes 219
Fig. 8.3 The activity of different grain refiners (butanediamine, ammonium ethylenediamine tetra-acetic acid, benzosulfimide) on the
nanostructure of gold deposits [29].
alloys are Permalloy (soft magnetic properties) and Invar (very low thermal expansion). The polycrystalline compounds can be prepared by melting processes or by
direct current plating [70]. The magnetic and mechanical properties of these alloys
can be designed by nanostructuring. In the case of alloy deposition the bath composition is an additional process parameter which can influence the nanostructure
of the deposit. An electrochemical DC current procedure was reported by Cheung
et al. [71]. One condition for preparing homogeneous alloys by electrochemical
methods is nearly equal electrode potentials of the components. Fe and Ni exhibit
standard reduction potentials of −0.44 and −0.22 V; in view of these similar values
alloy formation can be expected. Ni exhibits a more positive standard reduction
potential than iron and therefore the Ni/Fe ratio in the deposited alloys should be
higher than in the electrolyte. Actually, the literature reports opposite experimental
results [70]. This anomalous codeposition (ACD) was also observed for CoFe, ZnNi,
ZnFe and CuPb. Since the composition of the alloy depends strongly on the pH
value [72] the electrolyte has to contain a buffer system. For different concentrations
of iron salts (Figure 8.4) alloys (crystallite size: 16–19 nm) with iron contents up to
71 mol% have been obtained. Hessami et al. [73] explain the ACD by an increased
dissociation rate of the FeOH+ complex compared to that of the NiOH+ complex.
For this reason the concentration of free iron ions in the electrolyte increases and
therefore the alloys exhibit an increased iron content. There is the risk (or chance)
of preparing gradient materials in which the iron concentration increases with
deposition depth whereas laterally the deposited alloy is perfectly homogeneous.
Varying the crystallite size by varying the pulse parameters or the temperature is
impossible because these factors also change the alloy composition (Figure 8.5).
The best way to vary the crystallite size without changing the alloy composition is
by the addition of different amounts of grain refiners.
220 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
Fig. 8.4 The influence of the Fe3+ concentration of the electrolyte on
the alloy composition [29].
8.3
Special Features of Ionic Liquids as Electrolytes
The electrochemistry of ionic liquids is different in some essential features from the
electrochemistry of aqueous electrolytes. Particularly for electrodeposition, which
involves charge transfer from the electrolyte to the electrode, the double layer on
the electrode is of great importance. In general the cathode is negatively charged
for the electrodeposition of metals and therefore coated with a (Helmholtz-) layer of
cations at least 0.5 nm thick; but the metal species in most ionic liquids is anionic
(for instance AlCl4 − ). This makes the metal deposition process complicated, for
more details we refer to Chapter 2.
Fig. 8.5 The influence of the temperature (left axis) and the pulse
current density (right axis) on the alloy composition of FeNi
alloys [29].
8.3 Special Features of Ionic Liquids as Electrolytes 221
In ionic liquids the coordination chemistry and concentration of metal complexes
are also substantially different from those in aqueous electrolytes, with consequent
effects on both the thermodynamics, i.e., the redox potential, and the kinetics of
the deposition process. For details we again refer to Chapter 2.
Since the viscosity of ionic liquids is large in some cases and concomitantly the
diffusion is slow, ionic liquids generally exhibit a lower conductivity than aqueous
electrolytes. To improve the mass transport it has been suggested to add diluents
like benzene, toluene or acetonitrile. Water may also be a suitable diluent in some
cases, acting as both a ligand and a viscosity improver.
Brighteners are common additives in the electroplating industry and are mostly
based on empirical recipes. A well-known inorganic example is arsenic acid. Here
only the influence of organic compounds on the electrodeposition from ionic liquids
is discussed. The main effect of brighteners is adsorption on the electrode surface
thus impeding nucleation and influencing growth.
To determine the effectiveness of the brighteners Natter et al. [74] have examined
organic molecules with different structures (see Table 8.1). The first group comprises aromatic and aliphatic carboxylic acids with one or two carboxylic groups.
The second group consists of aromatic carboxylic acids with chlorine substituents
in different positions (2-, 3- and 4-chlorobenzoic acid). The third group comprises
carboxylic acids with one or two hydroxy substituents and the last group are substances with a sulfur-containing functional group (benzoic acid sulfimide, sodium
butanesulfonate, sodium dodecyl sulfate). Aliphatic carboxylic acids and also their
hydroxy substituted derivatives (tartaric, malonic, malic and salicylic acids) show
hardly any grain refining effect, maybe because these substances do not have free
electron pairs which are important for adsorption at many metals. The aromatic
salicylic acid reduces the crystallite size down to 54 nm. For this reason the aromatic
Table 8.1 Effect of different additives on the crystallite size (electrolyte:
63 mol% absolute dry AlCl3 , 37 mol% [EMIM]Cl, DC: 5 mA cm−2 , additive concentration: 4 wt.%).
Additive
benzoic acid
3-chlorobenzoic acid
2-chlorobenzoic acid
4-chlorobenzoic acid
benzoic acid sulfimide
phthalic acid anhydride
sodium dodecyl sulfate
sodium butane sulfonate
tartaric acid
salicylic acid
malonic acid
malic acid
D/nm
19
13
18
24
12
16
21
132
99
54
61
133
222 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
carboxylic acids seem to be the better grain refiners. The electronic interaction of
the aromatic ring can be enhanced with halogen substituents. The use of benzoic
acid chloroderivatives shows a strong crystallite refining effect which depends on
the position of the chlorine group. 3-Chlorobenzoic acid interacts strongly with
the metal surface and works as a good grain refiner. Saccharine (benzoic acid sulfimide) has additional nitrogen and sulfur atoms which increase the surface activity
and decrease the crystallite size. The same effect can be observed for long chain
sulfonates and phthalic acid anhydride.
8.4
Nanocrystalline Metals and Alloys from Chlorometallate-based Ionic Liquids
Chlorometallate-based ionic liquids are eutectic mixtures of an organic chloride,
RCl, and a metal chloride MClx , mostly with M = Al or Zn; other possible but
less commonly used metals are Sn, Ga, Fe, or Ge, i.e., chlorostannate, chlorogallate, etc. Chloroborate (BCl4 − ) or bromoaluminate (AlBr4 − ) systems are also
possible. In chloroaluminate-based ionic liquids the AlIII exists in complexes like
AlCl4 − , Al2 Cl7 − or Al3 Cl10 − , in chlorozincate-based ionic liquids the ZnII exists in
complexes like ZnCl2 − , Zn2 Cl5 − or Zn3 Cl7 − , depending on the concentration.
The mole fraction of AlCl3 can be denoted as r = [AlCl3 ]/([AlCl3 ] + [RCl]).
According to this ratio the chloroaluminate systems are categorized as follows:
r > 0.5 called Lewis-acidic, r = 0.5 called Lewis-neutral and r < 0.5 called Lewisbasic melts. These systems have the serious drawback of being extremely sensitive
to humidity/water which causes hydrolysis and the formation of HCl: for that
reason extreme care has to be applied (controlled inert gas atmosphere with a water
content below 1 ppm). However, under these demanding conditions, which might
be too difficult in industrial applications, the chloroaluminate-based ionic liquids
have attractive electrochemical properties. They allow the electrodeposition not only
of Al but also of a large number of other metals: salts or oxides of other metals can
be dissolved in chloroaluminate systems, and then the result is dependent on the
ratio r mentioned above (and of course on the potential): from Lewis-acidic systems,
with r > 0.5, Al alloys are deposited, and from Lewis-basic systems, with r<0.5, the
other metal can be deposited in a pure form. This flexibility of chlororaluminate
systems occurs because the potential determining electrochemical reaction changes
with the composition, therefore the molar ratio r is an important parameter for the
experimentalist.
Nanostructured aluminum [74–78], iron [74] and aluminum–manganese alloys
[74] have been prepared from a Lewis acid AlCl3 /[BMIM]Cl mixture (65 mol% AlCl3 ,
35 mol% [BMIM]Cl) whereas palladium alloys have been deposited from a Lewis
basic system (45 mol% AlCl3 , 55 mol% [BMIM]Cl). The electrochemical cell and all
parts which are in contact with the electrolyte have to be built from inert materials.
As cathode material glassy carbon can be used. A constant ion concentration in
the electrolyte can be realized by the use of a sacrificial anode consisting of the
8.4 Nanocrystalline Metals and Alloys from Chlorometallate-based Ionic Liquids 223
Table 8.2 The influence of different process parameters on the nano-
structure of aluminum deposits. In the case of additive addition benzoic acid was used.
Sample
2
3
4
5
6
7
8
Iaverage /mA cm−2
Ipeak /mA cm−2
ton /ms
toff /ms
additive/g L−1
D/nm
6.67
10
54
11.9
21.2
60.1
4.2
60.1
—
—
—
310
550
1570
—
1570
∞
∞
∞
2
2
2
∞
2
0
0
0
50
50
50
0
50
0
0
0
0
0
0
2.5
2.5
121
98
84
140
129
122
20
7
same material to be deposited on the cathode. In this way the electrolysis can run
for several days without a break. Grain refining additives like carboxylic acids have
been used for crystallite size reduction because these organic substances exhibit
very good solubility in the ionic liquid. The temperature of the bath should be kept
constant at 40 ◦ C to enhance the conductivity of the electrolyte. Temperatures above
60 ◦ C lead to crystallite growth during the deposition and have to be avoided.
The crystallite sizes of aluminum samples prepared with different process parameters [74] are summarized in Table 8.2. The electrolyte consisted of 5.0 g [EMIM]Cl
and 8.8 g absolutely dry AlCl3 . In the case of additive addition carboxylic acids, especially benzoic acid or nicotinic acid, were used. According to Eq. (8.1) the crystallite
size should decrease with increasing overpotential. This behavior can be observed
for samples 1–3. The use of 54 mA cm−2 decreases the crystallite size to 84 nm.
A further increase in the current causes decomposition of the electrolyte. For this
reason a pulsed current with high peak currents for a short time (ton ) and proper
breaks (toff ) can be used. Samples 4–6 were prepared with different peak currents at
constant pulse times. It can be observed that with increasing peak current density
the crystallite size decreases. This behavior can be explained by Ostwald ripening,
which sets in during the toff -time, and can be impeded by the use of organic additives which interact with the active growth sites by adsorption [48, 79] and thus
act as grain refiners. The freshly deposited adatoms are forced to form new nuclei,
resulting in the formation of a nanostructure. Sample 7 confirms this behavior. The
addition of just a small amount (2.5 g L−1 ) of benzoic acid reduces the crystallite
size to 20 nm. The additional use of pulsed current causes a further reduction in
the crystallite size.
From Table 8.2 it can be seen that sample 3 has a smaller crystallite size than
sample 6 although sample 3 was prepared with a lower average current density.
We consider the average current densities because sample 3 was prepared by DC
plating whereas sample 6 was prepared by PED plating. The larger crystallites in the
case of PED plating can also be explained by Ostwald ripening during the toff -time
224 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
Fig. 8.6 The influence of the benzoic acid concentration on the crystallite size of aluminum deposits [74].
whereas in the case of DC deposition a continuous deposition and growth of nuclei
takes place.
As mentioned above the organic additives block the active growth sites. In this
case the crystallite size should be a function of the additive concentration [80,
81]. Therefore the aluminum deposition from an AlCl3 /[BMIM]Cl bath has been
repeated with increasing benzoic acid amounts (see Figure 8.6). Only a small
concentration of this additive reduces the crystallite size to 40 nm, and a further
addition does not lead to any substantial further reduction in the crystallite size.
The limit is at about 1.5 wt.%. The addition of more benzoic acid shows no further
effect because all active sites are blocked by the additive molecules. Even an extreme
surplus of additive does not change the nanostructure [64].
Additive adsorption is evident from the results of cyclic voltammetry. Figure 8.7
shows the cyclic voltammogram of a Lewis acid electrolyte consisting of
[BMIM]Cl/AlCl3 electrolyte (66.7 mol% AlCl3 ) with and without additives. The
electrochemical window is about 2.3 V limited by the cathodic bulk deposition of
aluminum (peak 3) and the decomposition of the electrolyte at 1.9 V (vs. Al/Al3+ ).
Two additional peaks at −0.17 and +0.4 V (1 and 2) can be observed resulting
from underpotential deposition [82, 83]. During the inverse run the formation
of gold alloys (Al2 Au5 , AlAu2 ) [84] takes place at +0.33 and +0.37 V (Figure 8.7,
inset).
With increasing amounts of benzoic acid (0 to 0.38 wt.%) we observe a decreasing
peak current for the aluminum deposition (see Figure 8.7), the aluminum oxidation
peak disappears and the underpotential deposition (UPD) of aluminum is also
strongly diminished. These experiments lead to the conclusion that the additive
molecules block the active growth sites and therefore the peak current of aluminum
deposition increases with decreasing additive concentration. A further consequence
8.4 Nanocrystalline Metals and Alloys from Chlorometallate-based Ionic Liquids 225
Fig. 8.7 Cyclic voltammograms for aluminum deposition with increasing amounts of additive (benzoic acid) [74].
of this behavior is the irreversibility of the process (no aluminum oxidation) and
the strongly reduced UPD process.
The effect of temperature on the nanostructure of the deposits in the presence
of additives has been examined using an electrolyte with 3.5 wt.% benzoic acid.
The experimental details of the deposition are given in Figure 8.8. The crystallite
size increases from 23 to 72 nm between 40 and 63◦ C. This indicates a strong
Fig. 8.8 Temperature dependence of the crystallite size for Al samples
prepared from additive containing electrolyte [74].
226 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
Table 8.3 Saturation magnetization, relative remanence and coercivity
for different crystallite sizes of nanostructured iron.
D/nm
40
57
62
72
79
82
114
159
Ms /kA m−1
Relative remanence MR /MS
HC / kA m−1
1235
1372
1230
1235
1377
1179
1283
1656
0.010
0.121
0.094
0.089
0.060
0.074
0.067
0.015
151.2
135.3
111.4
103.5
79.6
87.5
45.3
31.9
temperature dependence of the interaction between metal surface and additive;
furthermore, Ostwald ripening is enhanced with increasing temperature since it is
based on surface diffusion of adatoms and additives which is strongly temperature
dependent.
Nanostructured iron was deposited from an electrolyte consisting of 5.0 g
[BMIM]Cl, 8.8 g absolute dry AlCl3 and 0.5 g anhydrous FeCl3 . Benzoic acid (40
g L−1 ) was used as grain refiner. The crystallite size of the deposits (Table 8.3) can
be adjusted by the variation of the DC-current density (0.05–10 mA cm−2 ). Magnetization measurements reveal a saturation magnetization essentially independent
of D, an increase of the coercivity field Hc proportional to D−1 , and a remanence
with a maximum where the grain size equals the magnetic exchange length.
Alloys like Alx Mn1–x or Alx In1–x are very difficult to prepare in nanostructured form because the substantial difference in physical properties of both metals (melting point, hardness) does not allow the use of ball milling methods,
chemical processes or inert gas condensation procedures. The preparation of
these alloys can, however, be realized with the electrodeposition technique, using an AlCl3 /[BMIM]Cl (53:47 wt.%) mixture with addition of the corresponding
metal salts (5.5 wt.% InCl3 or MnCl2 ). The alloys were deposited with Iaverage =
10 mA cm−2 at 50 ◦ C. In both cases nanostructured alloys with a crystallite size of
25 nm were obtained (Figure 8.9). The alloy composition can be controlled by the
composition of the bath [85–87]. A surplus of aluminum in the ionic liquid also
increases the aluminum content in the alloy. By the use of ionic liquids with AlCl3
contents between 20 and 39 mol% different Alx Mn1–x alloys can be prepared
In contrast to these electrodeposition experiments where nanocrystalline material in macroscopic quantities results there are a number of reports in the literature where, in an electrochemical scanning tunneling microscope (EC-SCM)
experiment, individual nanoclusters/nanocrystals were deposited and immediately
imaged. Aravinda and Freyland [88] have studied the electrocrystallization of Sb
and AlSb on Au single crystals from a chloroaluminate-based ionic liquid; the ionic
liquid was neither Lewis acidic nor Lewis basic but neutral. AlIII mainly exists as
AlCl4 − and is not reducible in the potential range of the Sb deposition; therefore, Sb
nuclei can be electrodeposited and coalesce to form deposit domains and eventually
grow according to the Stransky–Krastanov mode. On the other hand, at low potential
8.5 Nanocrystalline Metals from Air- and Water-stable Ionic Liquids 227
Fig. 8.9 X-ray diffraction pattern of an AlMn alloy prepared from a
Lewis acid AlCl3 /[BMIM]Cl electrolyte [74].
three-dimensional clusters of Alx Sby appear. With this EC-STM study the viability
of electrochemical methods to deposit less noble compound semiconductors like
AlSb on the nanometer scale is claimed.
A detailed review and discussion of EC-STM studies can be found in Chapter 9.
8.5
Nanocrystalline Metals from Air- and Water-stable Ionic Liquids
Nanocrystalline aluminum has also been electrodeposited, without any additives,
from the ionic liquid 1-butyl-1-methyl-pyrrolidinium bis(trifluoromethylsulfonyl)
imide saturated with AlCl3 [89]. This ionic liquid has, compared to chloroaluminate
ionic liquids, a number of advantages: it is water and air stable, easy to purify, and
easy to dry to water contents below 1 ppm. AlCl3 dissolves well and homogeneously
in it up to a concentration of about 1.5 mol L−1 giving a clear solution, from which Al
cannot be deposited. Upon further increase in the concentration of AlCl3 a biphasic
mixture is obtained, similar to the behavior
of several ionic liquid systems based
on the bis(trifluoromethylsulfonyl)imide (T f 2 N), as described by Wasserscheid
[90]. The lower phase is colorless while the upper one is pale and more viscous,
Figure 8.10(a). The biphasic mixture AlCl3 /IL becomes monophasic by heating to a
temperature of 80 ◦ C, Figure 8.10(b). Obviously, reducible aluminium-containing
species only exist in the upper phase of the AlCl3 /IL mixture, and only from that
phase does electrodeposition of Al occur.
Figure 8.11 shows the cyclic voltammogram of the upper phase of the biphasic
mixture of AlCl3 /1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)imide
on a gold substrate at room temperature. At a potential of −0.7 V (vs. Al), the
cathodic current rises with two small cathodic steps at −0.9 and −1.3 V which
are correlated to two different redox processes before the bulk growth of Al sets
228 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
Fig. 8.10 (a) A biphasic mixture of the ionic liquid 1-butyl-1-methyl
pyrrolidinium bis(trifluoromethylsulfonyl)amide containing 1.6 M AlCl3
at room temperature. (b) The biphasic mixture becomes monophasic
at 80 ◦ C [89].
in. A clear nucleation loop is observed in the forward scan which means that
the bulk deposition of aluminum in this complicated system seems to require a
certain activation energy/overpotential. The onset of bulk Al deposition occurs at
a potential of about −1.5 V. In the reverse scan, the cathodic current continues
to flow, forming a current loop which is typical for nucleation processes. A small
Fig. 8.11 Cyclic voltammogram recorded at the Au(111)
substrate in the ionic liquid 1-butyl-1-methyl pyrrolidinium
bis(trifluoromethylsulfonyl)amide containing 1.6 M AlCl3 (from the
upper phase of the mixture) at room temperature. The scan rate was
10 mV s−1 [89].
8.5 Nanocrystalline Metals from Air- and Water-stable Ionic Liquids 229
anodic peak is recorded on the reverse scan at a potential of about −0.17 V which is
attributed to the incomplete stripping of the electrodeposited aluminum. Stripping
seems to be kinetically hindered, which is a common phenomenon in ionic liquids.
The Al electrodeposits were investigated by means of a high resolution fieldemission scanning electron microscope (SEM) and by energy dispersive X-ray
analysis (EDAX), to explore the morphology and composition, respectively. Visually, the deposit appears thick and shiny. Figure 8.12(a) shows a high resolution
micrograph of a 10 µm thick layer of Al on a gold substrate, electrodeposited at
room temperature at a constant potential of −1.7 V for 2 h. The layer contains
very fine crystallites about 20 nm in size. It is also seen in the micrograph that
the deposited layer is a bit stressed. The quality of the deposit which is obtained
at a constant potential of −0.45 V for 2 h at 100 ◦ C is significantly better and the
crystallites also become finer, Figure 8.12(b). The EDAX profile taken for the area
shown in the micrograph of Figure 8.12(a) shows a strong Al peak and a weak Au
peak (originating from the gold substrate), Figure 8.12(c). This confirms that the
deposited Al layer is thick and dense.
Figure 8.13 shows the XRD patterns of a nanocrystalline Al film obtained at
a constant potential of −1.7 V for 2 h at 100 ◦ C in the ionic liquid [BMP]Tf2 N
containing 1.6 M AlCl3 on a glassy carbon substrate. The X-ray diffractogram shows
the characteristic pattern of crystalline Al, with broad peaks indicating the small
crystallite size of the electrodeposited Al. The inset of Figure 8.13 exhibits the
evaluation of the full-width half maximum (FWHM) of the Al(111) Bragg reflection,
as an example, of the electrodeposited Al film. The grain size of Al was determined,
using Scherrer’s equation [56], to be 34 nm.
The organic cation in these ionic fluids has an amphiphilic character: 1-butyl1-methyl pyrrolidinium, e.g., has a hydrophilic head (charged nitrogen atom) and
a hydrophobic butyl tail. Thus it interacts with the nuclei in a similar way to the
additives discussed in Section 8.3, i.e., it affects the surface energy of the growing
metal nuclei, thus the relative energetics of the nucleation and growth processes,
and thus the morphology of the deposit. Zein El Abedin et al. [91] actually found that
the organic cations of Tf2 N− salts strongly affect the morphology of the deposit.
The 1-butyl-1-methyl-pyrrolidinium cation and the trihexyl-tetradecylphosphonium cation have a pronounced amphiphilic character, and the aluminum
films electrodeposited from both these Tf2 N− salts exhibit very fine crystallites
with average sizes of about 20 nm (Figure 8.14), whereas the 1-ethyl-3-methylimidazolium cation evidently is not suffiently amphiphilic and the corresponding
Al deposits contain coarse cubic-shaped microcrystallites (Figure 8.15). Evidently,
this behavior of the first mentioned cations is due to adsorption phenomena that
prevent Ostwald ripening and growth.
Nanocrystalline copper with an average crystallite size of about 50 nm can
be obtained without additives in the ionic liquid 1-butyl-1-methyl-pyrrolidinium
bis(trifluoromethylsulfonyl)imide ([BMP]Tf2 N) [92]. Because of the limited
solubility of the tested copper compounds in this ionic liquid, copper cations
were introduced into the ionic liquid by anodic dissolution of a sacrificial copper
electrode. The electrodeposition of copper was also investigated in the ionic liquid
230 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
Fig. 8.12 SEM micrographs of electrodeposited Al on gold formed
after potentiostatic polarization for 2 h in the upper phase of the mixture AlCl3 /[BMP]Tf2 N: (a) at room temperature, E = −1.7 V; (b) at
100 ◦ C, E = −0.45 V. (c) EDAX profile for the area shown in the SEM
micrograph (a) [89].
8.5 Nanocrystalline Metals from Air- and Water-stable Ionic Liquids 231
Fig. 8.13 XRD patterns of an electrodeposited Al layer obtained
potentiostatically at −1.7 V for 2 h in the upper phase of the mixture AlCl3 / [BMP]Tf2 N at 100 ◦ C on a glassy carbon substrate. Inset:
FWHM of Al(111) peak of XRD patterns [89].
1-butyl-1-methylpyrrolidinium trifluoromethanesulfonate [BMP]TFO, with
Cu(TFO)2 as a commercial source of copper [93]. The employed electrolyte was
prepared by adding appropriate amounts of Cu(TFO)2 to the ionic liquid [BMP]
TFO until the saturation limit.
Figure 8.16 shows the cyclic voltammogram of the ionic liquid [BMP]TFO saturated with Cu(TFO)2 at a gold substrate. In the forward scan, two reduction steps are
recorded, c1 and c2. The first might be correlated to an underpotential deposition
process, whereas the second one is clearly due to the bulk deposition of copper.
The corresponding anodic peaks, a1 and a2, are recorded in the anodic branch of
the cyclic voltammogram. It was concluded that the deposition of copper in the
employed electrolyte is a quasi-reversible diffusion-controlled process. The SEM
micrograph in Figure 8.17(a) shows the surface morphology of an electrodeposited
copper layer on gold obtained at a constant potential in the ionic liquid [BMP]TFO
saturated with Cu(TFO)2 . As seen, the deposit is dense and contains fine crystallites
with average sizes of about 40 nm. The deposit was analysed as metallic copper, as
revealed by the corresponding EDX profile, Figure 8.17(b).
Ge(111) bilayers can be obtained by electrodeposition in the dry ionic liquid
[BMIM]PF6 containing GeI4 as a source of germanium [94]. This ionic liquid has
an electrochemical window of a little more than 4 V on Au(111). However, stable
232 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
Fig. 8.14 SEM micrographs of electrodeposited Al films on gold formed in the upper phase of the mixture AlCl3 /[BMP]Tf2 N
after potentiostatic polarization for 2 h at
(a) room temperature (E = −1.7 V, corresponding I = −0.5 mA cm−2 ); (b) 50 ◦ C (E
= −1.0 V, corresponding I = −1 mA cm−2 );
(c) 75 ◦ C (E = −0.75 V, corresponding I
= −1.7 mA cm−2 ) and (d) 100 ◦ C (E =
−0.45 V, corresponding I = −2 mA cm−2 )
[91].
nanoclusters/nanocrystals or thick layers could not be obtained from this system.
Recently, thick layers of germanium have been electrodeposited from the ionic liquid [BMIM]PF6 , saturated either with GeBr4 or GeCl4 [95, 96]. A thin germanium
layer with a rather metallic behavior and a maximum thickness of 300 pm forms
before bulk growth of germanium sets in. Bulk deposition starts with nanoclusters,
and nanosized micrometer-thick layers can easily be obtained. Moreover, Endres
et al. [97] have shown that germanium nanoclusters with a narrow height distribution can be made in a dilute solution of GeCl4 in the ionic liquid [BMIM]PF6 . The
lateral sizes of most of the clusters are in the range 20–30 nm, while their heights
vary from 1 to 10 nm, with most of them being between 1 and 5 nm.
Abbott et al. [98–103] reported the synthesis and characterization of new moisturestable, Lewis acidic ionic liquids made from metal chlorides and commercially available quaternary ammonium salts (see Chapter 2.3). They showed that mixtures of
choline chloride (2-hydroxyethyltrimethylammonium chloride, [Me3 NC2 H4 OH]Cl
and MCl2 (M=Zn, Sn) give conducting and viscous liquids at or around room temperature. These deep eutectic solvents/ionic liquids are easy to prepare, are waterand air-stable, and their low cost enables their use in large-scale applications. Furthermore, they reported [104] that a dark green, viscous liquid can be formed by
mixing choline chloride with chromium(III) chloride hexahydrate and that the
8.5 Nanocrystalline Metals from Air- and Water-stable Ionic Liquids 233
Fig. 8.15 SEM micrographs of electrodeposited Al films on gold formed after potentiostatic polarization for 1 h in the upper
phase of the mixture AlCl3 /[EMIM]Tf2 N at
(a) 25 ◦ C (E = −0.3 V, corresponding I =
−1 mA cm−2 ); (b) 50 ◦ C (E = −0.135 V,
corresponding I = −2.5 mA cm−2 ); (c)
75 ◦ C (E = −0.082 V, corresponding I
= −3.5 mA cm−2 ) and (d) 100 ◦ C (E =
−0.055 V, corresponding I = −4 mA cm−2 )
[91].
Fig. 8.16 Cyclic voltammogram of the ionic liquid [BMP]TFO saturated with Cu(TFO)2 on gold at room temperature. Scan rate
10 mV s−1 [93].
234 8 Nanostructured Metals and Alloys Deposited from Ionic Liquids
Fig. 8.17 (a) SEM micrograph of nanocrystalline copper obtained
potentiostatically on Au in the ionic liquid [BMP]TFO saturated with
Cu(TFO)2 at a constant potential for 2 h at room temperature. (b)
EDAX profile of the area shown in the SEM micrograph [93].
physical properties are characteristic of an ionic liquid. The eutectic composition
is 1:2 choline chloride/chromium chloride. Chromium can be electrodeposited
efficiently from this ionic liquid to yield a crack-free deposit [104]. Adding small
ions like Li+ to an ionic liquid decreases the Helmholtz layer thickness considerably
and should make ion reduction easier. This enhances nucleation, as has been
shown qualitatively in the chromium case, i.e., the deposition of chromium from a
eutectic mixture of chromium chloride and choline chloride. Up to 10 mol% LiCl
led to a change in deposit morphology from microcrystalline to nanocrystalline and
a change in visual appearance from metallic to black [105].
8.6
Conclusion and Outlook
Nano is a word that is no longer reserved for science but has entered the public
consciousness. The field of nanoscience with its promise of amazing nanotechnologies is one of today’s most challenging, exciting, multidisciplinary and competitive
References
fields. A prerequisite for all achievements and promises is a synthetic access to
nanomaterials. Electrodeposition of nanocrystalline metals and alloys has opened
the route to surface coatings with innovative functionalities and to bulk materials
with improved physical properties. The productive power and the possibilities of
electrodeposition have been expanded significantly by the recent introduction of
ionic liquids as electrolytes, and a variety of nanostructured metals, metal alloys
and even semiconductors, deposited from ionic liquids, will contribute to the 21st
century demand for new materials in the fields of nanotechnology, information
technology and even biomedicinal engineering. The transfer of the knowledge, accumulated so far and to be accumulated, from academic research into industrial
applications is the great challenge for the near future.
Acknowledgment
We gratefully acknowledge financial support by the Deutsche Forschungsgemeinschaft in the framework of Sonderforschungsbereich 277. For stimulating discussions and mutually fruitful collaboration in the past years we thank Dr. Bukowsky
and Professor Endres, and we are grateful for Professor Endres’ continuous encouragement while writing this chapter.
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239
9
Electrodeposition on the Nanometer Scale: In Situ
Scanning Tunneling Microscopy
Frank Endres, and Sherif Zein El Abedin
9.1
Introduction
Nanocrystalline materials have received extensive attention since they show unique
mechanical, electronic and chemical properties. As the particle size approaches
the nanoscale, the number of atoms in the grain boundaries increases, leading
to dramatic effects on the physical properties and on the catalytic activity of the
bulk material. Nowadays, there is a wide variety of methods for the preparation of
nanocrystalline metals such as thermal spraying, sputter deposition, vapor deposition and electrodeposition. The electrodeposition process is commercially attractive
since it can be performed at room temperature and the experimental set-up is less
demanding. Furthermore, the particle size can be adjusted over a wide range by
controlling the experimental parameters such as overvoltage, current density, composition, and temperature (see Chapter 8).
Nowadays ionic liquids can be regarded as potential electrolytes for the electrodeposition of nanomaterials. Therefore, the electrochemical processes and the
factors that influence the deposition and the stability of the structures have to be
investigated on the nanometer scale. Over the last decade, scanning tunneling microscopy (STM) has been widely used as a powerful tool for probing surfaces on
the nanometer scale. A combination of classical electrochemical methods with in
situ STM has facilitated the investigation of the electrodeposition process on the
nanometer scale and even on the atomic scale. This gives valuable information on
the phase formation and growth at the electrode/electrolyte interface. At the end
of the last decade, we started for the first time to carry out in situ STM studies on
electrochemical phase formation in ionic liquids. There was no knowledge of the
local processes of phase formation in ionic liquids, however, due to wide electrochemical windows, these systems give access to elements that cannot be obtained
in aqueous solutions, such as Al, Si, Ta and many more.
Apart from our in situ STM studies in ionic liquids, there are a few papers dealing with this subject, especially in chloroaluminate ionic liquids, although air- and
water-stable ionic liquids are now commercially available (see for example Refs.
[1–3]). As chloroaluminate ionic liquids are extremely hygroscopic, in situ STM
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
240 9 Electrodeposition on the Nanometer Scale: In Si tu Scanning Tunneling Microscopy
Fig. 9.1 Photos of a home-built STM head for in situ studies in ionic liquids.
measurements have to be performed under inert gas conditions, otherwise they
decompose, liberating HCl and different oxo-chloro-aluminates in the liquid which
makes reproducible fundamental experiments impossible. Evolution of HCl with
its corrosive action can damage the STM heads and formation of the oxochloroaluminates may lead to a contamination layer at the electrode/ionic liquid interface.
Air- and water-stable ionic liquids enable one to perform high quality in situ STM
studies. Figure 9.1 shows two photos of one of our STM heads for in situ studies.
The whole of the head is handled inside an inert gas glove box and put into a
vacuum-tight stainless steel vessel under inert gas with H2 O and O2 contents of
less than 1 ppm. This allows experiments to run for up to one week without the
risk of altering the ionic liquid quality. An important point is that we have a parallel
approach of the sample plate to the tip which allows complete remote control of
approach and experiment. For details on STM theory and principles we would like
to refer the reader to Ref. [4].
In this chapter we present a few selected results on the nanoscale electrodeposition of some important metals and semiconductors, namely, Al, Ta and Si,
in air- and water-stable ionic liquids. Here we focus on the investigation of the
electrode/electrolyte interface during electrodeposition with the in situ scanning
tunneling microscope and we would like to draw attention to the fascinating
9.2 In situ STM in [Py1,4 ] TFSA
electrode processes on the submicron/nanometer scale. The first step in investigation of a new system is always to investigate the surface of Au(111) in the liquid of
interest at variable electrode potentials. This is the only way to distinguish between
surface processes which are due to the liquid alone (restructuring/reconstruction,
impurities) and processes due to metal or semiconductor electrodeposition. Au(111)
on mica is our standard substrate as it is easily commercially available and delivers
high quality experiments. The purity of ionic liquids is a challenge for fundamental
physicochemical studies, especially with the in situ STM. It is tough to purify ionic
liquids as, hitherto, they can neither be distilled without decomposition nor recrystallized nor sublimed. It will be briefly discussed how even apparently ultrapure
ionic liquids can contain low amounts of inorganic impurities leading to unexpected behavior on the single crystalline surface of Au(111). Such impurities might
also affect the deposition of metals, semiconductors and conducting polymers, in
the initial stages. Due to their importance we focus in this chapter solely on the
third generation of ionic liquids, i.e. air- and water-stable ones. The third generation ionic liquids have received extensive attention, not only because of their low
reactivity with water but also because of their large electrochemical windows of up
to 6 V. Usually these ionic liquids can be well dried to water contents below 1 ppm
under vacuum at temperatures between 100 and 150 ◦ C. In the next section the electrochemical window of 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)
amide is discussed. The pyrrolidinium cation is one of the most stable against
cathodic breakdown and therefore it is interesting for electrochemistry. However,
there is also an unexpected anion breakdown and an influence of the cation on
the size of the deposited crystals. The local probe electrodeposition of Al, Ta and
poly-p-phenylene is briefly summarized.
9.2
In situ STM in [Py1,4 ] TFSA
Figure 9.2 shows a typical cyclic voltammogram of ultrapure 1-butyl-1methylpyrrolidinium bis(trifluoromethylsulfonyl)amide ([Py1,4 ] TFSA) on Au(111)
[5, 6]. Ultrapurity means that the supplier (here: Merck KGaA/EMD) guarantees
that water and halide impurities are below the 10 ppm level. Routinely the liquids
are dried to water contents below 3 ppm prior to use in our laboratory.
As shown in Figure 9.2 the electrochemical window of this liquid on Au(111)
can be determined by the extrapolation of the rising cathodic and the rising anodic
currents to zero. This is not a thermodynamically exact value but gives a good value
for the real thermodynamic window on the respective substrate. The cathodic limit
is mainly due to the irreversible reduction of the [Py1,4 ] to N-methylpyrrolidine and
butyl radicals which undergo further decomposition to butene(s) and hydrogen (see
also Chapter 10). The oxidation wave A4 is directly correlated with the breakdown of
the cation C4. The anodic limit is due to gold disintegration and partly to irreversible
anion oxidation. The potential regime in between is the maximum available electrochemical window and can be determined to be 5.6 V on Au(111). But what about
241
242 9 Electrodeposition on the Nanometer Scale: In Si tu Scanning Tunneling Microscopy
Fig. 9.2 CV of dry pure [Py1,4 ]TFSA on Au(111) at a scan rate of 10 mV s−1 .
the peaks/waves C1–C3 and A3? Macroscopically there is no surface modification
visible. In the potential regime between +2 V and −3 V vs. ferrocene/ferrocinium
(Fc/Fc+ ) gold looks like gold should look, and thats all. The quartz crystal microbalance does not show any mass effect in this potential window. One could argue
about soluble organic or inorganic impurities, but the liquids are ultrapure with
respect to the inorganics and, furthermore, purified by chromatography to remove
organic impurities, thus this cannot be the reason. Figure 9.3 shows the surface of
Au(111) under [Py1,4 ] TFSA at the open circuit potential, i.e. at around −0.4 V vs.
Fc/Fc+ .
It is quite interesting that gold does not show here the typical surface known
in aqueous electrochemistry with flat terraces separated by steps, it is, in contrast,
strongly structured with a wormlike pattern. Such patterns have also been described
by the Mao group in other ionic liquids [7]. Interactions of the gold surface with
ions of the ionic liquid lead to such restructuring phenomena. If the electrode
potential is reduced successively to −1.7 V the wormlike pattern disappears slowly.
There is a potential regime where only vacancy islands are observed (from −0.4
to −1.7 V), finally a flat terraced gold surface is obtained, as evidenced in Figure
9.4(a). In this picture the transformation is not yet complete (some vacancy islands
are still present), but on a timescale of 20–30 min these vacancy islands disappear
completely, i.e. in the potential regime of wave C1. Thus the peak C1 could be
correlated to the restructuring of the gold surface. However, the respective oxidation
peak is missing although MacFarlane et al. [8] have described that in this potential
regime the irreversible breakdown of the TFSA ion starts.
9.2 In situ STM in [Py1,4 ] TFSA
Fig. 9.3 STM image in pure [Py1,4 ]TFSA on an Au(111) surface. Instead of a typical flat Au(111) surface, a rough surface with worm-like
structures is observed at −0.4 V (OCP) .
If we perform the same experiment with ultrapure 1-ethyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)amide ([EMIM] TFSA), we do not see the same restructuring with the in situ STM, although MacFarlane describes that it is the anion
which is subject to irreversible breakdown in this potential regime. As described by
MacFarlane et al. [8] the reduction of TSFA weakens one of the N–S bonds leading
to its cleavage:
2−
−
−
−
•
N(SO2 CF3 )−
2 + e = N(SO2 CF3 )2 = NSO2 CF3 + SO2 CF3
(9.1)
The reduction products of TFSA can undergo further reduction reactions as
follows:
•
−
−
NSO2 CF−
3 + e = NSO2 + CF3
SO2 CF−
3
−
+e =
SO−
2
+
CF−
3
(9.2)
(9.3)
This led us to the conclusion that adsorption of the [Py1,4 ] cation (maybe together
with TFSA + TFSA breakdown products) is responsible for the wormlike pattern
and the formation of a flat surface finally [9, 10]. Between −1.7 and −1.9 V vs. Fc/Fc+
a flat gold surface can be probed and around −2 V (i.e. in the potential regime of
wave C2) we observed routinely that the picture quality got worse, see Figure 9.4(b).
In the first experiments we thought this would be due to a bad tip, a common
problem the experimentalist has to struggle with, but it was surprising that the noise
shown in Figure 9.4(b) disappeared again when the electrode potential was set back
to −1.7 V and reappeared at about −2 V. Such a reversible and reproducible behavior
excludes a “bad tip”. If in the in situ STM experiment the electrode potential is set to
values between −2.2 and −2.7 V it is evident that the picture quality is dramatically
reduced, as shown in Figure 9.4(c). It should be mentioned clearly that this is
243
244 9 Electrodeposition on the Nanometer Scale: In Si tu Scanning Tunneling Microscopy
Fig. 9.4 (a) When the electrode potential is
reduced to more negative values (−1.7 V)
the typical Au(111) terrace-like surface is
observed.(b) At −2.0 V the surface appears
reproducibly “noise”, which disappears
again when the electrode potential is set
to less negative values and thus this effect
does not correlate to a “bad” tip.(c) When
the electrode potential is set between −2.2
and −2.7 V, the intensity of the “noise” increases.(d) If the electrode potential is reduced to −2.9 V, a thin film is formed on
the gold surface, which makes the identification of the gold terraces very difficult.
definitely not due to a bad tip, and there is still a tunnelling contact between tip
and surface allowing one to probe the surface. At −2.9 V (in the potential regime
of wave C3), Figure 9.4(d), the surface is now obviously covered by a film which
makes probing of the surface difficult. Nevertheless, the steps can still be identified.
At lower electrode potentials, i.e. in the regime of the cathodic breakdown C4, the
tunnelling contact is finally lost. Together with the results from the MacFarlane
group it can be concluded that the TFSA is subject to a certain cathodic breakdown
and with the in situ STM this breakdown can be probed.
It should be mentioned, furthermore, that MacFarlane et al. reported recently on
the breakdown of films of the TFSA on magnesium and magnesium alloy surfaces
9.3 Electrodeposition of Aluminum 245
[8]. It can be concluded from our in situ STM experiments presented here that the
understanding of even apparently simple electrochemical windows of ionic liquids
can be a tough job, and we ourselves were not completely right with our first interpretation of the waves C1–C3 [6]. It is not sufficient just to have a look at the
cyclic voltammograms, it may be required first to acquire fundamental local probe
information of any ionic liquid which is to be employed for fundamental studies at
the interface electrode/ionic liquid. Furthermore, there are combined anion/cation
effects known in ionic liquids affecting chemical and electrochemical processes, so
it might not be sufficient just to exchange the anion for a more stable one. The surface behavior might be different. Fundamental local probe electrochemistry may
first require a detailed characterization of the potential-dependent interface effects.
It should be mentioned that inorganic impurities in ionic liquids (even in apparently ultrapure liquids) may lead to a complete misunderstanding of the surface
processes. We have discussed in Ref. [5] that liquids, made by a metathesis reaction
from [Py1,4 ]Cl and Li-TFSA, can contain low amounts of Li ions. Apparently perfect
on glassy carbon there is clear evidence for the underpotential deposition of Li on
Au(111). A further error in IL synthesis can originate from purification processes.
In order to remove the often yellowish color of ionic liquids after synthesis they are
commonly purified over silica or alumina. One of the dominant impurities, even in
high quality silica, is aluminum species, which can be washed off by the ionic liquid
and finally reduced to Al in the cathodic regime. As a consequence, in our laboratories any newly delivered ionic liquid is first tested by cyclic voltammetry and in situ
STM on Au(111) thoroughly before it is employed for fundamental studies. This
approach is somewhat time consuming, on the other hand it is currently the only
chance of avoiding misinterpretation of electrochemical experiments, especially
with the in situ STM. This is one of the challenges in ionic liquids electrochemistry
and makes in situ STM experiments extremely time consuming.
9.3
Electrodeposition of Aluminum
Aluminum can be well electrodeposited in first generation ionic liquids and there
are many papers available in the peer-reviewed literature. A main shortcoming of
these liquids is that the organic halides (e.g. [EMIM]Cl or [BMIM]Cl) are extremely
difficult to dry. Consequently these organic halides can easily contain more than
1000 ppm of water. A typical “synthesis” quality from the catalogue might contain
10 000 ppm of water. This water will consequently react with AlCl3 when the liquid
is made from the educts. A major problem is the evolution of HCl, which can
severely damage the STM heads, apart from impurity effects at the interface with
the electrode due to the oxochloroaluminates which produce a slushy film on the
electrode surface. In in situ STM experiments such a contamination layer at the
interface electrode/ionic liquid makes high quality experiments impossible. In
order to avoid such problems our aim was to investigate Al deposition from third
generation ionic liquids, as these liquids can easily be dried to water contents below
246 9 Electrodeposition on the Nanometer Scale: In Si tu Scanning Tunneling Microscopy
Fig. 9.5 Biphasic behavior of the ionic liquid [Py1,4 ] TFSA/ 2 M AlCl3 at 25 and 80 ◦ C.
3 ppm (the Karl–Fischer limit). Our original motivation was to try to deposit Al from
liquids that contain low amounts of Al(III) species. For this purpose we selected
again [Py1,4 ] TFSA. On the one hand this liquid has a low decomposition potential
for the cation, on the other hand the charge in the bulky TFSA ion is delocalized,
thus maybe avoiding strong complexation with AlCl3 . Unfortunately the latter was
not the case: AlCl3 can indeed be dissolved easily in concentrations between 0.01
and 1 mol l−1 giving a clear solution, but in no case can Al be deposited. Obviously
AlCl3 and the TFSA form a not yet fully characterized complex which cannot be
reduced to Al. Beginning with an AlCl3 concentration of 1.6 mol l−1 the solution
becomes biphasic at room temperature. Above 2.5 mol l−1 the mixture solidifies. At
80 ◦ C the mixture becomes monophasic again and at 100 ◦ C up to 5 mol l−1 AlCl3
can be dissolved in [Py1,4 ]TFSA. From the lower phase of this biphasic ionic liquid
Al cannot be deposited. Figure 9.5 shows the phase behavior of [Py1,4 ] TFSA with
2 mol l−1 AlCl3 at 25 and 80 ◦ C. Quite interestingly, electrodeposition in the upper
phase delivers, at temperatures between 25 and 125 ◦ C,clearly a nanocrystalline
aluminum with grain sizes of about 20–30 nm [10]. Although the upper phase is
rather emulsion-like in nature around room temperature with high viscosity it is
possible, as shown below, to perform in situ STM experiments with an acceptable
resolution in such a viscous liquid.
The cyclic voltammogram of this upper phase around room temperature on
Au(111) is represented in Figure 9.6. There are several reduction peaks A–E and a
very small oxidation peak E. The reduction peak E is correlated with the bulk electrodeposition of nanocrystalline aluminum, the peak E is due to some Al oxidation.
However, under the mentioned conditions the deposition is practically irreversible.
The gold surface is macroscopically not subject to any visible modification between
the peaks A and D. If the gold is polarized to −1.2 V for some time there is, visually,
no modification at the surface. Even after removing the liquid, the gold surface
still looks like gold. If the experiment is performed with [EMIM]TFSA and AlCl3 a
9.3 Electrodeposition of Aluminum 247
Fig. 9.6 Cyclic voltammogram for aluminum deposition on Au (111)
in the upper phase of the biphasic mixture of AlCl3 /[Py1,4 ] TFSA at
room temperature. Scan rate: 10 mV s−1 .
biphasic mixture is also obtained and the cyclic voltammogram of that upper phase
gives a fully reversible aluminum deposition with a clear underpotential deposition
of Al, see Ref. [11]. With the in situ STM the Al deposition starts with the growth
of islands in the underpotential deposition regime then, after deposition of nearly
3 monolayers, the bulk growth sets in (for details see Ref.[11]). Figure 9.7 shows a
sequence of STM images recorded on Au(111) in the upper phase of the biphasic
mixture of 5.5 M AlCl3 /[EMIM] Tf2 N. At a potential of 0.5 V, 2D Al islands are
formed, as shown in Figure 9.7(a). If the potential is decreased further to 0.4 V the
number of aluminum islands increase until – with the exception of a few vacancies –
one aluminum monolayer is formed (Figure 9.7(b)). These vacancies are still
present even after further reducing the potential to 0.3 and 0.2 V, at which points a
second and a third aluminum layer, respectively, start to grow. When the electrode
potential is further reduced to 0.1 V, for several minutes, the number of deposited
crystallites increases rapidly and a 3D growth sets in, Figure 9.7(c). Reducing the
electrode potential to −0.1 V leads to a crystal growth and the initial aluminum
deposit exhibits a granular structure, as shown in Figure 9.7(d). This picture shows
that the crystallites are homogeneously spread with an average diameter of about
15 nm. Nevertheless, when a bulk deposit is made in this liquid the obtained deposit
is clearly microcrystalline [10] showing that the results from in situ STM cannot
necessarily be transferred to bulk processes.
In the light of these results, it can be concluded that the Al deposition in
[EMIM]TFSA behaves on the nanoscale more or less as in the first generation
ionic liquids published earlier [12, 13], but it is much easier to keep the quality of
the STM experiments as the liquids are per se water free. With [Py1,4 ] TFSA the
behavior is more complicated. Figure 9.8 shows a set of STM images on Au(111) in
the upper phase of the [Py1,4 ]TFSA/AlCl3 mixture at room temperature at the open
248 9 Electrodeposition on the Nanometer Scale: In Si tu Scanning Tunneling Microscopy
Fig. 9.7 In situ STM images of Au (111) in the upper phase of the
biphasic mixture of AlCl3 /[EMIM] TFSA at (a) E = 0.5 V, (b) E = 0.4 V,
(c) E = 0.1 Vand (d) E = −0.1 V vs. Al/Al(III).
circuit potential (0.3 V). Quite similar to the pure liquid, the gold surface shows a
worm-like pattern but the quality of the STM pictures under practically the same
conditions is not as good as in the pure liquid, Figure 9.8(a). This might be due to
the remarkably higher viscosity of this liquid with AlCl3 dissolved in it or due to
the influence of halide ions. If the electrode potential is decreased to values as low
as −0.9 V the worm-like pattern disappears completely and the more or less typical
gold surface is obtained (Figure 9.8 (b)).
At slightly lower electrode potentials (−1.1 V) tiny two-dimensional islands start
growing (Figure 9.8(c)) and at −2 V the bulk phase of aluminum starts growing
giving, interestingly, crystals with sizes of 100–300 nm (see Ref. [11]), quite in contrast to the nanocrystals obtained by potentiostatic deposition in the overpotential
deposition regime. As already described above, this is a shortcoming of the STM
experiment which is performed at low overvoltages where there is a slow growth.
9.3 Electrodeposition of Aluminum 249
Fig. 9.8 In situ STM images of Au (111) in the upper phase of the
biphasic mixture of AlCl3 /[Py1,4 ] TFSA at (a) E = 0.3 V, (b) E = −0.9 V
and (c) E = −1.1 V vs. Al/Al(III).
250 9 Electrodeposition on the Nanometer Scale: In Si tu Scanning Tunneling Microscopy
The bulk deposits are, in contrast, made at higher overvoltages where the initial
nuclei are smaller than at lower overvoltages [14]. As there is no clear surface
process the peaks A–D are likely to be correlated with solution species and some
aluminum deposition. It is quite fascinating that the cation of an ionic liquid seems
to have such a dramatic influence on the electrodeposition of metals. In [EMIM]
TFSA/AlCl3 the electrochemistry is reversible, with underpotential and overpotential deposition giving a microcrystalline deposit. In [Py1,4 ] TFSA/AlCl3 there is no
clear underpotential deposition, the electrochemistry is practically irreversible and
the final deposit is nanocrystalline with two orders of magnitude smaller crystal
sizes. The reasons for this surprisingly different behavior are not yet understood
and will require, besides further in situ STM experiments with varying liquids,
some solution chemistry and simulation studies.
9.4
Electrodeposition of Tantalum
As was shown in Chapter 4, elemental tantalum can be electrodeposited in the
water- and air-stable ionic liquid [Py1,4 ] TFSA at 200 ◦ C using TaF5 as a source of
tantalum [15, 16]. The quality of the deposit was found to be improved upon addition
of LiF to the deposition bath. At room temperature only ultrathin tantalum layers
can be deposited as the element. The electrodeposition of tantalum was investigated
by in situ STM to gain insight into the electrodeposition process.
Figure 9.9 shows the cyclic voltammogram of [Py1,4 ] TFSA containing 0.5 M TaF5
on Au(111) at room temperature. As shown, two reduction processes are recorded
in the forward scan. The first starts at a potential of −0.5 V with a peak at −0.75 V,
this might be correlated to the electrolytic reduction of Ta(V) to Ta(III). The second
Fig. 9.9 Cyclic voltammogram of 0.5 M TaF5 in [Py1,4 ]TFSA on
Au(111) at 200 ◦ C. Scan rate: 10 mV s−1 .
9.4 Electrodeposition of Tantalum 251
Fig. 9.10 In situ STM images and I–U STS of an about 300 nm thick
layer of Ta deposited on Au(111) in [Py1,4 ]TFSA containing 0.5 M TaF5
at −1.25 V vs. Pt.
252 9 Electrodeposition on the Nanometer Scale: In Si tu Scanning Tunneling Microscopy
process starts at a potential of −1.5 V and is accompanied by the formation of a black
deposit on the electrode surface. This can be attributed to the reduction of Ta(III)
to Ta metal simultaneously with the formation of insoluble tantalum compounds
on the electrode surface. The anodic peak recorded on the backward scan is due to
the partial dissolution of the electrodeposit which, however, does not seem to be
complete. Then the anodic current increases as a result of gold dissolution at E >
1.5 V. In situ STM measurements under potentiostatic conditions were performed
on Au(111) in the ionic liquid [Py1,4 ] TFSA containing 0.5 M TaF5 . The STM picture
of Figure 9.10(a) shows the surface morphology of a deposited tantalum layer
obtained at a potential of −1.25 V (vs. Pt). As seen, the deposited layer is rough
and some triangularly shaped islands with heights of several nanometers, Figure
9.10(b), grow above the deposited layer. With time, these islands grow vertically
and laterally and finally merge together to a thick layer, Figure 9.10(c) and (d).
The thickness of the deposited layer obtained from the change of z-position of
the piezo, was found to be about 300 nm. In order to investigate whether the in
situ deposit is metallic or not, current–voltage (I–U) tunneling spectroscopy was
conducted. A typical in situ tunneling spectrum of the 300 nm thick layer of the
electrodeposit at different positions is shown in Figure 9.10(e). As seen, the I–U
spectrum exhibits metallic behavior with an exponential-like rise in the current,
indicating the formation of elemental Ta. There are approaches in the literature to
determine the apparent electron work function from such I–U spectra. As several
simplifications are required and as the influence of the ionic liquid on the distancedependent tunneling spectra are completely unknown we rather restrict ourselves
to a qualitative description.
In the light of the above results, it can be concluded that ionic liquids – due to their
wide electrochemical windows – not only give access to many elements like e.g. Al
and Ta and many others, they also show unexpected cation/anion effects, both at the
interface electrode/ionic liquid and on the electrodeposition of metals. Due to the
extremely high number of ionic liquids there is an enormous potential to perform
original electrochemical studies on both the microscale and the nanoscale. But,
there is a price to pay: as it is tough to purify ionic liquids – hitherto they can neither
be recrystallized nor distilled nor sublimed without remarkable decomposition –
only ultrapure liquids can be recommended for fundamental studies, and even these
liquids can contain very low amounts of inorganic impurities. The experimentalist
has to find first which impurities are present in a liquid of interest, then if they cause
any problems at all and if there is a certain level of impurities which maybe can be
tolerated. Thus, there is no quick (and dirty) electrochemistry in ionic liquids, and
the comparatively slow progress, especially with the in situ STM, is maybe one of
the greatest challenges for the experimentalist.
9.5
Electrodeposition of Poly(p-phenylene)
Conducting polymers have been extensively investigated due to their potential applications in supercapacitors, sensors, batteries, electrochromic devices and light
9.5 Electrodeposition of Poly(p-phenylene) 253
emitting diodes. Polypyrrole and polythiophene are the most studied conducting polymers due to their high stability and simple preparation [17–19]. Among
all conducting polymers, poly(p-phenylene) (PPP) is very interesting because it
is suitable for the fabrication of blue polymer light emitting diodes (PLED) [20–
22]. However, the electrochemical polymerization of benzene to PPP is still a
challenge as water in the solution has to be strictly avoided. Therefore, in the
past only solvents like concentrated sulfuric acid [23], liquid SO2 [24] or liquid
HF were feasible for the electropolymerization of benzene. In 1993, ionic liquids
based on AlCl3 were employed for the first time for the electropolymerization of
benzene [25]. Because of side reactions due to chlorine co-evolution during the
electropolymerization the quality of the deposits was not satisfactory. Recently,
we have reported for the first time that modern air- and water-stable ionic liquids are also well suited for the electropolymerization of benzene [26, 27]. In
contrast to the above-mentioned solvents, ionic liquids deliver much milder chemical conditions. The electropolymerization of benzene in the ionic liquid 1-hexyl3-methylimidazolium tris(pentafluoroethyl)trifluorophosphate [HMIM]FAP is well
reproducible and gives as a result poly(p-phenylene) of spherulitic morphology with
grain sizes as small as 500 nm [26]. Figure 9.11(a) shows three successive cyclic
voltammograms for the oxidation of benzene in the ionic liquid [HMIM]FAP. In
the anodic scan of the first cycle (curve 1) an anodic current starts at an electrode
potential of 1.8 V (vs. Pt-quasi reference). The current rises strongly, indicating
the oxidation of benzene to form a polymer. It is worth mentioning that the rise
of this current occurs at an electrode potential that is about 1 V below the anodic
decomposition limit of the liquid, see Ref. [26]. In the back scan a reduction process
at a potential of 0.75 V is recorded. In the second anodic cycle (curve 2) an oxidation
peak at an electrode potential of 1.25 V is obtained, and in the second cathodic scan
with reference to the first cycle a more pronounced current flows. In the third cycle
(curve 3) both anodic and cathodic currents increase. Such cyclic voltammetric behavior is typical for the electrosynthesis of conducting polymers. Visual inspection
during the deposition of the polymer reveals that a yellowish film has formed on
the platinum electrode after the second cycle. With subsequent cycles the color
changes from yellowish to brownish, and finally a black polymer film is obtained
on the electrode surface. Quite a similar observation was reported by Kowalski
et al. during the electrosynthesis of poly(p-phenylene) by polymerization of benzene in a mixture of glacial acetic acid and concentrated sulfuric acid [28]. Here,
a thick and well adhering polymer film was also obtained by applying, for a sufficiently long time, an electrode potential of 2.2 V vs. the quasi-reference to the
platinum working electrode. The surface morphology of an electrodeposited polymer film on the platinum electrode is shown in Figure 9.11(a). As seen, the film
consists of small spherical and globular grains with an average diameter of about
3 µm. The smallest grains that can be observed with the selected resolution of the
SEM have sizes of around 500 nm. IR spectra of the deposited polymer film reveal
the formation of poly(p-phenylene) as a result of the polymerization of benzene in
the employed ionic liquid [26].
In order to gain further insight into the growth and characterization of
the deposited polymer film we acquired in situ STM and STS measurements.
254 9 Electrodeposition on the Nanometer Scale: In Si tu Scanning Tunneling Microscopy
Fig. 9.11 (a) Cyclic voltammogram of 0.2 M benzene in [HMIM]FAP
on platinum. The numbers refer to the respective cycle. Scan rate:
10 mV s−1 . (b) SEM micrograph of the electropolymerized film on the
platinum electrode after synthesis in [HMIM]FAP.
Figure 9.12 shows a set of in situ STM images obtained on Au(111) in the ionic
liquid [HMIM]FAP containing 0.2 M benzene, as well as an I–U spectrum of a
deposited PPP layer. As shown in the STM image of Figure 9.12(a), obtained at
a potential of 0.9 V, the gold surface is subject to slight oxidation at such anodic
potential and the gold terraces are still distinct. At 1.3 V, a number of randomly distributed 2D-islands are formed on the gold surface, as manifested in the STM image
of Figure 9.12(b). This indicated the start of the formation of PPP. By setting the
potential at 1.9 V, a relatively thick polymer layer of PPP is obtained, Figure 9.12(c).
In order to measure the band gap of the deposited polymer layer, current–voltage
tunneling spectroscopy was performed. It has already been shown by us that I–U
9.5 Electrodeposition of Poly(p-phenylene) 255
Fig. 9.12 In situ STM images of Au(111) in the ionic liquid
[HMIM]FAP containing 0.2 M benzene at (a) 0.9 V, (b) 1.3 V and (c)
1.9 V vs. Pt-quasi ref.(d) In situ I–U spectrum of a PPP layer obtained
at a potential of 1.9 V (vs. Pt).
tunneling spectroscopy is a valuable technique for in situ characterization of electrodeposited semiconductors [29–31] and metals [15]. We could show with in situ
I–U tunneling spectroscopy that germanium with a layer thickness of 20 nm and
more is semiconducting with a symmetric band gap of 0.7 ± 0.1 eV. On the other
hand, we have found that very thin layers of germanium with thicknesses of several
monolayers clearly exhibit metallic behavior [30, 31]. Figure 9.12(d) represents an
in situ current–voltage tunneling spectrum of the deposited PPP layer shown in
Figure 9.12(c). As seen in the spectrum, a band gap of 2.1 ± 0.2 eV is recorded.
This value approaches the value reported in the literature for PPP, 2.7 eV, [32]. A
detailed in situ STM and STS study on the film formation stages during the electropolymerization of benzene in the ionic liquid [HMIM]FAP is now in preparation
and will be published elsewhere [33].
256 9 Electrodeposition on the Nanometer Scale: In Si tu Scanning Tunneling Microscopy
9.6
Summary
Ionic liquids represent a promising class of solvents with unprecedented properties
for electrochemistry. They can have wide electrochemical windows of 6 V, they
have – in most cases - practically no vapor pressure around room temperature,
they have wide thermal windows of 300 ◦ C and they show unusual cation/anion
interaction which can influence chemical and electrochemical reactions. Ultrapure
ionic liquids allow in situ STM experiments at deeply negative electrode potentials
with high quality and, consequently, give insight into the initial electrodeposition
of reactive elements like Si, Al, Ta and presumably many others. Some ionic liquids
like those with the bis(trifluoromethylsulfonyl)amide anion are, however, subject to
an unexpected anion breakdown which can alter the nanoscale processes. Varying
the cation or the anion of an ionic liquid might have a dramatic influence on the
surface electrochemistry, as shown with the example of Al deposition. This opens
the door to many fundamental studies, not only with classical electrochemistry but
also with in situ STM. As shown in the example of benzene polymerization, the
growth and in situ characterization of conducting polymers can be probed on the
nanoscale with in situ STM.
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259
10
Plasma Electrochemistry with Ionic Liquids
Jürgen Janek, Marcus Rohnke, Manuel Pölleth, and Sebastian A. Meiss
10.1
Introduction
Electrochemical reactions occur at the interface between two phases with sufficiently different conduction behavior, i.e. a predominantly ion-conducting electrolyte phase and an electrode phase with predominantly electronic conduction.
Among all possible types of interfaces the most intensively applied are solid
metal|liquid electrolyte and solid metal|solid electrolyte. Electrode systems which
have been much less studied are those formed by combining either a solid or liquid
conducting phase with a low-temperature gas discharge (plasma).
This chapter aims to discuss and summarize theoretical and practical aspects of
such plasma interfaces, presenting the existing examples from our own recent work
on plasma electrochemical reactions between typical ionic liquids and plasmas.
First, we address the plasma state and essential properties with respect to its
application in electrochemistry. Today, low temperature plasmas – mostly in the
form of radiofrequency or microwave plasmas – play an important role in the
treatment or modification of solid surfaces. However, as plasma chemistry is usually
not an element of chemistry curricula, we include a very brief introduction but refer
the reader to the literature for more detailed information.
Plasma electrochemical reactions have been studied by chemists for a surprisingly long time, with the first report on cathodic metal deposition at the free surface
of a liquid electrolyte with free electrons from a plasma dating back to 1887 [1],
long before the plasma state had been named by Langmuir in 1928 [2]. A short
survey of past work with more conventional liquid electrolytes is also included in
this chapter.
Typical low-temperature plasmas are usually only weakly ionized and quasineutral but are thermally in a non-equilibrium state, i.e. the different plasma
species (molecules, atoms, ions and electrons) possess different kinetic energy
distributions. Because of their small mass electrons acquire much more kinetic
energy than atomic or molecular species and thus show an energy distribution
which corresponds to a much higher temperature than in the case of much heavier
particles. The stability of ionic liquids towards reduction by these “hot” electrons
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
260 10 Plasma Electrochemistry with Ionic Liquids
and also towards reactions with other reactive plasma species is considered in
Section 10.4.
The central part of this chapter comprises a summary of our recent attempts to
deposit metals from ionic liquids by plasma cathodic reduction. The concluding
part then analyzes the plasma electrochemical approach with respect to possible
applications.
10.2
Concepts and Principles
Like other salt melts ionic liquids are characterized by a specific combination
of physicochemical properties: high ionic conductivity, low viscosity, high thermal
stability compared to conventional liquid solvents, wide electrochemical windows of
up to 7 V and – in most cases – extremely low vapor pressures. Due to their low vapor
pressure ionic liquids are not only well suited for the application of UHV-based
analytical techniques (e.g. photoelectron spectroscopy [3]), but also for use in plasma
reactors with typical pressures of the order of 1 Pa up to 10 kPa. Moreover, due to
their high electrical conductivity, ionic liquids may even be used as “electrodes” for
plasmas. To date there are just a few reports on the combination of low-temperature
plasmas and ionic liquids available in the literature [4–6]. Therefore, the essential
aspects of experiments with ionic liquids in typical plasma reactors are discussed
in this section.
10.2.1
Plasma Electrochemistry
As plasma chemistry deals with charged particles, there is no doubt that it can be
considered as plasma electrochemistry a priori. However, a comparison of electrochemistry and plasma chemistry [4–9] in more detail is instructive. Electrochemistry
deals with the interplay of electric fields (potential differences) and chemical reactions. Once electric fields get strong enough, e.g. at interfaces, electron transfer can
be enforced leading to reduction or oxidation of chemical compounds. Two routes
exist within this interplay: (i) external electric potential differences can be used to
control chemical processes; (ii) chemical processes can be used to generate electric
potential differences. The first (synthetic/charging) route is the basis of electrochemical synthesis and galvanic technology. The second (analytical/discharging) route is
the basis for batteries, fuel cell and sensor technology.
Obviously, plasmas can be used very efficiently within the synthetic approach (i),
and all examples given in this paper are assigned to the synthetic approach. It is
much less obvious whether plasmas can be used also in the counter-direction.
In order to measure a stable and reproducible electromotive force (EMF) the
corresponding electrochemical (galvanic) cell must be in (local) thermodynamic
equilibrium. Low-temperature plasmas represent non-equilibrium states and are
highly inhomogeneous systems from a thermodynamic point of view, often not
10.2 Concepts and Principles 261
fulfilling the conditions of local equilibrium (at least at interfaces). Therefore EMF
measurements in these plasmas will not provide easily accessible thermodynamic
information. Flames can, in small volumes, be considered as equilibrium systems
under certain conditions, and almost a century ago the first attempts to measure
EMF signals in flame plasmas were reported [10–12]. The theoretical analysis of the
EMF measurements in flames is complicated by complex electrode processes and,
to date, no adequate treatment has been published. First steps towards a correct and
comprehensive description have been published by Caruana et al. [13, 14]. Several
other researchers have tried to measure EMF signals in flames but got no stable
signal (e.g. Lorenz et al. [15]).
10.2.2
Low-temperature Plasmas: Electrodes or Electrolytes?
From the electrochemical point of view, the answer to this question appears, at
first glance, to be simple: electrolytes are usually defined as electrically conductive
media with a negligible electronic conductivity, for example as purely ionic conductors. In contrast, electrode materials have to be predominantly electronic conductors
(mostly metals). This definition originates from electrochemistry in the liquid state,
where an electronic contribution to the bulk charge transport in electrolytes is a
rare phenomenon (except in some well-known cases, e.g. sodium/ammonia solutions). In solid state electrochemistry we deal mostly with mixed ionic/electronic
conductors, as electronic charge carriers are a priori always present in solids, either
as intrinsic defects (electron–hole pairs due to a sufficiently small band gap) or as a
consequence of non-stoichiometry (metal excess: n-type doping, non-metal excess:
p-type doping). The mixed character of conduction in the solid state is the basis
for chemical diffusion of crystal components and, among other effects, is responsible for the occurrence of diffusion potentials. It becomes obvious in the Nernst
equation for the EMF of solid state galvanic cells, which contains the electronic
transference number as a factor.
A plasma is always a mixed ionic and electronic conductor, mostly with a small
ionic transference number. The majority of charge carriers in plasmas of electropositive gases (e.g. noble gases) are cations and electrons. The charge carriers
in electronegative gases (e.g. halogens) are cations, anions and electrons. In addition to their two orders of magnitude higher mean energy, the mobility of free
electrons is two orders of magnitude higher than the mobility of free ions and,
thus, the electronic partial conductivity is a priori much larger than the ionic partial
conductivity.
Following these arguments, plasmas should rather be regarded as electrodes
than as electrolytes. However, this simple analysis completely neglects the inhomogeneity of most plasmas, in particular in the boundary regions in front of walls.
Low-temperature plasmas are non-equilibrium systems which only exist in stationary states driven by a continuous conversion of electric energy into heat. Diffusion
of the charged plasma species from the plasma bulk to the plasma boundaries
establishes extended space charges with large diffusion potentials and leads to
262 10 Plasma Electrochemistry with Ionic Liquids
a negative charging of the plasma walls. The resulting (electron-poor) positive
space charge region acts as a rectifying element and leads to strongly non-linear
current–voltage characteristics. As a consequence, the total resistance of a plasma
is rather controlled by the mobility of ions in the plasma sheaths than by the mobility of electrons (see below). Depending on the applied potential to an electrode
in a plasma, the carrier concentration in the plasma sheath changes considerably.
Under negative polarization the plasma sheath is particularly poor in electrons and
therefore will act as an electrolyte rather than an electrode.
In the present context, we suggest considering low-temperature plasmas as fluid
mixed conductors with a small ionic transference number in principle – in particular for positive (anodic) polarization of the electrode in the ionic liquid and
corresponding negative (cathodic) polarization of the electrode in the plasma. As
exemplified below, plasmas are more often used as gaseous electrodes than as
electrolytes. One of the most important application of electrolytes, i.e. their use as
“electron filters” in galvanic cells, is hampered by the large electronic contribution
to the bulk conductivity and by the relatively large diffusion potential within the
plasma sheaths. There is also a practical aspect that further complicates quantitative
potential measurements in plasmas: the local plasma density and the corresponding charge carrier density depend on the boundary conditions set by the reactor
walls and the electrode arrangement. Often the wall and electrode surfaces are
slowly covered with thin sputtered films. Once these are electrically conducting the
local plasma state may change considerably with time.
10.2.3
The Plasma|Electrolyte Interface
One of the characteristic features of plasmas is their inhomogeneity at boundaries.
The faster electrons charge any wall of a plasma reactor negative and leave a positive
space charge in front of the wall, the so-called “sheath” [16]. This space charge can
formally be treated like a diffusion potential in conventional electrolytes. Depending
on the degree of ionization, the Debye length can take large values (up to several
millimeters) for dilute plasmas with large potential drops within this region. The
experimental study of these inhomogeneous plasma boundaries is hampered by
the fact that most methods interfere with the electric fields within the plasma.
Therefore, only a few methods can be used for the investigation. However, for a
qualitative discussion we can restrict ourselves to the general picture of plasma|solid
boundaries with a positive space charge in the plasma in front of the solid, extending
to the order of the Debye length. The situation at the plasma|wall interface is
depicted schematically in Figure 10.1. The negatively charged wall does not affect
neutral particles, while positive particles are accelerated towards it and negative
particles are repelled from the wall so that only highly energetic negative particles
can reach it. In a stationary state the fluxes of highly energetic negative and positive
particles towards the wall compensate each other.
The consequence of the plasma sheaths is two-fold: first, significant voltages
have to be applied to plasma electrochemical cells in order to draw sufficiently
10.2 Concepts and Principles 263
Fig. 10.1 Positive space charge layer at the
interface between a plasma and (a) a dielectric, (b) a metallic, (c) an electrolytic wall
with floating potential W . Due to the negative surface charge mainly neutral, positive
and only high energetic negative plasma
particles reach the wall ( je− : flow of electrons, jK+ : flow of cations, jK : flow of neutralized particles, jA− : flow of anions). The
potential difference between the zero poten-
tial 0 and the potential of the plasma p
is denoted as plasma potential p . If the
wall’s potential is not “free floating” (without an applied external potential) as shown
above, the characteristics of the potential in
the wall, the surface charge at the interface
and hence the characteristics of the sheath
and pre-sheath potential can be influenced
by the applied external potential.
264 10 Plasma Electrochemistry with Ionic Liquids
large currents for electrochemical reactions, see below. Only by a careful design
of the electrodes can we overcome the electric current limitations of the space
charge regions. Secondly, the large electric fields within the plasma sheaths lead to
a significant acceleration of ionic charge carriers and the resulting sputter effects.
10.2.4
Types of Plasmas and Reactors
The properties of plasmas vary strongly with gas composition, pressure and the
method and parameters of the plasma generation process. The charge carrier concentration depends on the pressure and the fractional ionization of the plasma, for
instance basically on the power density. The mobility of the electrons depends on
the electron temperature, which is typically several orders of magnitudes greater
than the gas temperature or the temperature of the ionized species in non-thermal
low temperature plasmas used for electrochemical purposes.
Both parameters, as well as the gas composition, can usually be changed relatively
easy within the limits of the given experimental set-up. When employing the plasma
merely as a gaseous but chemically inert electrode, one will choose a noble gas,
typically argon. For other purposes, reactive gases might be added to the noble gas
or replace it. Depending on the reaction at the plasma|electrolyte interface, gaseous
reaction products may emerge into the plasma, or components may disappear
from it due to reactions with the electrolyte or substances dissolved in it, therefore
changing its composition. As the pressure of a typical non-thermal low-temperature
plasma is two to four orders of magnitude smaller than atmospheric pressure, even
small absolute amounts of emerging or disappearing gas components can have a
relatively high effect on the gas composition of the plasma. Therefore, it is desirable
to design plasma electrochemical experiments in such a way that allows a high gas
throughput, especially near the interface, to retain a well defined gas composition.
The method of the plasma generation has a strong influence on the parameters
of the plasma, including the distribution of the various ionized species of the gas
components. Feasible plasma reactors are direct current (DC) discharge reactors
and inductively or capacitively coupled radio frequency (RF) discharge reactors, the
latter not being discussed here further. Microwave (µW) discharges are spatially
much more concentrated, due to the much smaller wavelength, and lead to a
considerable increase in temperature. They are often applied in the form of “remote
plasma sources”, i.e. using a gas flux with particles excited during their passage
through the plasma source rather than working in the center of the plasma source.
Figure 10.2(a) shows the set-up for a DC discharge, where the electrolyte (ionic
liquid) is part of the serial electric circuit. This set-up has the advantage that
it is possible to gain information about the formation rate of a product at the
plasma/electrolyte interface directly by measuring the electric current and having
information about the relevant transference numbers. However, it is not possible
to freely choose the applied voltage. It has a lower limit, given by the voltage of the
order of 100 –1000 V needed to sustain the discharge. Figure 10.2(b) shows a DC
discharge set-up where the electrolyte is not necessarily part of the electric circuit.
10.3 Early Studies 265
Fig. 10.2 Different types of plasma reactors employing the use of an IL: (a) DC discharge with the IL as an integral part of a
serial set-up, (b) DC discharge with the IL
as optional part of a parallel set-up, (c) in-
ductively coupled RF discharge with an electric circuit for electrochemical experiments
independent from the plasma generating
process.
It, rather, represents an ion-conducting wall of the plasma at a floating potential
and reactions are motivated by the plasma–wall interactions described earlier. It is
feasible to introduce a third electrode to the system, placing it in contact with the
electrolyte, but not with the plasma, and therefore gaining some control over the
potential difference between the electrolyte and the plasma. In the case of purely
ion-conducting electrodes, the electric current offers information about the reaction
rate at the plasma/electrolyte interfaces.
In an inductively coupled RF discharge (Figure 10.2(c)) the plasma is not in
contact with the external RF coil (“electrode-free discharge”). Again the ionic liquid
acts as a “wall” to the plasma, with the effects described earlier. Its floating potential
will be negative, due to the collected electrons, and a positive space charge is found
above the surface. Introducing an electrode to the electrolyte allows one to influence
its then no longer “floating” potential. A second electrode can be placed in the gas
phase, but often metallic parts of the reactor itself are used as the second electrode.
This set-up has been applied successfully in experiments with solid electrolytes and
typical I–U curves are reported by Vennekamp [17].
10.3
Early Studies
The use of gas discharges for electrochemical processes has been investigated for
more than 100 years, and a full account is beyond the scope of this chapter. We
will focus on a few innovative and seminal studies which can be regarded as major
advances. The first plasma electrochemical experiments were already reported in
1887 by Gubkin [1], in the same year when Arrhenius published his most influential paper on electrolytic dissociation of salts in water [18]. Gubkin investigated
266 10 Plasma Electrochemistry with Ionic Liquids
Fig. 10.3 Set-up of the reproduced Gubkin experiment: silver is dissolved at the anode inside of the liquid electrolyte and reduced at the
plasma|electrolyte interface; and photograph of the laboratory experiment.
the plasma-assisted cathodic deposition of silver, platinum and zinc oxide. For this
purpose an aqueous metal salt solution was put into a round flask fitted with two
platinum electrodes. The anode was coated with a layer of the same metal that was
dissolved in the form of its salt in the electrolyte. It was immersed in the liquid
electrolyte and a vacuum was generated above the electrolyte by cooling the bulb
with the boiling electrolyte after sealing the bulb. A glow discharge over the surface
of the liquid electrolyte was produced by applying a high voltage between both electrodes. Gubkin observed the deposition of clearly visible metal particles, formed
by reduction of the metal cations with free electrons from the plasma at the interface between the plasma and the liquid. The plasma electrochemical cell can be
summarised as:
metal Me (anode) | salt solution (Mez+ ) | plasma | inert metal Pt (cathode)
Figure 10.3 shows Gubkin’s original experiment, as it was reproduced in our
laboratory. A sketch of the experimental set-up and a photograph of the experiment
are depicted. It has to be mentioned that Gubkin was not the first person, who
reduced metal ions or metal compounds by using plasmas. Trasatti [19] reports
on experiments performed by Father Beccaria as early as 1750, who seemingly
observed the reduction of zinc oxide to zinc metal by an electric discharge.
In the 1920s the phenomenon of electrostenolysis was investigated by Söllner [20].
When a voltage higher than required for electrostenolysis (U > 20 V) was applied
to the cell
anode | transition metal salt solution || membrane ||
heavy metal salt solution | cathode
the deposition of metal was observed inside the membrane. In addition a light
emission was observed at higher voltages, which indicates the occurrence of micro gas discharges. At still higher voltages spark discharges were observed. This
10.3 Early Studies 267
phenomenon is explained by the electrolytic formation of gas in pores and cracks
of the glass membrane. Due to the high electric resistance of the gas bubbles, the
main electrical potential decay is assumed at the gas bubbles inside the pores. In
consequence, the electric fields at these pores are very high and micro-plasmas are
generated inside the bubbles. At the gas|electrolyte interface metal deposition takes
place. Similar phenomena of plasma formation in gas bubbles were observed in
electrolytic commutators and capacitors [21].
The study of micro-, spark or arc discharges in liquid electrolytes (usually referred
to as plasma electrolysis) has been continued by other groups [22], depositing either
metals or metal oxides. Here the metal or metal oxide is deposited cathodically or
anodically, respectively in the presence of a gas discharge in front of the electrode.
Shen et al. reported that the resulting metal or metal oxide layers are comparatively
dense and show better corrosion protection than conventionally deposited coatings
[23]. They propose this plasma-assisted deposition as a method with high potential
for industrial application in corrosion protection. The processes within micro arc
discharges in liquid electrolytes are complex and not yet fully understood. As the
properties of these discharges themselves cannot be controlled directly by well
adjustable experimental parameters, we exclude them at this point from further
consideration. However, ionic liquids may provide new opportunities for the further
development of spark electrolysis.
Gubkin’s simple plasma electrochemical experiment was reproduced and improved in the 1950s and 60s, mainly by Klemenc and Brenner [24–29]. A typical
experimental set-up of Klemenc is depicted in Figure 10.4. The process was named
glow discharge electrolysis or electrode-less electrolysis (which is a misnomer a priori,
as electrolysis always requires electrodes) and an attempt was made to explain
the phenomena occurring at the surfaces of the electrolytes. Surprisingly the observed yields of oxidation or reduction products were often higher than expected by
Faraday’s law (positive deviation), e.g. as reported by Klemenc in the case of the oxidation of hydrochloric acid [24]. This positive deviation from Faraday’s law caused
Fig. 10.4 Experimental set-ups for different plasma (electro)chemical
experiments: (a) DC set-up of Klemenc, (b) set-up for the recovery
of metal from slags, (c) vapor-phase electrolytic deposition set-up of
Ogumi et al.
268 10 Plasma Electrochemistry with Ionic Liquids
temporarily strong interest in glow discharge electrolysis. The effect was attributed
to reactions driven by local high-temperature spots (hot spots) or by additional reactions caused by UV emission from the plasma. Negative deviations are usually
caused by a partial electronic conductivity in the electrolyte or by sputter loss of the
product. Klemenc concluded that the interface between a liquid electrolyte and a
plasma is comparable to the metal|plasma interface [24].
The electrolytic decomposition of organic compounds at plasma electrodes was
investigated both as an important side reaction and as a possible application itself.
Compared to Gubkin’s original experimental arrangement the setups were improved, e.g. the electrode areas were separated spatially and the vacuum generation
was improved by using better vacuum pumps. In addition, by reversing the applied
potential to the electrode immersed in the plasma, plasma-anodic experiments
were performed [28], but the mechanism of plasma-anodic processes remained
unclear. Brenner and other authors investigated the glow discharge electrolysis of
metal salt melts [29, 30]. At the interface between the salt melt and the plasma they
deposited dendrites of zinc, cobalt, copper, silver and nickel. Their investigations
can be considered as the forerunner of our current studies of ionic liquids.
Glow discharge electrolysis reappeared in the 1990s as plasma electrolysis. New
types of plasma reactors and discharges were developed and introduced for the deposition of either metals or metal oxides. Ogumi focused on solid electrolytes and
developed a method for the plasma electrochemical deposition of ion-conducting
metal oxides without liquid electrolytes [31, 32]. For this purpose he injected
metal-containing precursors (e.g. ZrCl4 and YCl3 ) into a capacitively coupled RFdischarge. The experimental set-up is shown in Figure 10.4. Yttria-stabilised zirconia was then formed on an oxygen-conducting substrate by electrolytic deposition
applying direct current between the substrate and a counter electrode within the
plasma. Vennekamp et al. studied this approach more systematically and quantitatively. They proved both the Faradaic character of plasma electrochemical processes
and the specific surface morphologies of plasma electrochemically grown solid electrolyte films [33, 34]. Today, plasma electrolysis of liquid electrolytes is applied to
waste water treatment [35]. In these applications ozone is formed in the discharge
region, which then reacts with organic waste molecules in the liquid solution.
During the last few years another method for plasma electrolysis has been developed. Thermal plasmas in the form of a plasma torch are used to melt metal oxides
and salts [36]. By applying an additional DC voltage via the thermal plasma electrode
(anode), the pure metal or metal alloys are deposited at the cathode which is located
in the melt (see Figure 10.4). The advantage of this process is that, due to the high
temperature of the plasma discharge, metal oxides can be used as electrolytes. The
process allows the direct recovery of pure metals from a slag of metal oxides [37].
The electrochemical cell is:
anode | metal oxide slag | cathode
He et al. reported the production of noble metal nanoparticles (Ag, Au, Pd, Pt)
by using plasmas [38], but no external voltage was applied, and the reduction was
10.4 The Stability of Ionic Liquids in Plasma Experiments 269
achieved with free electrons from the gas discharge under a floating potential. They
incorporated noble metal cations into a titanium dioxide gel by ion exchange and
reduced the cations by hydrogen low-temperature plasma treatment in a commercial plasma etcher. Inside the matrix nanoparticles of 2–10 nm in diameter were
produced.
Directly applying Gubkin’s concept of a plasma cathode, Koo et al. produced
isolated metal nanoparticles by reduction of a platinum salt at the free surface of
its aqueous solution [39]. The authors used an AC discharge as cathode over the
surface of an aqueous solution of H2 PtCl6 . Platinum particles with a diameter of
about 2 nm were deposited at the plasma|liquid electrolyte interface by reduction
with free electrons from the discharge.
metal anode | aqueous H2 PtCl6 solution | plasma | metal cathode
As indicated by Koo et al. in their paper and as shown in Figure 10.3, the gas
discharge over an aqueous solution is a localised corona discharge rather than an
extended plasma. This leads to a spatially highly inhomogeneous reduction process.
As demonstrated in Section 10.5 the use of ionic liquids leads to homogeneous and
extended gas discharges, contacting the whole surface area of the electrolyte. To
our knowledge, this type of spatially extended and homogeneous plasma/electrolyte
interface has not been investigated before.
10.4
The Stability of Ionic Liquids in Plasma Experiments
The voltages which are applied in order to ignite a DC discharge or which exist
across plasma sheaths are far beyond the electrochemical window limits of any
ionic liquid. But only a small part of the applied voltage (several hundreds of
volts) actually drops across a pure ionic liquid or an ionic liquid containing an
arbitrary metal salt situated beneath the burning plasma. Nevertheless, one may
expect severe decomposition reactions and a number of questions can be raised:
first, does the possible decomposition of ionic liquids lead to impurities of the
obtained particles? And if so, to what extent? Secondly, does it affect the deposition
negatively in other ways, e.g. by inhibiting the desired reaction? Thirdly, does the
decomposition reduce the solubility of the metal salts and restrict the reusability
of the ionic liquid? This section discusses some of these questions on the basis of
reports on ionic liquid decomposition reactions.
Our key reaction is the reduction of a metal salt dissolved in the ionic liquid
with free electrons from plasmas in order to obtain metal particles. Processes in
lithium ion batteries which employ ionic liquids with dissolved lithium salts can be
considered as close relatives. In both cases the dissolved metal salts crucially affect
the stability of the ionic liquid. Only if the anion of the dissolved metal salt is more
easily oxidised and its cation is more easily reduced in comparison to the ions of the
ionic liquid, the decomposition of the ionic liquid would be negligible; of course
270 10 Plasma Electrochemistry with Ionic Liquids
this would be the ideal case. At present, the objective of numerous research groups
is to avoid the electrochemical decomposition of the ionic liquid or at least to reduce
the extent of this side reaction. Hence, some decomposition reaction pathways of
ionic liquid ions are already well investigated. Most examples stem from the field
of lithium ion batteries, where the electrochemical stability of the electrolyte is a
crucial point, e.g. with regard to the rechargeability of the devices. These examples
will be briefly reviewed in this section, as ionic liquids with good stability towards
lithium will probably be suitable candidates for plasma electrochemical reactions.
The decomposition of the trifluoromethanesulfonate anion, CF3 SO3 − (OTf), and its
derivatives, the imidazolium and pyrrolidium cations, are primarily considered. At
the end of this section some proposals are given for reduction of the decomposition
of the ionic liquid or perhaps even to avoid it completely.
In general the electrochemical stability of an electrolyte is experimentally evaluated
by means of cyclic voltammetry. However, the determination of the electrochemical windows exhibits several problems. First, the electrochemical degradation or
breakdown of an electrolyte is an irreversible reaction, thus there is no theoretical
redox potential [40, 41]. Passivation of the electrodes often makes it difficult to
identify the onset of the reaction due to inhibition of further reactions [40, 42].
Some of the already used electrolytes, and also future candidates for lithium ion
batteries, are based on organic solvents like propylene carbonate (PC), vinylene
carbonate (VC), 1,2-dimethoxyethane (DME), etc. containing a lithium salt instead
of an ionic liquid including a lithium salt. The organic solvent molecules of the
electrolyte decompose simultaneously beside the electrolyte ions [40, 43, 44]. Thus
different orders of anion stabilities were obtained for different electrolyte compositions [40, 45, 46]. Of course, impurities can lead to a similar phenomenon. The
connection between the disintegration reactions of electrolyte salt and electrolyte
solvent as well as the influence of their composition ratio was demonstrated, for
example, by Rahner [47]. Koch et al. tried to circumvent the problem using pure
ionic liquids to investigate the stability of anions, in order to find the most suitable
counterion for the lithium ion, without distortion by a solvent [45]. However, once
some ions of the ionic liquid were reduced or oxidised they could form neutral
organic molecules (as will be described below) acting as impurities and leading to
similar problems as for PC, VC, DME, etc. lithium salt solutions.
Another point is that the reduction and oxidation potential limits (electrochemical window) are defined as the potentials at which the current density reaches a
predefined value that is arbitrarily chosen [40, 48]. Ue et al. also mention that the
same problem arises in the choice of the sweep rate [40]. For example Egashira
and coworkers obtained a log I–U line shifted to a higher position at a faster potential scan in comparison to a slower scan because of non-Faradaic currents such
as the larger charging currents of the double-layer, and the decomposition of impurities [41]. The last factor affecting the electrochemical window is the electrode
itself, its composition and its morphological surface structure, which defines the
electrocatalytic properties [40].
Johansson compared several theoretical measures in order to find a theoretically
calculable substance property which correlates to the oxidation potential of anions
10.4 The Stability of Ionic Liquids in Plasma Experiments 271
and thus allows the prediction of the anodic stability limits of anions from different
anion families. First, he compared the highest occupied molecular orbital (HOMO)
energy, which he converted as all the other energy changes to electrochemical
potentials additionally corrected for the Li+ /Li0 electrode, with the experimental
literature oxidation potentials. Secondly, he used the vertical transition energy,
which is the energy difference between the anion and the corresponding unrelaxed
neutral radical following the Frank–Condon principle. The first two quantities are
gas phase energies by definition. In order to mimic real battery electrolyte species
better he carried out additional single point calculations for the anions and their
radicals using a self-consistent reaction field method to get the corresponding
vertical free energy [49]. All of the considerations above are also valid for the case
of a metal salt dissolved in an ionic liquid.
Nakajima et al. considered the decomposition of the trifluoromethanesulfonate
anion [OTf], in the context of aluminum corrosion in lithium ion batteries. As
a result of the electrochemical oxidation of [OTf], C–F active species like CF2
emerge, which either lead directly to corrosion of the aluminum or to a disproportionation reaction that forms atomic carbon which also corrodes the aluminum.
Additionally, S–O-containing species are created during the oxidation of [OTf].
The authors were able to confirm their suggested decomposition products by
energy dispersive X-ray (EDX) spectra. These findings should also be valid for
the corresponding bis(trifluoromethanesulfonyl)amide, (CF3 SO2 )2 N− [NTf2 ], and
tris(trifluoromethanesulfonate)methide, (CF3 SO2 )3 C− (CTf3 ), respectively, since
they consist of comparable building substructures [50].
Witkamp and coworkers investigated the reductive decomposition of 1-butyl1-methylpyrrolidinium bis(trifluoromethanesulfonyl)amide, [BMP][NTf2 ], and 1butyl-3-methylimidazolium tetrafluoroborate, [BMIM][BF4 ], by a combination of
simple and inexpensive semi-empirical calculations (Spartan ‘04 modelling program, PM3) and experiments, where a voltage (8 V) larger than the electrochemical
windows of the considered room-temperature ionic liquids was applied at room
temperature for 3 h [51]. Subsequently, the degradation product of [BMP][NTf2 ]
was analysed via gas chromatography and mass spectroscopy (GC-MS) as well as
nuclear magnetic resonance (NMR) spectroscopy, whereas for [BMIm][BF4 ] only
NMR spectroscopy was used. In general the cations were reduced more easily
than the anions on the cathodic limit. One exception is the heptachloroaluminate
(Al2 Cl7 − ) anion of the acidic chloroaluminate ionic liquid family. After electron
transfer from the electrode to the cation the obtained radical can undergo several
possible decomposition and rearrangement pathways. Witkamp et al. calculated
the energies of all conceivable breakdown products. The main pathway was then
found by comparison of the several product energies.
After formation of the analogue radical of 1-butyl-1-methyl-pyrrolidinium, it
can decompose into methylpyrrolidine and a butyl radical, whereupon the energy of the products amounts to −61 kJ mol−1 in vacuum (see Figure 10.5, Eq.
(2)). A second possible product represents the dibutylmethylamine radical (E =
−43 kJ mol−1 ) resulting from a ring opening reaction (see Figure 10.5, Eq. (1)). The
third and least likely product combination is butylpyrrolidine and a methyl radical
272 10 Plasma Electrochemistry with Ionic Liquids
Fig. 10.5 Decomposition pathways of 1,1-butylmethylpyrrolidinium according to Ref. [51].
(E = −21 kJ mol−1 ) (see Figure 10.5, Eq. (3)). For all decomposition pathways
Witkamp et al. found experimental evidence.
In the case of the 1-butyl-3-methylimidazolium cation a stable radical is obtained
[50], due to the stabilizing interaction of the singly occupied p orbital of the C2
carbon atom with the p orbitals of the free electron pairs of the two adjacent
nitrogen atoms. That is why the decomposition pathway would need 75 kJ mol−1
[51]. A dimerization needs only an energy of 33 kJ mol−1 , whereupon two 1-butyl-3methylimidazolium radicals are coupled to each other via their C2 atoms of the ring
system (see Figure 10.6, Eq. (4)). Another reaction could be a disproportionation,
i.e. a hydrogen abstraction from one radical to another, leading to 1-butyl-3-methyl2,3-dihydro-1-H-imidazole and a zwitterionic structure (see Figure 10.6, Eq. (5)).
Finally, they suggest that a radical addition of two imidazolium radicals to the C–C
double bond of the respective partner radical could take place, forming a cage-like,
Fig. 10.6 Decomposition pathways of 1-butyl-3-methylimidazolium according to Ref. [51].
10.4 The Stability of Ionic Liquids in Plasma Experiments 273
Fig. 10.7 Decomposition pathways of 1-butyl-3-methylimidazolium via
a biradical transition state.
neutral structure, where the former two independent imidazolium radicals are
connected to each other via two new bonds (see Figure 10.6, Eq. (6)). But the two
radicals cannot form a product like the suggested one where the two double bonds
remain.
After a radical addition of one imidazolium to the C–C double bond of a second
one, a biradical results. This biradical could react with the remaining double bond
to give a cage structure but with three bonds formed instead of only the two
proposed by Witkamp et al. and a concomitant vanishing of the two double bonds
(see Figure 10.7, (7)). Compound (7) would be highly strained, hence, very unlikely.
An internal rearrangement followed by a recombination of the two radical centers
of the biradical could be another possible product (see Figure 10.7, (8)) but, as it
is known, a transfer of a saturated alkyl group is also very unlikely to occur. Thus,
they found evidence only for the first two reaction pathways in the NMR spectra of
the decomposition of 1-butyl-3-methylimidazolium tetrafluoroborate [51].
In the case of metal deposition at the ionic liquid|plasma interface two possible
reduction processes can conceivably take place. First, the metal cations of the
dissolved metal salt can be reduced. Secondly, the cations of the ionic liquid can
be reduced to neutral radicals, which can further react as described by Witkamp
and as summarized above. As a first guess of which process is preferred, the rate
constants of the reaction for example of silver ions (k ≥ 3.2 × 1010 L mol−1 s−1 )
and imidazolium ions (k ≤ 4.3 ×109 L mol−1 s−1 ) with hydrated electrons, taken
from the data collection of Buxton et al., can be considered [52]. Thus, as long as
sufficient silver ions are still in solution the reduction of the imidazolium cations
of the ionic liquid represent the minor reaction pathway and the ionic liquid should
not decompose significantly.
How can disintegration of the ionic liquids be avoided or reduced? The cations
of the ionic liquid have to be stabilized, e.g. via delocalisation of the positive charge,
so that they are less eager for electron uptake in comparison to the dissolved metal
274 10 Plasma Electrochemistry with Ionic Liquids
salt cations. Once a reduction of the cation of the ionic liquid occurred, it would
be advantageous if the radical undergoes a quenching reaction with a metal ion,
i.e. an electron transfer, instead of decomposing or forming a dimer. To prevent
oxidation of the anions a metal salt should be used with an anion that is much
easier to oxidise than the anion of the ionic liquid.
A completely different approach might be the use of radio frequency plasma
instead of a DC plasma. The ignition and sustainment of the plasma is decoupled
from the application of voltages to the electrodes that are now used only for electrochemical reactions. Another method which has been proven to be quite successful
is the application of an U-shaped tube in order to avoid an IR-drop over the ionic
liquid (see Figure 10.2). Unfortunately, this set-up led to a large size distribution
of the obtained particles but it showed that RF plasma could further improve the
stability of the ionic liquids during the metal deposition process.
10.5
Plasma Electrochemical Metal Deposition in Ionic Liquids
Considering the different general reaction schemes for processes at plasma|ionic
liquid interfaces, the plasma-cathodic reduction of compounds dissolved in an ionic
liquid is the most obvious application. In fact, the plasma-cathodic reduction of dissolved metal salts has recently emerged as a first example of plasma electrochemical
processes with ionic liquids [53–55]. Up to now deposition of the metals Ag [53, 54],
Pt, Cu and Pd [55] from different ionic liquids has been tested. The experimental
approach is based on previous work on processes at the interface between a solid
ionic conductor and a plasma [33, 56, 57] but it can, in principle, also be directly
traced back to original works on “glow discharge electrolysis” of aqueous solutions
by Gubkin [1], Klüpfel [58] and Klemenc [59], as summarised in Section 10.3. As
shown schematically in Figure 10.8, the prototype experiment represents basically
a cathodic reduction of a precursor (starting material), dissolved in the ionic liquid,
with free electrons from the plasma phase – driven by the external electric field.
Electrons are generated in the cathode region of the plasma and are driven towards
the surface of the ionic liquid, where they reduce the dissolved metal compounds.
In essence, we use the free surface of the ionic liquid in contact with a plasma as
the electrode interface, leading to the deposition of solid products dispersed in the
ionic liquid at the surface.
The minimal experimental set-up (Figure 10.9) for a DC glow discharge experiment consists of a glass tube with two electrodes, of which the bottom electrode is
made of either an inert metal like Pt or of a consumable bulk metal or semiconductor, thus keeping the concentration of the electroactive cation in the ionic liquid
constant (side reactions at the anode are neglected). The pressure in the reactor is
controlled by adjusting the mass flow of a gas (mostly argon in the case of metal
and semiconductor deposition) and by a vacuum pump.
Deposition of silver metal: In order to exemplify the plasma electrochemical deposition (PECD) technique in ionic liquids we first deposited silver
10.5 Plasma Electrochemical Metal Deposition in Ionic Liquids 275
Fig. 10.8 Schematic experimental set-up for the deposition of metal
nanoparticles by plasma electrochemical reduction of a metal salt dissolved in an ionic liquid at room temperature.
nanoparticles from both a AgNO3 and a AgCF3 SO3 solution in ultrapure 1-butyl3-methylimidazolium trifluoromethanesulfonate ionic liquid ([BMIM][TfO]) by the
use of an argon plasma. Similar experiments with 1-ethyl-3-methylimidazolium
trifluoromethanesulfonate ([EMIM][TfO]) and 1-butyl-1-methylpyrrolidinium trifluoromethanesulfonate ([BMP][TfO]) have also been successful. In the first
Fig. 10.9 DC discharge over [BMIM][TfO].
276 10 Plasma Electrochemistry with Ionic Liquids
experiments saturated solutions of silver nitrate in the ionic liquid were used,
later we used silver trifluoromethanesulfonate (Aldrich, ≥99 %) due to its better
solubility in the ionic liquids. The solutions typically contained 0.3 g of CF3 SO3 Ag
in 10 ml [BMIM][TfO], that corresponds to a concentration of about 0.15 mol l−1 or
a molar fraction of 0.026.
A 1 × 1 cm2 platinum sheet was used as anode and a hollow platinum cylinder
of about 1.5 cm height and 0.75 cm diameter was used as cathode, placed in the
gas phase above the ionic liquid at a distance of typically 10 cm. The glass reactor
(2.5 cm diameter) was filled with 10 ml of the silver salt solution, fully covering the
anode with approximately 2 cm liquid phase. The reactor was then evacuated, and
the pressure was controlled to 100 Pa (argon atmosphere). Ascending bubbles were
usually observed for some minutes due to emerging gases originally dissolved in
the ionic liquid. After this outgassing the electric voltage was switched on. Drawing
a current of 10 mA under galvanostatic conditions (corresponding to 2 mA cm−2 ),
the voltage stabilised typically at about 470 V. The formation of a small number of
bubbles at the anode could be observed after some time during the electrochemical
experiment. This effect was more distinct with [EMIM][TfO] as solvent.
The solution of the silver salts in the ionic liquids was almost transparent and
colorless before the deposition experiment was started (Figure 10.10 (a)). After the
onset of the glow discharge process the following observations were made: (i) A
homogeneous plasma burnt with a pale pink/blue optical emission between the
Fig. 10.10 Plasma electrochemical deposition of silver nanoparticles
at the free surface of [BMIM][TfO].
10.5 Plasma Electrochemical Metal Deposition in Ionic Liquids 277
upper electrode and the surface of the ionic liquid (Figure 10.10 (b)). During the
initial period of the reaction the optical emission of the plasma changed slightly,
indicating a change in the plasma composition. (ii) Starting from the surface of
the ionic liquid a dark cloud appeared reproducibly in the ionic liquid (see. Figure
10.10 (b) and (c)). Upon longer reaction time this dark region widened until a
completely dark ionic liquid was obtained (see Figure 10.10 (d)). Thus, the product
phase spreads completely across the ionic liquid. (iii) At the platinum anode we
observed the formation of gas bubbles. The amount of gas bubbles corresponds to
the electric current across the cell, and we assume at this point that either the NO3 −
or the CF3 SO3 − anion is oxidised, liberating oxygen and/or the products suggested
in Section 10.4. After 5 to 10 min of reaction time the plasma electrolysis was
stopped. After some minutes the homogeneous product region started to disperse
and later to sediment at the bottom of the ionic liquid. Using an ultracentrifuge, the
sedimentation process could be accelerated and the liquid phase of the dispersion
could be removed and replaced easily by distilled water. Using ultrasound, the
sediment could be dispersed again. Several of these cleaning steps were used to
remove fully any remnants of the ionic liquid, thus purifying the reaction product.
Images obtained by high resolution scanning electron microscopy (HRSEM)
and high resolution transmission electron microscopy (HRTEM) (Figures 10.11
and 10.12) show aggregates of particles with average sizes in the nanometer region.
Energy dispersive X-ray (EDX) spectra were recorded in scanning and transmission
mode, both confirming that the aggregates mainly consist of silver with traces of
ionic liquid. In transmission mode, we were able to focus on single nanocrystals,
thus evidencing that they consist of pure silver; in particular no oxygen impurities
could be detected. This finding was supported by selected area electron diffraction
(SAED) and HRTEM. The diffraction patterns recorded on the aggregates show
Bragg reflections located on concentric rings. The d-values determined from the
diameter of the rings are fully consistent with those of pure silver (Figure 10.12)
the profile of the reflections underlines the high crystallinity of the silver particles.
Fig. 10.11 SEM image of the silver nanoparticles.
278 10 Plasma Electrochemistry with Ionic Liquids
Fig. 10.12 TEM image and size distribution of the silver nanoparticles.
The substance was found to consist exclusively of silver nanoparticles (for particle
size distribution see Figure 10.12). Silica nanoparticles were frequently found in
the product. We attribute this to sputter effects of the glass reactor walls. These
sputter effects can easily be reduced by a more sophisticated design of the plasma
reactor.
Deposition of copper metal: Since Cu(II) is the preferred oxidation state of copper,
Cu2+ salts are more stable and more available, hence, in a technical application it
would be favorable to use them as starting material. We tried to reduce Cu(CF3 SO3 )2
dissolved in [EMIM][TfO], [BMP][TfO] and [BMIM][TfO] with an argon plasma (gas
pressure 100 Pa) as well as with a nitrogen plasma (100 Pa), respectively. Additional
experiments with Cu(CF3 SO3 )2 dissolved in [EMIM][TfO] and Ar/H2 plasmas were
carried out, with the distance between the hollow cathode in the gas phase and the
surface of the ionic liquid metal salt solution being 3, 45 and 100 mm. Moreover,
for the 3 mm distance several experiments with different gas pressures from 50 to
500 Pa were carried out.
Virtually all observed reactions proceeded in the same manner: A platinum
electrode was located in the middle of the ionic liquid, and a platinum hollow
electrode was placed in the gas phase above, as described for silver. After 5 minutes
a light brown cloud appeared in the upper half of the ionic liquid – Cu(CF3 SO3 )2
solution. During the next 5 minutes the triple phase boundary between the glass
wall, the ionic liquid and the blue–pink plasma grew darker and brown threads
starting from this dark region spread down into the light brown area. Then black
particles emerged at the triple phase boundary and some of them finally sank down
to the bottom of the ionic liquid (see Figure 10.13). Later during the reaction, the
10.5 Plasma Electrochemical Metal Deposition in Ionic Liquids 279
Fig. 10.13 Ar/H2 plasma (3 parts Ar to 1 part H2 ) burning over
Cu(CF3 SO3 )2 dissolved in [EMIM][TfO] with a brown cloud and black
deposits. The distance amounts to only 4.5 cm, thus the plasma consists mostly of dark space (Faraday space).
lower half of the ionic liquid also became brown, but even after 1 h there remained
a distinction between the upper and the lower half of the ionic liquid phase in terms
of the brightness of the brown color.
Subsequent investigation of the obtained deposit with EDX revealed that it indeed
consisted mainly of carbon and the residues of decomposed ionic liquid. Only a
small amount of copper was found so the question remains as to whether this is
copper metal or merely enclosed Cu+ or Cu2+ . Hence, at this point we conclude
that copper deposition from Cu(II) salts does not easily result in Cu(0) deposition.
Why was the reaction not successful in the case of copper? Can the rate constants
(k) for the reaction of ions with hydrated electrons tabulated by Buxton et al. be used
again, as in Section 10.4 in the case of silver ions, to estimate whether this reduction
is kinetically reasonable at all? In general the reaction of (metal) ions with hydrated
electrons is significantly affected by the counter ions of the considered ions and
their complexation ligands. Moreover, the rate constants are given only for the
reaction of Cu(I) with a hydrated electron to Cu(0), where k amounts to 2.7 ×
1010 L mol−1 s−1 , and the reaction of Cu(II) with a hydrated electron to Cu(I) with
k ≥ 2.9 ×1010 L mol−1 s−1 in the neutral/acid pH range, but not for the reaction of
Cu(II) to Cu(0). A two-electron process is much less likely to occur and one would
expect that the rate constant of this process would be lower than the k values for the
two single reduction steps mentioned. The first k value [Cu(I) → Cu(0)] suggests
that Cu(I) salts could be a proper starting material. The disadvantage of Cu(I) salts is
that the stable ones, like the Cu(I) halides, are inherently insoluble in ionic liquids
due to their covalent bonding character which leads to a diamond analogue zinc
blende (sphalerite) structure. Those Cu(I) salts which do not possess this highly
polymeric structural character are very sensitive to air and moisture. However, if a
Cu(I)-containing ionic liquid is made by electro-oxidation of metallic copper directly
280 10 Plasma Electrochemistry with Ionic Liquids
in the ionic liquid the plasma electrochemical reduction to elemental copper should
be feasible.
Deposition of platinum metal: In the case of platinum no solid product was found.
The ionic liquid darkened more and faster the smaller the distance between the surface of the ionic liquid [EMIM][TfO] containing tetrabutylammonium hexachloroplatinate ([n-Bu4 N]2 [PtCl6 ]) and the Ar/H2 -plasma (3:1, overall pressure 100 Pa)
was chosen. So far no other ionic liquid has been tested. The rate constant for the
reduction of the tetrabutylammonium ion with a hydrated electron is only 1.4 ×
106 L mol−1 s−1 , hence the main rival pathway for reduction of platinum(IV) is the
reduction of the imidazolium ion of the ionic liquid. As in the case of copper, a
suitable platinum salt – maybe made by electro-oxidation of metallic platinum in a
suitable ionic liquid – has to be found.
Deposition of palladium metal: It was possible to deposit palladium nanoparticles by reduction of ammonium tetrachloropalladate ([NH4 ]2 [PdCl4 ]) dissolved in
[EMIM][TfO]. The product yield was only about 2.5% compared with the theoretical
value and was reached after 25 min. A longer application of the plasma did not
lead to a significant increase in the amount of product. The lower yield and the
slower reaction process might be a sign of the much more difficult two-electron
reduction process compared to the one-electron process in the case of silver. Koo
et al. stated that they did not obtain any platinum nanoparticles using a plasma
without H2 gas [39]. In the case of palladium, nanoparticles formed when a Ar/H2
plasma (3:1, 100 Pa) was applied and when a pure Ar plasma was used. The HRSEM
(Figure 10.14) picture reveals the high homogeneity of the particles which are all
about 5 nm in diameter (HRTEM, Figure 10.15), only a few are bigger but that can
be neglected. In Figure 10.15 is an additional SAED picture that exhibits the diffuse
rings which are typical for equally seized nanoparticles.
The successful deposition of silver and palladium nanoparticles proves the applicability of the PECD concept in ionic liquids. We expect that the cathodic deposition
of other elements can be run in the same way under comparable conditions from
Fig. 10.14 HRSEM picture of the obtained palladium nanoparticles.
10.5 Plasma Electrochemical Metal Deposition in Ionic Liquids 281
Fig. 10.15 HRTEM/SAED picture of the palladium nanoparticles. The
letters in the SAED picture represent the lattice indices: a = 111, b =
200, c = 220, d = 311.
suitable starting materials. In the case of highly reactive materials like titanium a
hydrogen plasma may be used in order to avoid immediate re-oxidation by residual
oxygen in the plasma phase. We believe that PECD in ionic liquids represents a versatile method with a potentially broad field of applications in the synthesis of metal
and semiconductor micro- and nanoparticles. It can be expected that the physical
properties of different ionic liquids, electric current density, the temperature, the
chemistry of the plasma phase and also the convection in the liquid phase will
influence the morphology of the reaction product and, thus, may be used profitably
as control parameters.
Comparing this approach with previous work – except the studies on solid electrolytes – ionic liquids have two distinct advantages over aqueous or organic solvents: (i) Due to their extremely low vapor pressure ionic liquids can be used without
any problem in standard plasma vacuum chambers, and the pressure and composition in the gas phase can be adjusted by mass flow controllers and vacuum pumps.
As the typical DC or RF plasma requires gas pressures of the order of 1 to 100 Pa,
this cannot be achieved with most of the conventional liquid solvents. If the solvent
has a higher vapor pressure, the plasma will be a localised corona discharge rather
than the desired extended plasma cloud. (ii) The wide electrochemical windows
of ionic liquids allow, in principle, the electrodeposition of elements that cannot
be obtained in aqueous solutions, such as Ge, Si, Se, Al and many others. Often
this electrodeposition leads to nanoscale products, as shown e.g. by Endres and
coworkers [60].
The development of methods for the reproducible and continuous production of
metal and semiconductor particles with a typical size on the nanoscale is still an
active field of research [61–65]. The existing synthetic methods for isolated nanoparticles can be categorised into two major groups: (i) Gas or plasma phase-based
preparation from gaseous or liquid precursors, (ii) preparation of nanoparticles in
282 10 Plasma Electrochemistry with Ionic Liquids
liquid solution by reduction or by precipitation, often with the help of template
molecules or micelles [66]. Electrochemical methods are hardly used for the preparation of isolated nanoparticles, mainly because the reaction products are usually
deposited as compact materials at a solid electrode rather than as free particles.
A particularly successful method is pulsed electrodeposition (PED), a well-known
technique in galvano plating [67], which was introduced into nanoscience by Erb
et al. [68], mainly for n-Ni deposition. This concept was expanded by Natter and
Hempelmann to deposit n-Pd [69], n-Cu [70], n-Fe [71] and n-Cr [72]. They were
also able to deposit alloys like for example Nix Fe1–x or Nix Cu1–x [73]. Thus all
metals with E > 0 V (vs. NHE) can be electrodeposited in this way from aqueous
electrolytes [74, 75]. The electrodeposition of nanocrystalline metals and nanoscale
semiconductors in ionic liquids is summarized in Chapters 6 and 8 .
Koo et al. have recently published results from PECD of Pt nanoparticles in an
aqueous solution of H2 PtCl6 [39]. They also observe the formation of relatively
small particles with a typical diameter of 2 nm. From the electrochemical point of
view, water is not a suitable solvent for plasma electrochemical processes, due to
its relatively high vapor pressure, even at low temperatures.
10.6
Conclusions and Outlook
The plasma|ionic liquid interface is interesting from both the fundamental and the
practical point of view. From the more fundamental point of view, this interface
allows direct reactions between free electrons from the gas phase without side reactions – once inert gases are used for the plasma generation. From the practical
point of view, ionic liquids are vacuum-stable electrolytes that can favorably be used
as solvents for compounds to be reduced or oxidised by plasmas. Plasma cathodic
reduction may be used as a novel method for the generation of metal or semiconductor particles, if degradation reactions of the ionic liquid can be suppressed
sufficiently. Plasma anodic oxidation with ionic liquids has yet to be explored. In
this case the ionic liquid is cathodically polarized causing an enhanced plasma ion
bombardment, that leads to secondary electron emission and fast decomposition
of the ionic liquid.
Currently only a few exploratory experimental studies have been reported, and
much work has still to be done in order to explore fully the properties and characteristics of plasma|ionic liquid interfaces. Currently it is still too early to comment
if technical applications will be found. From the economic point of view, both ionic
liquids and plasmas are comparatively expensive media, therefore only applications
which show significant advantages compared to more conventional routes will be
successful.
Introducing reactive gases to the plasma phase may even lead to the formation of
metal or semiconductor compounds, extending the experimental possibilities even
further. From the physicochemical point of view, plasma electrochemical deposition is a highly interesting interfacial phenomenon, linking plasma chemistry and
References
electrochemistry and utilizing nucleation under conditions far from equilibrium.
A systematic investigation of this process is required in order to understand the
nucleation and growth process in detail.
Acknowledgments
Parts of the experimental work have been performed in close collaboration with
Dr. Lorenz Kienle at MPI für Festkörperforschung Stuttgart and Professor Frank
Endres at TU Clausthal-Zellerfeld, Germany. The support of the DFG (Priority
program “Ionic Liquids”, projects Ja 648/13-1 and En 370/16-1) and the Funds of
the Chemical Industry (FCI) is gratefully acknowledged.
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287
11
Technical Aspects
Debbie S. Silvester, Emma I. Rogers, Richard G. Compton, Katy J. McKenzie,
Karl S. Ryder, Frank Endres, Douglas MacFarlane, and Andrew P. Abbott
11.1
Metal Dissolution Processes/Counter Electrode Reactions
While the subject of this chapter may seem counter to the title of the book, metal
dissolution is vital in numerous aspects of metal deposition, counter electrode
processes, pre-treatment protocols and electropolishing. This chapter outlines the
current state of understanding of metal dissolution processes and discusses in
some detail an electropolishing process that has now been commercialised using a
Type III ionic liquid.
11.1.1
Counter Electrode Reactions
Little or no information is available in the open literature about counter electrode
reactions occurring during deposition processes in ionic liquids. No data exist on
anodic dissolution efficiencies and hence many practical issues associated with
process scale-up are unknown at present. In ionic liquids the issues associated with
pH can largely be ignored since the passivating layers either dissolve, e.g. in high
chloride media, or trans-passive corrosion occurs at high enough over-potentials.
This means that even metals such as Cr and Al have been used as soluble anodes
as they can be readily oxidised in ionic liquids.
Most work to date has either used soluble anodes or has not considered the
anodic reaction. A limited amount of information has been collated on the electrochemical windows of ionic liquids but this tends to be on either platinum or
glassy carbon, which is not necessarily suitable for practical plating systems [1].
The anodic limits of most liquids are governed by the stability of the anion, although pyridinium and EMIM salts are sometimes limited by the stability of the
cation. The widest electrochemical windows are obtained with aliphatic quaternary
ammonium salts with fluorous anions. A selection of potential windows is given in
Chapter 3.
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
288 11 Technical Aspects
The optimum process would ideally involve the use of soluble anodes, as the
over-potential required to drive the deposition process will be small. This is especially important with ionic liquids because the ohmic loss across the cell can be
significant. In aqueous solutions the use of soluble anodes is not often possible
due to passivation of the electrode surface at the operating pH.
While no systematic studies have been carried out, to date the only metal that
we have been unable to electrochemically dissolve in eutectic-based ionic liquids is
iridium. Although our research has not studied all metals we have found that even
Pt, Au and Ti can be made to dissolve in eutectic-based liquids. Hence, in principle,
soluble anodes could be used for the deposition of most pure metals from ionic
liquids. This is, however, a considerable over-simplification and a number of factors
need to be considered before employing a soluble metal anode. In ionic liquids with
discrete anions attention needs to be given to the ligand present that will solvate
the dissolving metal. It is highly unlikely that an unsolvated anion could exist in an
ionic liquid and no evidence has been obtained to date to suggest otherwise. Metals
are known to be soluble in ionic liquids based upon Tf2 N− and BF4 − anions but
the nature of the metal complexes is unknown [2]. It could be that dative bonds are
formed with oxygen or fluorine moieties or it could be that trace water acts as a
ligand.
In eutectic-based ionic liquids, the chloride ions act as strong ligands for the
oxidized metal ions, forming a range of chlorometallate anions. The free chloride
ions are present in very low concentrations as they are complexed with the Lewis
acidic metal ions and so the dissolution of metal ions must lead to a complex series
of equilibria such as
−
2+
4 ZnCl−
↔ Zn2 Cl−
3 + Zn
5 + Zn3 Cl7
(11.1)
Therefore it can be seen that metal dissolution is easier in Lewis basic melts. The
zinc and aluminum deposition processes, which are by far the most frequently
studied, are almost totally reversible. Since these metals have no other stable oxidation states the deposition and dissolution processes are very efficient [3–6]. This has
the distinct advantage that the composition of the ionic liquid remains constant and
the process becomes the removal of metal from one electrode and its deposition on
the other electrode.
Graphite has been used, but it fragments following electrolysis at high overpotentials leaving a black powdered residue at the base of the cell. Glassy carbon
has been used extensively in voltammetric studies, but its stability at high applied
current densities has not yet been tested. While anodic dissolution of metals may
be advantageous for some metal deposition processes, for others it may prove
problematic e.g. for alloy deposition or for electrowinning applications. In some
cases, e.g. chromium, it may be impossible to obtain electrodes because the metals
are not commercially available in a suitable form from which to make electrodes.
Another issue that needs to be considered is metals that exist in different oxidation states, e.g. Cr and Mn. The use of inert anodes could potentially lead to the build
up of metals in a higher oxidation state. However, unlike aqueous solutions, ionic
11.1 Metal Dissolution Processes/Counter Electrode Reactions 289
liquids tend to lack strong ligands such as oxygen which can stabilise higher oxidation states and this tends to negate the potential problem. Again no information
exists in the open literature but experiments carried out in our laboratory showed
that this was not an issue. The type II ionic liquid choline chloride: 2CrCl3 ·6H2 O
was studied for the deposition of chromium using a soluble chromium anode [7, 8].
Prolonged electrolysis was carried out over several months using the sample liquid
and at the end of this period the liquid was analyzed. It was found that there was
no discernible breakdown of the choline cation and no chromium species other
than Cr(III) was detected. The chromium content of the liquid was approximately
the same as the initial sample and the only discernible change was that the water
content of the liquid had decreased, presumably due to both anodic and cathodic decomposition. The chromium rod anodes were also severely etched over the process
confirming that they can act as soluble anodes.
The anodic processes occurring in the ionic liquids containing discrete anions
have not been well characterized. They will be extremely complex as the fluorinated
anions tend to be very stable and act as poor ligands. This means that both metal
dissolution and solution oxidation will be difficult. If inert anodes, e.g. iridium
oxide coated titanium, are used then it is difficult to envisage what the anodic
process will be and this is important to determine as the systems will have to
operate at relatively high current densities. Electrolysis of the ionic liquid itself
must be avoided from the obvious economic viewpoint but also from the practical
perspective that most electrolytes will give off toxic fluorinated products. This is
analogous to the primary production of aluminum by the Hall–Héroult process
where perfluorocarbons (PFCs) CF4 and C2 F6 are produced at the anode. In the
United States, aluminum smelting is the primary source of PFC emissions [9, 10].
Hence it can be concluded that little or nothing is known about the practical issues
associated with suitable anode design. With such an array of ionic liquids, metals
and deposition conditions available it is impossible to make specific predictions of
how all anodic materials will behave. Some general conclusions can, however, be
drawn, which should be good starting points from which to design specific processes. Where possible soluble anodes should be used as these improve process
efficiency and bath longevity. Decomposition of the ionic liquid should be avoided
at all times as it is naturally costly to reprocess the liquid and shortens its use. Processes are in general more current efficient than corresponding aqueous systems.
11.1.2
Pre-treatment Protocol
The surfaces of metal substrates require preparation and cleaning in order to ensure adhesion and effectiveness of the finishing or coating treatment. Cleaning is
also employed for the removal of oil, grease or scale from metal surfaces. Abrasive blasting, acid washes, multi-stage chemical cleaning and priming are some
of the techniques used for surface preparation and cleaning [11]. Typical surface
preparation and cleaning operations such as abrasive blasting are used for removal
of paint, rust and scale prior to painting or refinishing. Organic solvents are used
290 11 Technical Aspects
for degreasing; aliphatic petroleum, aromatics, oxygenated hydrocarbons and halogenated hydrocarbons are all applied to metal surfaces [11].
Electrocleaning techniques make use of a direct, reverse, or periodic reversed
electric current, in combination with an alkaline cleaning bath for the removal of soil
and smut and the activation of the metallic surface. The workpiece may be set up as
cathode or anode. Electrocleaning baths contain a solution with ingredients similar
to those of alkaline cleaning and can be operated either at ambient temperatures
or in the range 40–80 ◦ C [12]. To date no processes have demonstrated this in
conjunction with an ionic liquid but there is no technical reason why this should
not be possible and in cases where the substrate etching is not reversible it may be
advantageous (vide infra).
In principle. there is no difference between the pretreatment that a metal should
undergo before immersion in an ionic liquid or in an aqueous solution. The sole
difference is that the workpiece must be dry before immersion in the ionic liquid.
The sensitivity of the ionic liquid to water content is dependent upon the ionic
liquid. Eutectic-based ionic liquids are less sensitive to water content than liquids
with discrete anions. This is thought to be due to the ability of the chloride anions
in the former interacting strongly with the water molecules, decreasing their ability
to be reduced. Especially with AlCl3 -based ionic liquids water has to be strictly
avoided.
Most pre-treatment protocols studied so far follow the aqueous protocol quite
closely. Good adhesion is obtained by degreasing in a chlorinated solvent, followed
by an aqueous pickle in aqua regia, a water rinse and drying. A typical pre-treatment
protocol is shown schematically in Figure 11.1.
Several groups have then used an anodic etch in the ionic liquid prior to deposition. Anodic etch potentials and times are dependent on the substrate and the
Fig. 11.1 Flow chart for the pretreatment of substrates before electrodeposition in ionic liquids
11.1 Metal Dissolution Processes/Counter Electrode Reactions 291
Fig. 11.2 AFM image of aluminum etched for 20 s at 10 V in a Type
III eutectic of 1 choline chloride: 2 ethylene glycol (right side of image
masked during experiment).
ionic liquid used but generally less than 1 min is required to achieve a suitably
etched substrate. The etch process has the dual purpose of removing any remaining oxide film and roughening the surface to act as a “key” for the coating layer.
Metal oxide dissolution is easier in ionic liquids containing metal ions that are good
oxygen scavengers, e.g. Type I eutectics, because the oxygen scavengers ‘mop’ up
any oxygen moieties which have been generated during the etch process.
Figure 11.2 shows an atomic force microscopy (AFM) image of an aluminum
electrode etched for 20 s at 10 V in a Type III eutectic of 1 choline chloride: 2
ethylene glycol. The right side of the image was masked with a lacquer during
the experiment which was then removed before the sample was imaged. It can be
seen that the left side of the sample was significantly etched even during the short
duration of the anodic pulse. Dissolution rates of between 50 and 150 µm min−1
are observed under these conditions and result in a pitted surface. The sample in
the image is too well etched for practical purposes and hence shorter times or lower
over-potentials should be employed.
Figure 11.3 shows an analogous experiment for a copper electrode and it can
be seen that significantly less metal is removed in the same period. The surface
has approximately the same roughness as the original sample but there are more
micro-pits on the sample, leading to a better key with the subsequently deposited
film. Figure 11.4 shows that the etch rate for aluminum is almost three times
that of copper under the same conditions. These figures show that in ionic liquids
passivating films on electrode surfaces play a smaller role in controlling metal dissolution kinetics. The metals behave more characteristically, as would be predicted
by their standard reduction potentials, i.e. metals with a more negative reduction
potential are easier to etch.
Many plating protocols advocate the use of a ‘flash’ step where a significantly
higher overpotential is applied to ensure that the entire substrate is covered with
metal before the potential is reduced to the plating potential. This has been shown
to be effective in ionic liquid and significantly improves the corrosion resistance of
the coatings [7].
292 11 Technical Aspects
Fig. 11.3 AFM image of copper etched for 20 s at 10 V in a Type III
eutectic of 1 choline chloride: 2 ethylene glycol (right side of image
masked during experiment).
One issue that has to be addressed is the reversibility of the dissolution and
deposition of the substrate. If the dissolution of the substrate is reversible, i.e. all
the metal dissolved can be redeposited, then etching in situ in the plating liquid
is possible. If the substrate cannot be redeposited then the metal will clearly build
up its concentration as the ionic liquid is used and this will significantly shorten
Fig. 11.4 Line traces taken through Figures 11.2 and 11.3; aluminum
(dotted line) and copper (solid line).
11.1 Metal Dissolution Processes/Counter Electrode Reactions 293
the life of the bath. In this case a pre-etch in a different liquid should take place
before the substrate is transferred to the ionic liquid. This is shown schematically
in Figure 11.1.
No literature has been published in this area but, as a rule of thumb, metals
which dissolve to give complexes that have linear or tetrahedral geometries, e.g.
Cu, Ag, Zn, Sn, Pb, can be reversibly deposited and etched. Those with octahedral
geometries, e.g. Fe, Ni, Co and Cr, are less reversible. The exceptions to this are the
very electronegative metals, most notably Al which is difficult to electrodeposit from
some ionic liquids. The reversibility is also dependent upon the type of ionic liquid
and the metal being deposited. Endres has shown that the adhesion of aluminum
to mild steel is greatly enhanced by an anodic pulse prior to deposition. It has been
shown that this alloy was formed between the steel substrate and the aluminum
coating [1].
11.1.3
Electropolishing of Stainless Steels
Electropolishing is the controlled corrosion of a metal surface to bring about a
reduction in surface roughness and an increase in corrosion resistance of the
components. Electropolished pieces also decrease wear and increase lubricity in
engines, thus reducing a major cause of failure, and offer several other functional
benefits. The first systematic study of electropolishing was carried out by Jacquet
and led to a patent in 1930 [13]. The majority of studies have been carried out on
stainless steel although metals such as copper, nickel and titanium have also been
studied [14–16]. The current stainless steel electropolishing process is performed
worldwide on a commercial scale and is based on concentrated phosphoric acid and
sulfuric acid mixtures. The polishing process is thought to involve the formation of
a viscous layer at the metal surface and many processes employ viscosity improvers
such as glycerol. The practical aspects of electropolishing have been reviewed by
Mohan et al. [17], whereas the more fundamental aspects are covered in a review
by Landolt [18]. While electropolishing is an extremely successful process there are
major issues associated with the technology, most notably that the solution used is
highly corrosive and extensive gassing occurs during the process, which results in
very poor current efficiency.
As explained previously, electrodissolution in ionic liquids is a simple and efficient process, particularly in chloride-based eutectics. Type III eutectics based
on hydrogen bond donors are particularly suitable for this purpose. However, it
has been noted that the polishing process only occurs in very specific liquids and
even structurally related compounds are often not effective. It has been shown
that 316 series stainless steels can be electropolished in choline chloride: ethylene
glycol eutectics [19] and extensive electrochemical studies have been carried out.
The dissolution process in aqueous solutions has been described by two main
models; the duplex salt model, which describes a compact and porous layer at the
iron surface [20], and an adsorbate–acceptor mechanism, which looks at the role
of adsorbed metallic species and the transport of the acceptor which solubilises
294 11 Technical Aspects
them [21]. Voltammetry and impedance spectroscopy have been used to confirm
that the dissolution mechanism in an ionic liquid is different from that in aqueous
acidic solutions. Preliminary results suggest that a diffusion-limited process in the
viscous ionic liquid appears to be responsible for electropolishing [22]. Impedance
spectroscopy has also shown that one of the main differences between the electropolishing mechanism in the ionic liquid and the aqueous solution is the rate
at which the oxide is removed from the electrode surface. The electropolishing
mechanism in the ethylene glycol eutectic is described in more detail in two recent
publications [21, 22].
Highly polished surfaces were obtained with current densities between ca. 70
and 50 mA cm−2 with an applied voltage of 8 V. Below this current density a milky
surface was obtained and above this range some pitting was observed on an otherwise bright surface. It should be noted that the polishing region was narrower
than that in aqueous phosphoric/ sulfuric acid mixtures, but the current density
requirements were considerably lower using the ionic liquid. In acidic solutions
typical current densities are 100 mA cm−2 but much of this results in gas evolution
at the anode. With the ionic liquid no gas evolution was observed, suggesting that
there are negligible side reactions occurring with the ionic liquid. The current efficiency of the 1ChCl: 2EG electrolyte has been determined using coulometry and
gravimetry and was found to be in excess of 90%, which is significantly higher than
the aqueous-based electrolytes which are typically ∼30%. Given that the current
density used for the 1ChCl: 2EG electrolyte is considerably lower than that used in
the aqueous solution the slight difference in the conductivity of the two solutions
does not lead to a significant difference in the ohmic loss through the solution.
In fact preventing a passivating layer at the electrode surface during polishing decreases the overall ohmic resistance of the non-aqueous system. Hence the current
going to metal dissolution is probably similar for the two systems, explaining why
the polishing process takes approximately the same time.
Analysis of the polished surface and residue left in the polishing tank showed
conclusively that no dealloying of the surface took place and AFM analysis of preand post-polished samples showed that the polishing process was effective at preventing corrosion because it removed the micro-cracks from the steel surface [22].
Various pre-treatment protocols have been developed including pickling and
anodic/cathodic pulses to remove the oxide films. It was apparent that different
types of steel require different pre-treatments, i.e. cast pieces behave differently to
rolled pieces. Significant success was achieved in electropolishing cast pieces and
the finish obtained with the ionic liquid was superior to that with phosphoric acid,
however, the converse was true for rolled pieces because the oxide film is thicker
in the latter samples and hence slower to dissolve in the ionic liquid.
Similar electropolishing experiments were carried out using different grades of
stainless steel (410, 302, 304, 316 or 347) and it was found that the mechanism of
metal dissolution and the oxidation potentials for the metals were very similar. The
slight exception was the 410 series steel (which has no Ni, unlike the 300 series
steels which have 8–14%). The 410 steel required a more positive oxidation potential
to break down the oxide in the ionic liquid whereas once the oxide was removed the
11.1 Metal Dissolution Processes/Counter Electrode Reactions 295
Fig. 11.5 AFM image of a 316 stainless steel sample in which one
side has been electropolished while the other has been masked with
lacquer.
metal was more easily oxidised than the other grades of steel. This shows why the
410 steel was more likely to pit during the polishing process. The pitting could be
reduced, however, by chemically pickling the steel with a proprietary phosphoric
acid etch before electropolishing [23].
This technology was scaled-up to a 1.3 tonne plant by Anopol Ltd (Birmingham,
UK). Results have shown that the technology can be applied in a similar manner to
the existing technology. The ionic liquid has been found to be compatible with most
of the materials used in current electropolishing equipment, i.e. polypropylene,
nylon tank and fittings, stainless steel cathode sheets and a titanium anode jig.
Extended electropolishing using the same solution leads to a dark green–brown
solution arising from the dissolved iron, chromium and nickel. The solubility of
the metals in the ionic liquid is relatively high and a dense sludge forms in the base
of the tank when the saturation concentration is exceeded. The metals are present
as glycolate and chloride complexes and numerous solvents have been tested to
determine their efficacy at precipitating the metal salts. (See SubChapter 11.3).
Water is completely miscible with the spent ionic liquid but the resulting mixture leads to a completely transparent liquid and almost all the metal complex is
precipitated to the base of the cell. The water can be distilled from the mixture
to leave a dry ionic liquid which has lost only ca. 15% ethylene glycol, mostly in
the form of the metal complex. The residual concentration of each metal in the
ionic liquids was less than 5 ppm. Hence, not only has it been demonstrated that
electropolishing can be carried out in this non-corrosive liquid, but also that the
liquid can be completely recycled and all of the metal can be recovered. Figure 11.6
shows a variety of stainless steel pieces electropolished using the choline-based
ionic liquid.
296 11 Technical Aspects
Fig. 11.6 A variety of pieces electropolished using a choline-based ionic liquid.
This subchapter has shown that metal dissolution processes are important to
numerous aspects of metal plating, however, very few concerted studies have been
made in this area. An understanding of dissolution rates and processes, together
with information on the stability of oxide films in ionic liquids, is essential for the
development of successful metal finishing processes.
11.2
Reference Electrodes for Use in Room-temperature Ionic Liquids
Voltammetric, electrodeposition, electrosynthetic and electroanalytical studies are
carried out in room-temperature ionic liquids (RTILs) by a significant and increasing number of both industrial and academic laboratories [23–25]. Such studies,
when carried out at anything other than a very empirical level, require the use of a
‘reference electrode’. The purpose of this chapter is to address the special problems
this poses and their solutions. First, however, we start by considering the essential
features of a reference electrode in general.
11.2.1
What is a Reference Electrode?
Consider a generalized electrode process: A ± ne → B, in which A is electrolytically
converted into B at a suitable electrode. The rate at which this happens is measured
by the current, I, flowing through the electrode, via. Faraday’s Law:
I = nF Aℑ
(11.2)
11.2 Reference Electrodes for Use in Room-temperature Ionic Liquids 297
where F is the Faraday constant (96 485 C mol−1 ), A is the electrode area, and ℑ is
the flux of species A to the electrode (mol cm−2 s−1 ) averaged over its surface. This
rate of reaction is controlled by the magnitude of the electrical potential applied to
the electrode of interest (often referred to as a “working” electrode). The application
of this potential, self-evidently, requires the presence of at least a second electrode
in the solution of interest, so that a defined potential can be applied between the
two electrodes. This second electrode is known as a “reference” electrode. In order
that a fixed and known driving potential is applied to the working electrode, it is
a minimum requirement that the reference electrode maintains a fixed, constant
potential difference between itself (M) and the electrolytic solution (S) with which
it is in contact, (M – S )ref [26].
The potential difference (M − S )ref is established by means of a suitable electrochemical equilibrium being established at the surface of the reference electrode:
Ox +e ↔ Red, so that the reference potential difference of interest is quantified by
means of the Nernst equation [27]:
(M − S )ref = a constant −
RT
ared
ln
F
aox
(11.3)
where R is the universal gas constant, T is the temperature in Kelvin, and ared and
aox are the activities of the reduced and oxidized species, respectively.
Examples of reference electrode systems which operate successfully in aqueous
solutions include the following ‘potential determining equilibria’:
1
e− + H+ ⇋ H2
2
1
e− + Hg2 Cl2 ⇋ Hg + Cl−
2
e− + AgCl ⇋ Ag + Cl−
(11.4)
(11.5)
(11.6)
In the first example, a platinized platinum electrode is immersed in a solution of
strong acid. The purpose of the platinizing procedure is to ensure the kinetics of the
electrochemical processes are rapid enough to sustain the process at equilibrium.
In the other two examples, a metal salt, highly insoluble in water (AgCl or Hg2 Cl2 )
is in contact with a solution containing chloride ions and the corresponding metal
(Ag or Hg). The corresponding appropriate forms of the Nernst equations are:
pH1/2
RT
2
ln
F
a H+
RT
(M − S )ref = A′ −
ln aCl−
F
RT
(M − S )ref = A′′ −
ln aCl−
F
(M − S )ref = A −
(11.7)
(11.8)
(11.9)
In experimental practice, the reference electrode will most likely be used in conjunction with a three-electrode potentiostat with a third electrode, a counter (or
298 11 Technical Aspects
auxiliary) electrode, completing the circuit so that negligible currents pass through
the reference electrode. The latter feature is of crucial importance otherwise electrolysis perturbs the concentrations (‘activities’) of the desired species establishing
the potential determining equilibrium and hence the quantity (M – S )ref which,
under conditions of sustained current flow, is neither fixed nor constant [28]. That
said, in the limit of microelectrodes, the currents passed can be sufficiently small
that a two-electrode system becomes viable. In this arrangement, a single electrode
can act as a reference electrode and a counter electrode.
Classically, in relation to conventional solvent media, three classes of reference
electrodes are recognised [29]:
1. Electrodes of the first kind: These are based on a potential determining equilibrium such as Ag+ + e− ⇋Ag or 12 Cl2 + e− ⇋Cl− where, for “cationic electrodes”,
equilibrium is established between atoms or molecules and their corresponding
cations in solution or, for “anionic electrodes”, their corresponding anions.
2. Electrodes of the second kind: These consist of three phases. A metal is covered by
a layer of its sparingly soluble salt, and immersed in a solution containing the
anion of this salt. The Ag/AgCl/Cl− and Hg/Hg2 Cl2 /Cl− electrodes referred to
above are of this type.
3. Redox electrodes: In this case, an inert, non-reactive metal such as platinum or gold,
is immersed in a solution containing both species contributing to a redox couple.
For example in water: 12 BQ + e− + H+ ⇋ 21 H2 Q, where BQ is benzoquinone and
H2 Q is hydroquinone or, in acetonitrile, Cp2 Fe+ + e− ⇋Cp2 Fe, where Cp2 Fe is
ferrocene and Cp2 Fe+ the ferrocenium cation.
11.2.2
Essential Characteristics of a Reference Electrode
In the context of classical solvent media, Butler [30] suggests:
“A satisfactory reference electrode must show one or more of the following properties: (1)
have a potential stable with time, (2) return to the same potential after polarization,1)
(3) obey the Nernst equation with respect to some species in the electrolyte, and (4) if it
is an electrode of the second kind, the solid phase must not be appreciably soluble in the
electrolyte.”
In the context of RTILs the criterion (3) raises considerable problems since the concept of activity and activity coefficients of ions is largely unexplored in such media.
Accordingly, validation of the applicability of the Nernst equation in such media is a
non-simple exercise, given that RTILs are likely to exhibit gross non-ideality. Rather,
electrochemical measurements based on otherwise validated reference electrodes,
may likely in the future provide a methodology for the study of RTIL non-ideality.
1) By ‘polarization’ is meant the application of a
voltage perturbing the equilibrium potential of
the electrode.
11.2 Reference Electrodes for Use in Room-temperature Ionic Liquids 299
Accordingly, for our present purposes, namely the identification of satisfactory reference electrodes, the pragmatic criteria of (1), (2), and (4) are pertinent, and (1)
in particular is paramount since, in essence, (2) and (4) merely indicate means by
which (1) might fail. Underpinning the requirement for a stable electrode potential
is, of course, the need for relatively fast electrode kinetics to establish the potential
determining equilibrium. To quote Ives and Janz [31]:
“Exchange current densities for various kinds of metal–solution interfaces cover a range of
about 10 −2 to 10 −18 A cm−2 , but the useful range for reference electrodes is normally much
more restricted than this; it will be in part dependant upon the sensitivity of the measuring
instrument to be used. One of the highest i0 values is for hydrogen ion discharge at platinum,
which is one reason why the hydrogen electrode2) is one of the most satisfactory of all.”
11.2.3
Pseudo-reference Electrodes and Internal Redox Reference Couples
Butler [30] says:
“If one is not too critical, many metal electrodes show relatively stable potentials in various
electrolyte solutions.”
Accordingly, much voltammetry in non-aqueous solvents has been conducted using
a ‘pseudo’-reference electrode (alternatively labelled a ‘quasi’-reference electrode)
constituting, quite simply, a metal wire, most often silver or platinum. It is then
expected (hoped) that the potential of the wire remains constant throughout the
voltammetric experiment. This may be a realistic hope if, as Bard and Faulkner
[32] point out, the composition of the bulk solution is essentially constant during
the period of experimentation, as may be realized during voltammetric studies but
certainly not in electrosynthetic work.
When a pseudo-reference electrode is used, good practice [32] dictates that its
actual potential is calibrated by measuring, voltammetrically or otherwise, the formal potential of an electrochemically reversible couple. IUPAC recommend the
use of either the ferrocene/ferrocenium, Cp2 Fe/Cp2 Fe+ couple [33], alternatively,
the cobaltocenium/cobaltocene Cp2 Co+ /Cp2 Co (where Cp ≡ C5 H5 ) has been suggested [34, 35]. In experimental practice, this simply involves measuring the voltammogram of either Cp2 Fe or Cp2 Co+ using the selected metal wire as the pseudoreference electrode before (and after) recording that of the species of interest in the
same medium. Since the couples Cp2 Fe+ + e− ⇋Cp2 Fe and Cp2 Co+ + e− ⇋Cp2 Co
are electrochemically reversible in most media and at most electrodes3) , comparison of the measurements allows redox data to be reported against either of the two
2) In aqueous solution.
3) Note that couples which show electrochemically reversible behavior at macro-electrodes
may display quasi-reversibility or irreversibility
at very small electrodes (ultramicro-electrodes,
nano-electrodes)
300 11 Technical Aspects
couples. The possibility of using solid Cp2 Fe+ + e− ⇋Cp2 Fe in the specific context
of RTIL voltammetry has been noted by Zhang and Bond [36, 37].
11.2.4
Liquid Junction Potentials
Measurements of electrode potentials using reference electrodes are of two general
types: those that involve liquid junctions and those that do not. An example of a
cell which does not have a liquid junction is:
Pt | H2 (g) | HCl(aq) | AgCl | Ag
where | denotes a phase boundary. In contrast the cell
Pt | Cp2 Fe, Cp2 Fe+ (CH3 CN) || AgNO3 (CH3 CN) | Ag
has a liquid junction ( ) since two liquid phases of different compositions are
brought into contact. The liquid phases may differ in terms of solvents and/or
solutes.
When liquid junctions exist, liquid junction potentials (LJPs) can arise due to
differing ion mobilities across the interface, leading to charge separation and the
development of a potential difference across the liquid junction. These can amount
to some tens of millivolts and add a corresponding uncertainty in any voltammetric
measurement. It follows that systems that avoid LJPs are generally preferable;
otherwise some consideration of their likely magnitude is desirable (see below).
11.2.5
Reference electrodes in RTILs: What has been used?
Table 11.1 presents the results of a literature survey to establish which reference and
pseudo-reference electrodes have been and are being used in RTILs. The structures
of the constituent anions and cations are shown in Figure 11.7. It is clear that the
majority of researchers favor the use of pseudo-reference electrodes but that not all
take the trouble to calibrate using internal standards such as Cp2 Co+ or Cp2 Fe. In
the latter case, the philosophy is nicely and honestly summarized by Welton and
colleagues [38]:
“The electrochemistry was performed on the neat ionic liquid. In such a set-up, with
no recognised background electrolyte or redox standard, the potential vs. the platinum
pseudo-reference is difficult to compare with standard potentials, however, in such unusual
conditions it is the qualitative nature of the electrochemistry that is important.”
The most popular pseudo-reference electrodes are Pt or Ag wires. Other pseudoreference electrodes have employed coating, for example Pt with polypyrrole [39]
or Ag with AgCl [40] (but in the absence of deliberately added solution phase Cl− ,
11.2 Reference Electrodes for Use in Room-temperature Ionic Liquids 301
Table 11.1 A cross-section of the different types of reference electrodes
that have been used by various researchers in a range of different
RTILs
RTIL solution
Reference electrode material
Referenced to . . .
[C4 mim][PF6 ]
[C2 mim][NTf2 ]
[M(MePEG-bpy)2+ ][DNA][b]
Various [NTf2 ]
[C4 mim]Cl/AlCl3
DIMCARB
[C4 mPyrr][NTf2 ]
[N4,1,1,1 ][NTf2 ]
[N6,2,2,2 ][NTf2 ]
[C4 mim][BF4 ]
[C4 mim][PF6 ]
[C4 dmim][BF4 ]
[C4 mim][PF6 ]
[C4 mPyrr][NTf2 ]
[C4 mim][BF4 ]
[C4 mim][PF6 ]
[C4 mim][NTf2 ]
[C4 mPyrr][NTf2 ]
[C4 mim][Co(CO)4 ]
[C2 mim][NTf2 ]
[C4 mim][NTf2 ]
[C4 mim][PF6 ]
[N8,8,8,1 ][NTf2 ]
[C2 mim][BF4 ]
Ag wire
Cc+ /Cc[a]
Fc/Fc+[a]
Fc/Fc+[a]
NR
NR
DMFc/DMFc+[a]
NR
[C2 mim]Cl/AlCl3
[PP13 ][NTf2 ]
[C6 dmim][NTf2 ]
[C6 dmim][CTf3 ]
[C6 dmim][PF6 ]
[C6 dmim][AsF6 ]
[C4 mim][PF6 ]
[C2 mim][BF4 ]
[C3 mim][BF4 ]
[C4 mim][BF4 ]
[C2 mim][BF4 ]
[C4 mim][BF4 ]
[C4 dmim][BF4 ]
[C4 mPyrr][NTf2 ]
[C2 mPip][F(HF)2 ]
[C4 mPip][F(HF)2 ]
[C4 mPyrr][F(HF)2 ]
[C2 mim][F(HF)2 ]
Ag wire
Ag wire
Ag wire
Ag wire
Ag wire
Ref.
[37]
[52]
[23, 25]
[53]
[36]
[54]
Pt wire
Fc/Fc+[a]
(E o = 0.3, 0.39 and
0.49 V/ Ag/AgCl)
Fc/Fc+
[24, 55]
Pt wire
NR
[45]
Pt wire
Fc/Fc+[a]
[56]
Pt wire
Pt wire coated in polypyrrole
NR
Fc/Fc+[a]
(E o = 0.405 V/ SCE)
[38]
[39]
Al wire immersed in a 1.5:1.0
acidic chloroaluminate
melt (frit)
Al wire in an 0.6 M solution
of RTIL [C2 mim]Cl/AlCl3
(porous tip)
Mg ribbon in Mg(CF3 SO3 )2
Li foil in Li+ salts
n/a
[57]
n/a
[41]
n/a
n/a
[47]
[48]
Saturated calomel (aq)
Ag/AgCl KCl (sat., aq)
n/a
n/a
[43]
[44]
Ag/AgCl Na (sat., aq)
n/a
[42]
Ag wire in 0.1 M AgNO3 in
RTIL [C4 mim][NO3 ] (glass
frit)
Ag wire in 0.05 M AgBF4 in
RTIL [C2 mim][BF4 ]
n/a
[58]
n/a
[50]
Ag wire coated with AgCl
[40]
302 11 Technical Aspects
Table 11.1 (Continued)
RTIL solution
Reference electrode material
Referenced to . . .
Ref.
[C4 mim][NTf2 ]
[C4 mim][BF4 ]
[C4 dmim][PF6 ]
[C4 mPyrr][NTf2 ]
Ag wire in 0.1 M AgNO3 in
RTIL and Ag/AgCl in
0.1 M Bu4 NCl in RTIL.
Ag wire in 0.01 and 0.1 M
AgTf in RTIL (frit)
Pt wire in 0.06 M
N(n-C3 H7 )4I, 0.015 M I2 in
[C2 mim][NTf2 ]
n/a
[49]
n/a
[51]
n/a
[46]
[N3,1,1,1 ][NTf2 ]
[TES][NTf2 ]
[TBS][NTf2 ]
a
Cc+ =Cp2 Co+ , Cc=Cp2 Co, Fc=Cp2 Fe, Fc+ =Cp2 Fe+ , DMFc=(C5 Me5 )2 Fe, DMFc+ =(C5 Me5 )2 Fe+ .
M= Fe, Co and MePEG-bpy = 4,4’-(CH3 (OCH2 -CH)OCO)-2,2’-bipyridine).
NR = no calibration vs. internal reference reported. All RTIL structures are given in Figure 11.7.
b
although some may arise locally from dissolution of AgCl). In both of these cases,
the Cp2 Fe/Cp2 Fe+ couple was used as an internal reference for the purposes of
calibration.
Another type of apparently “pseudo”-reference electrode involves the use of Al
wires in contact with solutions containing AlCl4 − ions [41]. A further group of
researchers simply use conventional aqueous solution-based calomel or silver/silver
chloride/aqueous chloride ion reference electrodes [42–44]. These are included in
Table 11.1 for illustration and completeness. The use of such electrodes is highly
likely to lead to the introduction of water into the RTIL system in contact with the
reference electrode, as well as to unknown problems in respect of LJPs. Properties
such as voltammetric windows, diffusion coefficients and RTIL viscosity are all
likely to be highly sensitive to trace amounts of water [45].
The following systems, in contrast to the above, are based on well-defined
potential-determining equilibria established within a RTIL.
3 −
−
1. The iodide/tri-iodide system: 12 I−
3 + e ⇋ 2 I has been used by Matsumoto et al.
[46]. The electrode was prepared by dissolving 60 mM N(n-C3 H7 )I and
15 mM I2 in 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
[C2 mim][NTf2 ] and placing a platinum wire in the solution. It is highly likely,
but not explicitly reported, that such a reference electrode was used to study the
voltammetry in various RTILs based on triallylsulfonium cations. If so, there
would be an unknown, but probably not too large and reasonably constant, liquid junction potential between the RTIL under study and the reference electrode
cell.
2. The couple 12 Mg2+ + e− ⇋ 12 Mg, with the cation present as the salt Mg(CF3 SO3 )2
(1 M) has been used as a reference electrode in N-propyl-N-methylpiperidinium
bis(trifluoromethylsulfonyl)imide [C3 mPip][NTf2 ] [47]. The authors also considered the use of a magnesium ribbon as a pseudo-reference electrode. The RTIL
11.2 Reference Electrodes for Use in Room-temperature Ionic Liquids 303
Fig. 11.7 Structures of all RTILs listed in Table 11.1.
304 11 Technical Aspects
in the reference electrode and in the bulk solution used for voltammetry were
the same so that liquid junction potentials were relatively minimized.
3. The couple Li+ + e− ⇋Li has been used for RTILs based on [1, 2-dimethyl-3propylimidazolium] [X]− (where [X]− = [NTf2 ]− , [CTf3 ]− , [PF6 ]− and [AsF6 ]− )
[48]. The electrode took the form of Li foil in the same ionic liquids to which was
added 0.02 M LiAsF6 , LiPF6 , Li[NTf2 ] or Li[CTf3 ] according to the nature of the
anion in the RTIL of interest. Again, this arrangement led to a minimization of
liquid junction potentials.
4. Josowicz et al. [49] have developed a reference electrode for use in RTILs based
on the equilibrium AgCl + e− ⇋Cl− + Ag, in which a chlorinated silver wire is
placed in a solution of 0.1 M Bu4 N+ Cl− in the RTIL of interest. The latter solution
was separated from the sample under study by a double junction arrangement
in which a further compartment contained only the RTIL of interest.
5. Several researchers have used the following potential determining equilibrium:
Ag+ + e− ⇋Ag
as the basis for well-defined reference electrodes. Josowicz et al. [49] dissolved
0.1 M AgNO3 in the RTIL of interest and inserted a silver wire. This was used in
a similar double junction arrangement as described above. Hagiwara et al. [50]
used 0.05 M AgBF4 in the ionic liquid 1-ethyl-3-methylimidazolium tetrafluoroborate [C2 mim][BF4 ]. Finally, Snook and colleagues [51] devised and voltammetrically characterized a Ag/Ag+ reference electrode which incorporated a known
concentration (usually 10 mM) of silver trifluoromethanesulfonate (AgTf; Tf =
CF3 SO3 − ) in 1-butyl-1-methylpyrroloidinium bis(trifluoromethylsulfonyl)imide
[C4 mPyrr][NTf2 ]. A stable and reproducible potential was reported. In a careful
and thorough study, the electrode Ag/Ag+ (10 mM AgTf, [C4 mPyrr][NTf2 ]) was
found to be stable to within a millivolt over a period of around three weeks, when
used in an argon atmosphere at room temperature. This is a highly important
and useful observation since the characterization of the RTIL-based reference
electrode (see Section 11.2.2) was significantly more rigorous than in any other
study of which the present authors are aware. Specifically, for a high concentration of Ag+ , close to Nernstian behavior was seen and measurements showed the
electrode to be significantly more stable than a Ag pseudo-reference electrode,
even when the latter was separated by a salt bridge. Above all, voltammetric
data recorded in a range of ionic liquids against the Ag/Ag+ (10 mM AgTf,
[C4 mPyrr][NTf2 ]) reference electrode showed apparent liquid junction potentials
of no more than a few tens of millivolts.
11.2.6
Recommendations and Comments
It is evident from the previous section that a range of approaches have been, and can
be, adopted by experimentalists wishing to conduct voltammetric or other studies.
The aim of this section is to answer some likely questions.
11.2 Reference Electrodes for Use in Room-temperature Ionic Liquids 305
11.2.6.1 When and How Can I Use a Pseudo-reference Electrode in Voltammetry?
The use of a Pt or Ag wire as a pseudo-reference electrode is attractive because
of its sheer simplicity and the fact that possible contamination of the test solution
is avoided. The issue as to whether this provides a stable reference potential is an
important consideration. To illustrate this, the electrochemistry of 5 mM Ferrocene
(Cp2 Fe) in the RTIL [C4 mim][PF6 ] was studied on a platinum microdisk electrode
(d = 10 µm). The pseudo-reference electrode used in this set-up was simply a
platinum wire inserted into a glass tube (standard non-aqueous reference electrode
kit from BAS) in the same RTIL [C4 mim][PF6 ], separated from the main solution
via a Vycor plug. The reference electrode was, prior to recording the voltammetry
of Cp2 Fe, pre-oxidized for different times by holding the potential at ca. +1.75 V in
blank [C4 mim][PF6 ] vs. a silver wire pseudo-reference. It may be possible that the
pre-oxidizing experiment deposits a layer of some species on the Pt wire, leading to
significant shifts in potential. Figure 11.8 shows this effect: with no pre-oxidizing
(a), the half-wave potential of Cp2 Fe is +0.275 V, which systematically shifts to
more negative potentials with increased pre-oxidizing time (+0.255 V for 5 min (b),
+0.225 V for 10 min (c), +0.185 V for 20 min (d), and +0.165 V for 40 min (e)).
The Pt wire that had been pre-oxidized for 20 min was then left in air for a further
1 h, after which the half-wave potential of Cp2 Fe had shifted back to a potential
(+0.245 V (f)) close to that observed with no pre-oxidizing. The same experiments
were repeated with a silver wire inside the reference compartment, and the results
are shown in Figure 11.9. Here, although there was no systematic shift in peak
potentials with pre-oxidizing time: (+0.385 V for 0 min (a), +0.365 V for 5 min (b),
+0.415 V for 10 min (c), and +0.385 V for 20 min), there was still a significant
difference in the half-wave potential of Cp2 Fe under different conditions.
It is clear that Pt or Ag wires can show significant drift (which depends in part
on their recent history as well as the solution in which they are immersed) and that
if such pseudo-reference electrodes are used, the regular internal calibration using
Cp2 Co+ or Cp2 Fe, as advocated by IUPAC [33–35] and Zhang and Bond [37], is
essential if anything other than the qualitative data is sought. Other redox couples
with Nernstian characteristics may also be suitable. Examples might include:
(i) The benzoquinone/benzoquinone radical anion couple (BQ/BQ•− ):
(ii) the N,N,N ′ ,N ′ -tetramethylphenylenediamine radical cation / N,N,N ′ ,N ′ tetramethylphenylenediamine couple (TMPD•+ /TMPD):
306 11 Technical Aspects
Fig. 11.8 Cyclic voltammograms for the oxidation of 5 mM ferrocene in [C4 mim][PF6 ]
on a platinum microelectrode (diameter
10 µm) at 100 mV s−1 . Reference electrode
was a Pt wire inserted into [C4 mim][PF6 ]
contained in a glass tube, separated by a
Vycor frit. Pre-oxidation of the reference
electrode took place for (a) 0 min, (b)
5 min, (c) 10 min, (d) 20 min, (e) 40 min
and (f) 20 min with 1 h ‘rest’.
Figures 11.10 (a) and (b) show that the voltammetry of these couples in a range
of RTILs is nearly electrochemically reversible. Note however that, unlike the
ferrocene- and cobaltocenium-based couples, the reduction potentials are likely
to vary significantly from one RTIL to another. In experimental practice it is also
important to verify that the calibration molecules do not interfere chemically with
the voltammetric process under study. For example, we have investigated the oxidation of molecular hydrogen in the presence of TMPD and observed a reaction of
the two species, as noted by the disappearance of the reverse-peak of the first redox
couple (see Figure 11.11). This implies that the peak potentials of TMPD•+ /TMPD
are no longer obvious, and that this redox couple cannot be used as an internal
reference in this type of experiment.
11.2 Reference Electrodes for Use in Room-temperature Ionic Liquids 307
Fig. 11.9 Cyclic voltammograms for the oxidation of 5 mM ferrocene in [C4 mim][PF6 ] on a platinum microelectrode (diameter
10 µm) at 100 mV s−1 . Reference electrode was a Ag wire inserted into
[C4 mim][PF6 ] contained in a glass tube, separated by a Vycor frit. Preoxidation took place for (a) 0 min, (b) 5 min, (c) 10 min, (d) 20 min.
11.2.6.2 How Do I Conduct an Electrosynthetic Experiment under Potential Control?
In this case, since the aim of the experiment is the bulk concentration of the
material being electrolyzed, then any attempts to maintain a fixed potential using
a pseudo-reference electrode will likely be hopeless. A properly defined and wellcharacterized electrode is essential and the present authors consider that described
by Snook et al. [51] to be very probably the best currently available. Note that
Fig. 11.10 Cyclic voltammograms for (a) the reduction of 12.5 mM
benzoquinone (BQ) in [C4 mim][NTf2 ] on a platinum microelectrode
(diameter 10 µm) at 100 mV s−1 and (b) the oxidation of 20 mM
N,N,N′ ,N′ -tetramethylphenylenediamine (TMPD) in [C4 dmim][NTf2 ]
on a platinum electrode (diameter 10 µm) at 4 V s−1 .
308 11 Technical Aspects
Fig. 11.11 Cyclic voltammetry of 20 mM TMPD in [C4 dmim][NTf2 ] on
a platinum electrode (diameter 10 µm) at 100 mV s−1 in the presence
of 0% and 100% hydrogen.
the electrode can be constructed in the form of a separate probe as shown in
Figure 11.12.
11.2.6.3 What Options Are Available for Rigorous, Quantitative Voltammetry?
For most voltammetric purposes, the Ag/Ag+ (10 mM AgTf, [C4 mPyrr][NTf2 ]) electrode discussed above can be recommended as a general, stable and well characterized reference electrode although issues of the possible photo-instability of AgTf
in daylight may need to be addressed in some applications. If an electrode of this
type is introduced into a RTIL other than [C4 mPyrr][NTf2 ], the most likely source
of error will occur from liquid junction potentials at the [C4 mPyrr][NTf2 ]/RTIL
interface. As Snook and co-workers [51] point out, these may amount to a few tens
of millivolts, but probably no more.
It follows that for the more rigorous work, it is worth developing reference
electrodes which minimize the liquid junction potentials. This is probably best
achieved by using the RTIL under study as the solvent in the reference system.
Building on published experiments, the latter is probably most securely based on
the Ag/Ag+ system. Thus, for example, in an RTIL in which the anion is [BF4 ]− ,
Fig. 11.12 Outline of components of Ag/Ag+ reference electrode, and
the reference electrode inserted into a salt bridge compartment.
11.2 Reference Electrodes for Use in Room-temperature Ionic Liquids 309
Fig. 11.13 (a) Cyclic voltammograms
for the reduction of 84 mM AgTf in
[C4 mPyrr][NTf2 ] on a platinum microelectrode (diameter 10 µm) at scan rates
of 200, 400, 700 mV s−1 , 1, 2, 4, 7 and
10 V s−1 . The pseudo-reference electrode
used was a silver wire. (b) Cyclic voltammetry for the reduction of 84 mM AgTf in
[C4 mPyrr][NTf2 ] on silver wire (diameter
0.5 mm) at 10 mV s−1 with Ag/AgNO3 reference electrode as in Ref. [58].
the Ag+ could most beneficially be introduced as AgBF4 (as in Ref. [50]). Similarly,
in [NO3 ]− -based RTILs, AgNO3 might be a recommended source of Ag+ . We note
that the following Ag salts (with anions corresponding to common RTIL anions)
are commercially available from Aldrich: AgTf (silver trifluoromethanesulfonate),
AgNTf2 (silver tri(fluoromethylsulfonyl)-imide), AgBF4 (silver tetrafluoroborate),
AgPF6 (silver hexafluorophosphate), AgNO3 (silver nitrate), AgCl (silver chloride),
AgMeSO4 (silver methanesulfonate), AgSCN (silver thiocyanate), AgHF2 (silver
hydrogenfluoride), AgAc (silver acetate), AgTFA (silver trifluoroacetate).
In Cl− -based RTILs the Ag/AgCl/Cl− system can be used to generate a reference
electrode relatively free of liquid junction potentials [49]. Finally, in generating
new reference electrodes based on the Ag/Ag+ system, it is worthwhile pointing
out that the electrode kinetics of this system are certainly unexplored in almost
any RTIL medium. Prudence dictates that some brief study of the aspect precedes
any application of newly developed reference systems. Usually, a voltammogram
(recorded against a pseudo-reference electrode!) will suffice to show that the Ag/Ag+
couple does or does not possess sufficiently fast (“Nernstian”, “reversible”) electrode
kinetics. Figure 11.13(a) illustrates the concept in respect of Ag metal deposited
on a Pt microelectrode (d = 10 µm) from AgTf in [C4 mPyrr][NTf2 ]. The relative
closeness of the peaks suggests quasi-reversible electrode kinetics and hence that a
likely satisfactory reference electrode system can be based on Ag/AgTf in the RTIL
of interest. In essence, this approach is equivalent to the micro-polarization test
advocated by Ives and Janz in their classic text [31]. Figure 11.13(b) shows similar
data to Figure 11.13(a), except performed using a Ag wire electrode; the near lack of
hysteresis confirms the near electrochemical reversibility of the system and hence
the validation of the system as the basis of a reference electrode, as advocated by
Snook et al [51].
310 11 Technical Aspects
11.3
Process Scale Up
11.3.1
Introduction
Although the deposition of metals from ionic liquids has been possible for over
50 years, to date no processes have been developed to a commercial scale. There
are numerous technical and economic reasons for this, many of which will be
apparent from the preceding chapters. Notwithstanding, the tantalizing prospect
of wide potential windows, high solubility of metal salts, avoidance of water and
metal/water chemistry and high conductivity compared to non-aqueous solvents
means that, for some metal deposition processes, ionic liquids must be a viable
proposition.
To assess the issues that need to be addressed before commercialization can
be implemented it is probably easier to analyse the current and future markets
for electroplating and compare the limitations of the current technology for various metals. The main metals of interest are Cr, Ni, Cu, Au, Ag, Zn and Cd,
together with a number of copper and zinc-based alloys. The electroplating industry, which dates back well over 100 years, is based, naturally, on aqueous solutions
due to the high solubility of electrolytes and metal salts resulting in highly conducting solutions. Water does, however, suffer from the drawback that it has a
relatively narrow potential window and hence the deposition of electronegative
metals such as Cr and Zn is hindered by poor current efficiencies and hydrogen
embrittlement of the substrate. In addition there are specific difficulties with certain
metals.
The most obvious case is that of chromium plating. The major disadvantage
of the current process of chrome plating is that it requires the use of chromic
acid-based electrolytes comprising hexavalent chromium, Cr(VI). The toxicity and
carcinogeneity associated with Cr(VI) [59] has resulted in wide-ranging environmental legislation in the USA (OSHA, EPA) and Europe (IPPC) to reduce its use.
For example, the EU End-of-Life Vehicles (ELV) Directive aims to ban the use of
Cr(VI) in the manufacture of vehicles, although limits of 2 g per car are to be permitted for the foreseeable future. In addition, the Directive on Waste, Electrical and
Electronic Equipment (WEEE) aims to ban the use of Cr(VI). In the US, compelling
health data and legal suits are forcing OSHA regulators to lower the exposure limit
to chromic acid and it is anticipated that future exposure limits could be established
at levels between 20- and 200-fold below the current level. Past work carried out
in the US and UK has generally examined the viability of reducing emissions of
chromic acid (air pollution control techniques and chemical fume suppressants)
rather than applying fundamentally novel chemistries for chrome plating [60].
However, environmental and social pressures of operating chromic acid-based processes are imposing demands upon the industry, which cannot be met through
effluent reductions alone. In answer to this, at least three types of aqueous trivalent
chromium baths have been developed industrially [61–63]. However, finish quality,
11.3 Process Scale Up 311
cost and a perceived difficulty of operation has hindered the general acceptance of
these commercially available baths. Ionic liquid-based processes could provide a
route to overcome these problems.
Other disadvantages of the existing aqueous technology are economic in nature,
such as the low current efficiency of the reduction of Cr(VI) in acid media. In addition, the difference in over-potential between chromium and hydrogen reduction
results in the evolution of hydrogen gas, which can lead to hydrogen embrittlement
in the substrate.
A more general issue associated with aqueous solutions is that at some point
all water must return to the watercourse and hence the contamination with toxic
metal salts e.g. Cd (II) or Ni (II), complexing agents e.g. CN− or brighteners must be
minimized. Hence, while the use of ionic liquids to replace aqueous technology may
not seem to have an urgent technological or economic driver there are numerous
circumstances where the use of ionic liquids has specific advantages, such as the
deposition on passivated substrates, e.g. Al, or the efficient deposition of specialist
alloys that could not be carried out in aqueous solutions because of the chemistry
of water.
It is most probable that practical plating liquids for Cr, Ni, Cu, Au, Ag and Zn
will use eutectic-based ionic liquids. Numerous Zn-based Type I eutectics have
been applied to small scale deposition studies but it is less likely that these will
be viable due to the comparatively low conductivities and high viscosities of these
liquids. Several companies are currently using Type III-based eutectic ionic liquids, primarily those with urea and ethylene glycol as the hydrogen bond donor,
to electrodeposit zinc and zinc-based alloys [64]. This is at the 10–25 l scale using
soluble zinc anodes. High current efficiencies can be obtained at low current densities but the morphology and current efficiency deteriorate as the current density
increases. The main technological difficulty associated with the further scale up
of these plating baths is the development of effective brighteners that function in
ionic liquids.
Outside of these seemingly niche markets the main driving force for using nonaqueous electrolytes has been the desire to deposit refractory metals such as Ti,
Al and W. These metals have numerous applications, especially in the aerospace
industry, and at present they are deposited primarily by PVD and CVD techniques.
The difficulty with using these metals is the affinity of the metals to form oxides.
All of the metal chlorides hydrolyze rapidly with traces of moisture to yield HCl
gas and hence any potential process will have to be carried out in strict anhydrous
conditions. Therefore the factor most seriously limiting the commercialization of
aluminum deposition is the engineering of a practical plating cell.
Notwithstanding the perceived difficulties with commercializing such technology, a commercial aluminum electroplating process is already in existence and
has operated for over 10 years [65, 66]. It is based on triethylaluminum in organic
solvents such as toluene and although precise technical details are not given in the
open literature it is apparent that the process is successful. It is also highly probably
that a plating bath based upon a chloroaluminate ionic liquid is less water sensitive
than the organics solution.
312 11 Technical Aspects
11.3.2
General Issues
There are several general issues where ionic liquids differ from aqueous solutions.
Some of these are discussed in greater detail in the preceding chapters and all
are discussed in more detail in a recent review [67]. The key issues are clearly
associated with developing non-aqueous processing protocols and accounting for
the differences between the physical properties of a non-viscous polar fluid and a
viscous ionic liquid.
11.3.2.1 Material Compatibility
In general, ionic liquids tend to be non-corrosive towards most metallic and polymeric materials that would normally be encountered in electroplating or electropolishing situations so there is no reason why they could not be simple “drop-in”
replacements for aqueous systems. The majority of plating plants are constructed
from polymers such as polyethylene, polypropylene, nylon and PVC, all of which
are stable in the majority of ionic liquids.
The only large scale tests that have been carried out using ionic liquids were in
collaboration between Anopol Ltd and the University of Leicester for the electropolishing of stainless steel. Figure 11.14 shows a 1.3 m3 tank that was constructed from
polypropylene with polypropylene, nylon and polyethylene fittings and run as a pilot plant. It has a standard 3 kW heater to maintain the liquid at 50 to 60 ◦ C. Tank
agitation was achieved by recirculation of electrolyte via eight banks of inductor
nozzles [68].
As with aqueous solutions, “dip coatings” can be obtained when more electronegative metals are placed in ionic liquids containing more electropositive metal ions
e.g. silver ions will be deposited onto copper metal. Unlike aqueous solutions,
Fig. 11.14 Electropolishing bath (1300 L) operating at Anopol Ltd.
(Birmingham, UK) based on an ethylene glycol:choline chloride
eutectic.
11.3 Process Scale Up 313
however, these dip coatings tend to be more adherent. It should also be noted that
the redox potentials of some metals can be significantly shifted from the standard
aqueous redox potentials due to the differences in metal ion speciation.
The main differences will occur with the design of baths suitable for aluminum
and other water-sensitive metal salts. The evolution of HCl will require materials
which are more corrosion resistant, and the main difficulty will be in the development of plants which will allow the transfer of pieces in and out of the liquid under
strictly anhydrous conditions.
11.3.2.2 Pre-treatment Protocols
The aim of any pre-treatment protocol is clearly to remove non-metallic detritus
from the surface and will naturally involve a wash with a solvent to remove organic
residues and an acidic or alkaline clean to dissolve inorganic residues. Chapter ?
discusses the different approaches that can be used but, in principle, these are the
same that are currently employed with standard aqueous electroplating baths. The
key issue is to introduce a dry substrate into the ionic liquid and this will involve
either a drying stage or a rinse in an ionic liquid prior to immersion in the plating
liquid. Several methods have been studied but by far the best adhesion is obtained
by degreasing in a chlorinated solvent, followed by an aqueous pickle, rinse, dry and
then anodic etch in the ionic liquid prior to deposition. Anodic etch potentials and
times are dependent on the substrate and the ionic liquid used. Metals such as Al
and Mg will require a larger anodic pulse for a longer period than other metals such
as Cu or Ni. Metal oxide dissolution is easier in ionic liquids containing a metal that
is a good oxygen scavenger. Endres has shown that the adhesion of aluminum to
mild steel is greatly enhanced by an anodic pulse prior to deposition. It was shown
that an alloy was formed between the substrate and the coating metal, improving
adhesion [69]. In situ anodic etching may not always be feasible if the substrate
is difficult to re-deposit e.g. steel. In this situation a build-up of contaminating
metal ions in the ionic liquid could change the physical properties of the liquid and
damage the quality of the coating. Anodic etching should take place in a compatible
ionic liquid and the etched substrate is then transferred to the plating tank.
11.3.2.3 Conductivity
The conductivity, κ, of an ionic liquid is also strongly dependent upon temperature
and in an analogous manner to viscosity it is found to change in an Arrhenius
manner
ln κ = ln κ0 −
E
RT
(11.10)
where E is the activation energy for conduction and κ 0 is a constant. It has been
noted that the empirical Walden rule (η = constant) is applicable to ionic liquids,
where is the molar conductivity and η is the viscosity [70]. Deviations from
the Walden rule have previously been used to explain ionic association in proton
transfer ionic liquids [71, 72]. The Walden rule is normally only valid for ions at
314 11 Technical Aspects
infinite dilution where ion–ion interactions can be ignored but with ionic liquids
it is the availability of holes that allow ion migration that limits charge flow. Since
the fraction of suitably sized holes in ambient temperature ionic liquids is very low
the movement of holes can be defined by a combination of the Stokes–Einstein and
Nernst–Einstein equations [73, 74]:
λ+ = z2 F e/6πη R+
(11.11)
where z is the charge on the ion, F is the Faraday constant and e is the electronic
charge. This explains why so many studies of conductivity in ionic fluids have noted
that the empirical Walden rule is valid. Since the Stokes–Einstein equation is valid
for both ions then the conductivity of the salt can be determined since
(11.12)
= λ+ + λ−
an expression can be written for the conductivity, κ
κ=
z2 F e
6πη
1
1
+
R+
R−
ρ
Mw
(11.13)
where ρ is the density and Mw is the molar mass of the ionic fluid. Hence all
of the theories developed for limiting molar conductivities in molecular solvents
are also applicable to ionic liquids where there is an infinite dilution of suitably
sized holes [74]. Using this theory it is possible to estimate the limits of viscosity and conductivity that an ionic liquid can achieve. It is difficult to foresee an
ionic liquid that has a conductivity significantly in excess of EtNH3 + NO3 − (ca.
150 mS cm−1 at 298 K) and this must be viewed as a probable upper ceiling without
modification [75].
11.3.2.4 Added electrolytes
The conductivities of most aqueous electroplating solutions are in the region of
100 500 mS cm−1 because they are mostly high strength aqueous acids [76] and
this allows high current densities to be applied with only limited ohmic loss.
Significantly lower conductivities are obtained with ionic liquids and one way to
increase the conductivity could be to add a small cation such as Li+ that could
have better mobility compared to the large organic cation. This has been attempted
by a number of groups, particularly those developing lithium ion batteries, but
the effect on the conductivity has not been as significant as expected [77, 78]. The
viscosity and freezing point of the liquid are, however, affected as the small cation
will be strongly associated with the anions and little increase in the conductivity
is generally achieved. Other salts such as Na+ and K+ have negligible solubility
in most ionic liquids. The addition of electrolytes is clearly an area that requires
considerable investigation in the future.
The structure of the double layer is also affected by the addition of lithium ions.
Few studies have been carried out on the structure of the double layer in an ionic
11.3 Process Scale Up 315
liquid but those that have tend to suggest that the models used for aqueous solutions
are inappropriate in ionic liquids [79, 80]. If the metals are reduced at potentials
below the potential of zero charge then the electrode must be coated with a 6–7 Å
thick layer of cations. Adding small ions such as Li+ to an ionic liquid will decrease
the Helmholtz layer thickness considerably and should make metal ion reduction
easier. This should simplify nucleation and it has been shown qualitatively to be the
case for the deposition of chromium from a eutectic mixture of chromium chloride
and choline chloride. The incorporation of up to 10 mol% LiCl led to a change in
deposit morphology from microcrystalline to nanocrystalline and a change in visual
appearance from metallic to black [81]. It has also been shown that the addition of
LiF to 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)imide allows the
deposition of dense, thick, corrosion resistant coatings of tantalum [82].
11.3.2.5 Brighteners
Brighteners are essential to most electroplating systems and act to decrease the
surface roughness and improve reflectivity. Brighteners are thought to function
by either forming metal complexes, which shift the reduction potential and hinder
metal nucleation, or by adsorption on the electrode surface blocking nucleation and
hindering growth. In aqueous solution most brighteners are complex mixtures of
components, many of which are derived by serendipity, but most have the function
of viscosity modifiers or amphiphillic molecules that can specifically interact with
the metal surface. No systematic studies have been carried out in ionic liquid using
the types of brighteners used in aqueous solutions and this is clearly an area that
needs to be addressed to see if the brighteners function in the same way as they do
in water. The Abbott group has carried out studies using commercial brighteners
for zinc plating from Type III eutectics but to date none of these have shown any
improved surface finishes. To some extent this is not surprising given:
Ĺ The viscosity of the ionic liquids is much higher than aqueous solutions affecting
mass transport,
Ĺ The double layer structure is totally different in the two liquids and hence the
surface potential will differ, meaning that specific adsorption of organics will
differ,
Ĺ The metal speciation is different and hence the reduction potential will be shifted,
Ĺ Electrode processes will be different due to the lack of proton or hydroxide ions
in ionic liquids.
It may seem to be an impossible task to find a brightener compatible with an
ionic liquid but comparison of practical aqueous plating solutions with current
ionic liquids shows a fundamental difference in the metal speciation. In aqueous
solutions most plating is carried out with either strong bases e.g. KOH for zinc
plating, strong complexing agents e.g. CN− for silver plating or metals in the oxide
form e.g. CrO3 for chromium plating. These will tend to shift the reduction potential
to more negative values, decreasing the rates of nucleation and growth.
316 11 Technical Aspects
Applying the same principle to ionic liquids a number of compounds containing
nitrile, carboxylate and amine functionalities have been tested. Limited success has
been achieved with Ni, Ag, Cu and Zn baths. Brighteners that involve a complexation with a solution-based species will depend upon the comparative strength of
the ionic liquid–metal interactions. It would therefore be logical to suppose that
ionic liquids with discrete anions would be likely to work directly with brighteners used in aqueous solutions, as the interaction between the metal salt and
the anion will be considerably weaker than those between the metal salt and the
brightener. In eutectic-based ionic liquids the chloride anions act as strong Lewis
bases and could decrease the relative interaction between the metal salt and the
brightener.
Brighteners which rely on electrostatic or hydrophobic interactions may function
in ionic liquids but their efficacy is likely to be surface and cation/anion specific.
As with other solutes in ionic liquids, the general rule of like dissolving like is
applicable i.e. ionic species will generally be soluble as will species capable of
interacting with the anion. Aromatic species tend to exhibit poor solubility in ionic
liquids consisting of aliphatic cations and vice versa.
We have also studied the use of brighteners in Type III-based ionic liquids, as
well as the majority of brighteners that are used in aqueous zinc plating solutions
and none of them are active in ionic liquids. Some success has been achieved using
complexing agents such as ethylenediamine and acetonitrile but this has not been
a significant improvement. Figure 11.15 shows an AFM image of silver deposited
from a ChCl:2urea eutectic. It can be seen that in the absence of any brighteners
a relatively rough surface is obtained whereas the addition of ethylenediamine acts
as a brightener producing a much smoother surface finish. Endres studied the use
of nicotinic acid for the deposition of Pd and Al/Mn alloys from an AlCl3 -1-butyl-3methylimidazolium chloride ionic liquid and showed that, in contrast to producing
a brighter surface finish, it aided the formation of nanocrystalline deposits.
Fig. 11.15 Silver coating deposited from a urea:ChCl eutectic in the
absence (a) and presence (b) of ethylenediamine as a brightener.
11.3 Process Scale Up 317
11.3.2.6 Counter Electrode Reactions
As outlined in previous chapters the counter electrode reactions occurring in ionic
liquids will be significantly different from those in aqueous solutions. Given the
increased ohmic resistance that will be encountered compared to aqueous solutions
it will be preferable to use soluble anodes which will decrease the over-potential that
needs to be applied between the electrodes. Soluble anodes will also minimize the
breakdown of the ionic liquid itself and retain the bath composition in its original
state.
Anodic dissolution of most metals will occur in ionic liquids due to the absence
of passivating films on the electrode surface. Hence metals such as Al and Cr could
potentially be used as anodic materials. While this is potentially useful it should
also be noted that caution should be exercised when choosing a suitable material
for jigs or connectors that will be immersed in the ionic liquid.
No systematic study of inert electrode materials has taken place to date and
nothing is known about the anodic processes taking place in ionic liquids. It is
probable that noble metal oxide coatings should be suitable but processes such as
chlorine evolution will clearly have to be avoided for eutectic-based ionic liquids.
The breakdown products of most cations are unknown but it is conceivable that
some of them could be potentially hazardous.
11.3.2.7 Post-treatment Protocols and Waste Treatment
Treatment of the sample following electrodeposition has primarily been carried out
using a simple aqueous washing procedure. While this is an extremely effective
method it may not ultimately be applicable to large scale production due to toxicological issues with some of the anions or cations. Some ionic liquids have been
developed with biodegradable cations and anions but the liquids will still contain
large metal ion concentrations and some complexing agents which would be better
to keep separate from aqueous systems. The amount of “drag-out” and the extent
of the issue will depend upon the viscosity of the liquid. To circumvent the need to
process large volumes of rinse water it may be more practicable to rinse the piece
with a liquid that is immiscible with the ionic liquid, which will allow the separation of the ionic liquid in a settling tank. The most appropriate washing liquid will
depend upon the nature of the ionic liquid and the phase behavior of most ionic
liquids is well documented.
11.3.2.8 Supply
The majority of aqueous plating solutions are supplied as finished products by
major distribution suppliers. No such analogue exists for ionic liquids as no plating
processes have been developed with a sufficiently good surface finish to replace the
aqueous competitor. A number of companies make or distribute ionic liquids on
the > 100 kg scale. These include BASF, Merck, Scionix and Solvent Innovation
although laboratory scale amounts can now be obtained from a wide range of
chemical supply houses.
318 11 Technical Aspects
Fig. 11.16 Recycling of ionic liquid (1 ChCl : 2 ethylene glycol) used
to electropolish stainless steel; (a) used liquid containing Fe Cr and
Ni salts, (b) as (a) with 1 equiv. v/v added water, (c) as (b), after gravity filtration and subsequent removal of residual water by
distillation.
11.3.2.9 Recycling
Given the cost and environmental compatibility of most ionic liquids, recycling
protocols will be essential. Many of the issues will be associated with separating the
metals from the ionic liquids. The same issues also exist in aqueous solutions and
they are usually addressed by the addition of concentrated base which precipitates
the metals as an oxide or hydroxide. The solutions then need to be filtered and neutralized before disposal. Similar ideas will need to be developed for ionic liquids
i.e. ligands that can be added to precipitate the metals. An alternative approach
is to add sufficient diluent to the ionic liquid, thus changing the solvent properties such that the specific metal becomes insoluble. This idea has been applied to
the recycling of the commercial electropolishing solution. The electrodissolution
of stainless steel produces an ionic liquid that contains high iron, chromium and
nickel concentrations. The metals are present as glycolate complexes and the addition of water renders the complexes insoluble, Figure 11.16. This has the advantage
that it decreases the viscosity of the mixture and permits easier filtration. The water
can be distilled from the mixture with minimal loss of the ionic component. While
this process will only be applicable to a limited number of ionic liquids analogous
processes should be possible using other solvents.
11.3.3
Conclusions
Although no plating processes have been developed to date using ionic liquids,
it is clear that the advantages afforded by this new technology will certainly have
commercial applications. There are some issues associated with process scale up
but these are only analogous to the aqueous solutions and are not insurmountable.
The potential high current efficiency and longevity of the ionic liquids should make
the economics of the processes beneficial and with the current groundswell of
interest in the area it is highly likely that the plating industry will see at least some
processes entering the market within the next ten years.
11.4 Towards Regeneration and Reuse of Ionic Liquids in Electroplating 319
11.4
Towards Regeneration and Reuse of Ionic Liquids in Electroplating
Daniel Watercamp, and Jorg Thöming
In electroplating, impurities can be assumed to interfere with the intended deposition, deteriorating surface qualities or narrowing drastically the potential window
available. Due to their relatively high price and the anticipated cost for discharge of
spent liquors a breakthrough of ionic liquids in electroplating applications can be
expected to be linked to successful regeneration options. Because ionic liquids are
non-volatile and typically one to three orders of magnitude more viscous than water,
their regeneration by separation from mixtures and purification is a challenging
task.
However, it can be assumed for most electrochemical applications of ionic liquids, especially for electroplating, that suitable regeneration procedures can be
found. This is first, because transfer of several regeneration options that have been
established for aqueous solutions should be possible, allowing regeneration and
reuse of ionic liquid based electrolytes. Secondly, for purification of fresh ionic
liquids on the laboratory scale a number of methods, such as distillation, recrystallization, extraction, membrane filtration, batch adsorption and semi-continuous
adsorption in a chromatography column, have already been tested. The recovery
of ionic liquids from rinse or washing water, e.g. by nanofiltration, can also be an
important issue.
11.4.1
Introduction
For electroplating purposes ionic liquids show several attractive properties, such as
large electrochemical windows, specific solvent characteristics and extremely low
vapor pressures compared to ordinary solvents. When used as base electrolytes
in electroplating, ionic liquids can allow new processes that are impossible in
conventional electroplating where the main solvent used is water.
Despite the great electrochemical potential offered, these new compounds have
to compete with water in terms of “greenness”, given that water is itself the most
environmentally benign solvent. Nonetheless, an advantage of ionic liquids is that
they do not evaporate, even at elevated bath temperatures, avoiding heat and mass
losses during processing. However, “greenness” should not be attributed to a compound due to a single characteristic, the whole system has to be considered and
every single chemical entity has to be assessed with respect to its entire life cycle.
Based on the findings of a case specific analysis, a relative degree of greenness can
be attributed to comparative process and compound alternatives. In principle, the
degree of greenness can be determined through the following four main aspects:
Ĺ Greenness of the manufacturing process of the ionic liquid
Ĺ Risk potential of the technical application of the ionic liquid (leakages, toxic and
eco-toxic effects, fate of the compounds in the environment)
Ĺ Possibility of regeneration, recycling and reuse of the ionic liquid
Ĺ Waste treatment options.
320 11 Technical Aspects
In general, recyclability is crucial for the design of sustainable chemical processes
[83]. The aspect that should be elaborated here is the possibility of regeneration and
reuse of the ionic liquid, depending on the type of impurity and the sensitivity of
the specific application towards contamination.
Despite the huge number of publications dealing with the application of ionic
liquids, there are only a couple that include reuse aspects. To the best of our
knowledge, there is none that deals with regeneration of spent ionic liquid based
electrolytes. The intention of this contribution is to bridge this gap and suggest
potential concepts for ionic liquid regeneration.
In this chapter, an introduction to the principles of regeneration as they have been
developed in the field of water-based electroplating is given. With this background,
a discussion of the purification options for ionic liquids is presented, followed by a
first case study.
11.4.2
Recovery, Regeneration and Reuse of Electrolytes in Electroplating
11.4.2.1 The Concept
A general approach towards both more economical and more environmentally
benign applications of electrolytes in electroplating is the minimization of losses
and purge stream optimization. Losses are caused by drag-out, i.e. electrolyte that
clings to workpieces when they are removed from the plating bath. This makes
subsequent rinsing of the workpieces necessary, through which the losses are
diluted and discharged into the wastewater. Purge streams could be necessary as
a measure for product quality assurance. This implies that, by replacing all these
losses with fresh electrolyte, the so called make-up, relevant contamination can be
kept below critical levels.
To reduce the consumption of fresh electrolyte, diverse general approaches are
possible, such as the recovery and reuse of the losses from product and wastewater
streams, the recycling of spent liquid back to the manufacturing of the electrolyte
and the reuse of purge streams within the plating process. Fundamentally, all these
approaches require regeneration of the electrolyte prior to reuse or recycling. Ideally,
the regeneration makes use of the selective separation of the minor compound, the
impurity, from the electrolyte. Without such a regeneration step, impurities would
accumulate and eventually interfere with the intended functions of the electrolyte,
thereby reducing product quality.
There are several possible sources of impurities in the electrolytes and reasons
for their potential accumulation during use. Key amongst the sources, are the
unavoidable side-reactions. Others include the widespread practice in electroplating
processes of using the more convenient open systems that allow easier handling of
workpieces. Consequently the absorption of atmospheric gases and particles might
introduce impurities.
The overall concept for recovery, regeneration and reuse in electroplating is
shown in Figure 11.17. It includes the recovery stage, in which the workpieces are
rinsed for further cleaning and the diluted electrolyte received. The diluted solution
11.4 Towards Regeneration and Reuse of Ionic Liquids in Electroplating 321
Fig. 11.17 Concept of sustainable use of process bath liquors in electroplating: recovery (rinsing and concentration), regeneration (concentration and purification) and reuse.
is then concentrated using membrane systems [84] producing wastewater. However, there exists another strategy for avoidance of the production of wastewater and
the reuse of the diluted stream in rinsing, thereby achieving zero-water discharge
systems [85]. In this case, both, the concentration and the purification units can be
part of the regeneration, as shown in Figure 11.17.
More significantly, the design of the whole system should be subject to an
optimization process with respect to sustainability aspects such as cost and wastes
[86]. As shown in industrial applications, there are surface finishing systems for
which the recovery of electrolytes is feasible and economically attractive [87].
11.4.2.2 Regeneration Options for Water-based Process Liquors
Chemical and electrochemical surface treatment processes such as electroplating,
pickling, and etching often have a high consumption of chemicals and produce a lot
of wastewater and heavy metal wastes. Consequently, cost saving and environmental compatibility lead to the necessity of applying purification and concentration
units. Purification units can be divided into two groups. The first group treat spent
plating solutions, while the latter treat rinsing discharges.
Regenerators for Spent Process Liquors. Most effort in developing regeneration
methods for water-based process liquors in metal finishing has been spent
on chromium plating baths. These solutions contain a significant amount of
chromium and a lesser amount of other heavy metals, which make them a significant environmental concern and obvious targets for regeneration and reuse.
Typically a two-chamber electrolytic cell is applied and different electrode materials
have been tested [88]. The cell allows oxidation of Cr(III) to regenerate Cr(VI) in
the anode compartment. The removal of dissolved metal impurities such as Fe(II),
Fe(III), Cu(II), and Ni(II) from contaminated chromic acid solutions can be performed through electrodialysis in the same two-chamber cell as the chromic acid
recovery, where the impurities that electromigrated into the cathode compartment
are deposited or precipitated.
To achieve chemically robust low-cost separators, ceramic membranes have been
suggested by Sanchez et al. [89]. A Nafion 117 membrane and a ceramic diaphragm
322 11 Technical Aspects
separator were compared by Huang et al. [90]. Their results indicated that a system
using the Nafion separator and a small catholyte/anolyte volume ratio was best
suited for removing impurities from concentrated plating solutions. Similarly,
using a ceramic membrane for chamber separation, Jegadeesan et al. found up to
69% impurity removal [91].
To reduce the energy demand in such a system Huang et al. [92] modified the
set-up successfully and found that the average removal rate of each impurity was
approximately proportional to the product of its initial concentration and the separator area/anolyte volume ratio. More detailed investigations have been reported
in Huang et al. [93].
In the case of dissolved metal as major additive compounds, a combination of
precipitation and redissolution can be applied for recovery from spent solutions.
Gyliene et al. [94] found, for recovery of the main additive in nickel electroless
plating, that the Ni(II)-citrate complex could be precipitated with alkali followed by
redissolution in citric acid for reuse in electroless nickel plating after separation of
the precipitate. Additionally, for decontamination of spent electroless nickel plating
solutions Fe(III) can be used to precipitate the pollutant.
To simultaneously recover the metal and sulfuric acid from spent process liquors
of nickel electrolysis, Xu and Yang [95] tested diffusion dialysis successfully. The
membrane used was surface-cross-linked with aqueous ammonium to decrease
waste volume expansion caused by the water osmosis. They could control nickel
leakage within 4% and recover about 70% of the acid.
Alternatively to diffusion dialysis, Pierard et al. [96] suggested electrodialysis as a
regeneration process. In the case study involving acid pickling before electroplating,
they demonstrated the selection of ion-exchange membrane couples as well as the
development of tools to promote the use of electrodialysis in industrial applications.
For removing organic compounds, adsorption could be a good choice. This
applies to many decomposition products that might occur during electroplating
as well as for additives such as polyethylene glycol, a major organic additive in
copper electroplating solution, used as a brightening and stabilization agent in the
low ppm concentration range. Chang et al. [97] reported a successful application of
activated carbon, Calgon Filtrasorb 400, to remove polyethylene glycol from used
electroplating solution in order to reuse it.
Other unit operations that can be used in this field of process liquor treatment
are evaporation and crystallization. Both were tested by Ozdemir et al. [98] for
regenerating waste pickling liquors from hydrochloric acid pickling baths and
are reported to be suitable for small to mid-scale plants, currently neutralizing
and discarding waste pickling liquors. Even the relatively expensive crystallization
process, which can be used for removal of ferrous chloride to enable the recycling
of unused acid, was found to bring some improvement.
Purification Units for Rinsing Solutions. The second group of purification units
comprises those for treating rinsing discharges. For example, dissolved metals can
be separated by applying a combination of electrodeposition and electrodialysis,
as reported by Bolger and Szlag [99]. They recovered nickel from the rinse water
11.4 Towards Regeneration and Reuse of Ionic Liquids in Electroplating 323
cathodically in an electrolytic cell separated by an anion exchange membrane.
Depending on the anions used in the electrolyte such a process generates anodically
a sulfuric/hydrochloric acid mixture. However, additives like boric acid that are
characterized by high acid dissociation constants cannot be recovered by anion
exchange.
A widespread technology for purifying diluted aqueous solutions and even electroplating waste solutions is ion exchange [100]. This is also true for rinsing solutions [101, 102]. In technical systems a set of ion exchange columns is applied [85].
Usually liquid for regenerating the columns is discharged afterwards, but in some
cases recovery of valuables is also possible [103].
For purification of aqueous solutions the use of adsorption processes for cationic
impurities is also common. As economical adsorbents, montmorillonite, tobermorite, magnetite and silica gel were found sufficient for the removal of Cd(II),
Cr(VI) and Cu(II) in rinsing wastewater from a plating factory [104]. From this
investigation, it was found that the removal efficiency tended to increase with increasing pH and decrease with increasing metal concentration. This method allows
the realization of a rapid, simple and cheap rinse water treatment system for the
removal of heavy metals.
A complete process scheme for regeneration and reuse of spent final rinse water
from an electroless plating operation has been developed by Wong et al. [105]. It
includes (i) pre-treatment by microfiltration, UV irradiation, carbon adsorption; (ii)
heavy metal removal by nanofiltration and (iii) polishing using an ion exchange
mixed bed. The results of a pilot study showed that high quality product water with
an overall water recovery of 90% could be produced with an estimated payback
period of less than 18 months.
Concentration Units. Typical concentrators for rinsing solutions are membrane
filtration units, which split the feed into diluate and concentrate streams, meaning
purification and recovery, respectively [106]. Both nanofiltration and reverse osmosis might be applied, depending on the physico-chemical properties of the solutes.
To produce highly concentrated solutions suitable for re-use in plating baths, high
pressure reverse osmosis might be necessary [84].
A combination of electrodialysis with a concentrator media, ion-exchange resins
or activated carbon in the catholyte chamber has been suggested by Chaudhary
et al. [107]. Besides anodic chromium regeneration about 90% of dissolved copper
could be recovered.
Another approach to achieve purification of rinses and recovery in one step,
electrodialysis has been suggested for chromic acid recovery and removal of metallic
impurities [108]. As the authors point out there are two main process limitations:
first, the poor stability of most anion-exchange membranes against the oxidative
chromic acid solution and secondly the increase in membrane resistance due to
the formation of polychromates in the membrane.
Recovery of Minor Compounds. Extraction and separation of nickel(II) and its
recovery from spent electroplating bath residue is reported by Singh et al. [109].
324 11 Technical Aspects
Along with Cr(III), Fe(III), Mn(II), Co(II), Cu(II) and Zn(II), Ni(II) was removed
from sulfuric acid media, employing a Cyanex 301–toluene system. The success
depended on various parameters such as the concentration of the acid, metal ion
and extractant and the nature of the diluent.
A more selective recovery of nickel from plating wastewater was described by Eom
et al. [103].They used a column packed with strongly acidic cation resin through
which over 99% nickel ion was removed. In this process, sulfuric acid was employed
with a reagent in order to regenerate nickel ions from the resin adsorbed. Moreover,
the nickel ions recovered by sulfuric acid were obtainable up to 120 g-Ni L−1 allowing
reuse in the plating bath.
Investigations by Malinowska et al. [110] have shown that absorption can be used
to recover 90% of ammonia that is vaporized during chemical bath deposition of
cadmium sulfide thin layers from which concentrated solutions with more than
10 mol L−1 of pure ammonia can be obtained. Additionally, a cake with mixed
cadmium sulfide–cadmium cyanamide is produced, from which cadmium can be
recovered hydrochemically as cadmium sulfate [111]. The global process recovers
up to 99.999% of cadmium and generates only solid sulfur and a liquid effluent
containing traces of cadmium.
Finally, impurities that accumulate during usage of electrolyte can also be recovered. For example, Ni–Cu–Zn ferrite powder can be prepared from steel pickled
liquor and electroplating waste solutions by a hydrothermal process [112].
Transfer from Water-based to Ionic Liquid Based Liquors. In the case of water-based
electrolytes, there are two economic incentives for the above mentioned approaches:
the recovery of valuables and the avoidance of wastes and wastewaters. Despite the
environmental attractiveness of such measures economic constraints may become
an obstacle in industrial application.
For ionic liquid based process liquors, the contrary can be assumed. Due to
their relatively high prices and anticipated costs for discharge of spent liquors a
breakthrough of ionic liquids in plating applications can be expected to be linked
to successful regeneration options.
Even though regeneration units have not yet been reported for ionic liquid based
electrolytes, it is most likely that some of those mentioned above could be transferred to this new field. For example, the application of electrodialysis could presumably allow removal of ionic impurities from ionic liquids. As in water-based
electrolytes, it should be possible to separate small and relatively highly charged
metal cations across cation exchange membranes and then to precipitate them out
in an alkaline catholyte. But for such a method the complementary anodic process
has to be designed carefully. For example, there should be another species to be
oxidised such as Cr(III) in spent chromic plating baths or the separated cations
could be replaced, for example, by anodic dissolution of the metal that is to be
plated. However, electroneutrality has to be guaranteed as a crucial constraint in
electric field driven separation processes.
Other unit operations that have been established for aqueous solutions could
be considered, to allow regeneration and reuse of ionic liquid based electrolytes.
11.4 Towards Regeneration and Reuse of Ionic Liquids in Electroplating 325
Actually, as can be seen in the following section, several of the separation methods
mentioned above have already been tested in the purification of at least fresh
ionic liquids. However, there is still some development necessary to come up with
sustainable regeneration units.
11.4.2.3 Regeneration Options for Ionic Liquids in Electroplating
Despite the huge number of publications dealing with the application of ionic liquids, to the best of our knowledge, there is only one paper [113] that mentions
general problems related to purification of ionic liquids for electrochemical applications and it appears that there is none so far that deal with regeneration of spent
ionic liquid based electrolytes. This is amazing, considering that the influence of
impurities often narrows drastically the potential window available, as illustrated
by Zhang and Bond [113]. However, a number of purification procedures have
already been tested on the laboratory scale for fresh ionic liquids with respect to
their downstream processing but little is known about efficiency on a technical
scale.
The reason for this lack of experience in large scale purification is quite simple:
downstream processing is avoided so as to minimize the production cost of ionic
liquids. On a commercial scale separation processes needed for purification can be
assumed to be more costly than improvements in the synthesis stage [114].
Regeneration Options for Ionic Liquids in Other Fields of Application. In fields of
application other than electroplating several examples of ionic liquid regeneration
and reuse are described in the literature. For example, in the field of new reaction
media [115, 116] or in the field of catalysis [117–121]. Even though they do not
deal with electrolytes, they are a useful guide to learning about possible concepts
and challenges. For example, Song et al. [122] described the reuse of an aminofunctionalized ionic liquid applied as a nucleophilic scavenger in solution phase
combinatorial synthesis. Here regeneration was necessary to remove extracted
electrophiles, such as benzoyl chloride and phenyl isocyanate, by a combination
of extraction and phase separation steps, such as decanting and filtration. Neither
FTIR nor 1 H NMR spectra showed any significant differences between the freshly
prepared and the regenerated ionic liquid. Here, the reusability of the regenerated ionic liquid was demonstrated by reusing it three times as scavenger with
comparable activity in terms of product yield and purity.
Thermal Unit Operations. The easiest case for regenerating ionic liquid electrolytes
is when the impurity is volatile. This is due to the negligible vapor pressure of the
ionic liquid and the resulting extreme vapor pressure difference. For such a task,
simple distillation in a single step is sufficient. If more than one volatile solute
is present in the solution from which one is to be removed selectively, the task
becomes more demanding. In this case, the other solutes would be lost through
the simple distillation process. Alternatively, the volatile components could be
separated from each other by repeated vaporization–condensation cycles within
a packed fractionating column. If the other solutes show lower boiling points
326 11 Technical Aspects
another method should be considered. Finally, the method is chosen on economic
grounds.
The same is true for the technical application of vacuum distillation that can be
performed by means of a rotary evaporator. For relatively high boiling temperature
compounds such as water, that is a common impurity in ionic liquids from many
applications, this technique is in general very useful, as it is for removing compounds with boiling points near or beyond the decomposition temperature of the
ionic liquid at atmospheric pressure. For high purity purposes, the target concentrations of contaminants are extremely low and vacuum distillation might also be
an option. For example, vacuum distillation at 120 ◦ C resulted in ionic liquids with
moisture content below 10 ppm, as reported by Appetecchi et al. [123]. Scott et al.
[115] found that successful reuse of an ionic liquid in a new synthetic route required
regeneration by removing the methanol, which was used as a precipitating agent,
under vacuum,.
As an alternative to simple distillation, pervaporation could be used [124]. This
technique makes use of non-porous membranes with a selective layer consisting
of hydrophilic or hydrophobic polymer. Those compounds, which are volatile and
soluble in the membrane, are evaporated into the vacuum on the permeate side.
By this means, selective separation, for example of volatile impurities from volatile
auxiliary agents in the ionic liquid, should be possible.
To the best of our knowledge this possibility has not yet been shown to work. The
two major challenges are the relatively high Reynolds numbers necessary inside
the membrane module and the need to find selective membranes suitable for ionic
liquids. Ceramic membranes show great potential for this application but so far
there are only a few choices available on the market.
Another thermal separation unit often used for the laboratory scale purification
of ionic liquids is recrystallization [125]. It is an attractive option for those ionic
liquids that can form solids with a high degree of crystallinity. Crystals of ionic
liquids are expected to be pure because each molecule or ion must fit perfectly into
the lattice as it leaves the solution. Impurities preferentially remain in solution as
they do not fit as well in the lattice. The level of purity of the crystal product finally
depends on the extent to which the impurities are incorporated into the lattice or
how much solvent is entrapped within the crystal formed.
In single-solvent recrystallization, the impure ionic liquid is dissolved in the
minimum amount of a single solvent necessary to give a saturated solution; the
solution is then allowed to cool. As cooling progresses, the solubility of the compounds in solution drops, resulting in the desired recrystallization. To enhance the
process, a seed crystal of the pure ionic liquid is preferably added to the saturated
solution resulting in these crystals forming first and thus leaving a greater ratio of
impurity in solution. In the case of unsatisfactory separation factors, multi-solvent
recrystallization can be tested. Here a second solvent, in which the impurities are
soluble and the ionic liquid is not, is added carefully to the solution.
As mentioned above, a possible drawback of recrystallization is the potential
presence of solvent traces in the ionic liquid. This might result in the formation of
yellowish compounds, as was reported by Appetecchi et al. [123].
11.4 Towards Regeneration and Reuse of Ionic Liquids in Electroplating 327
Extraction Processes. The extraction procedure usually applied for hydrophobic
ionic liquids containing hydrophilic impurities is “washing” with water. However,
this method is a problem for certain types of ionic liquids that undergo hydrolytic decomposition, such as those containing hexaflurophosphate. At first glance, washing
appears to be a simple and cheap method, but in large scale applications problems
related to the wastewater issues may arise. Even though hydrophobic ionic liquids have a low solubility in water, the concentrations are relatively high, typically
ranging from 0.1 to 10 g L−1 in the discharge. The fact that most cations and hydrophobic anions do not show significant biodegradability, coupled with the loss of
costly materials in the discharge, explains the problem in large scale applications.
A potential solution of this problem lies in the application of nanofiltration. The
success of solvent extraction to remove polar or non-polar compounds from ionic
liquids appears to depend strongly on the system for which it is used. While in
some cases only “mixed success” is reported [126], in other applications solvent
extraction has been shown to lead to excellent results, for example extraction with
hexane [127].
It has also been shown that in some cases consecutive removal of the extractant
is necessary if it partly dissolves in the ionic liquid, as Zulfiqar and Kitazume [116]
reported for the application of diethyl ether. They purified the ionic liquids after
extraction by distillation at 80 ◦ C. Therefore, before planning for a process scale up,
there are some questions that need to be answered such as: (i) How often could the
solvent be reused directly? (ii) By what means could the impurity be removed from
the solvent? (iii) To what extent does the ionic liquid accumulate in the solvent?
and (iv) How does this accumulation influence the performance of the intended
separation?
The separation of the auxiliary agent can be easily handled on a technical scale
if it forms a pure phase. Otherwise more sophisticated separation methods are
needed. In the case of ionic liquids a process termed organic solvent nanofiltration
has been tested successfully [120, 128].
Adsorption Processes. Since adsorption processes often show high distribution
coefficients, several adsorbents are favorite candidates for removing low concentrations of impurities. An important group are chromophores. In the synthesis of
ionic liquids the formation of color, generally ranging from yellowish to orange,
has been attributed to side reactions, e.g. from excessive heating during synthesis
[129]. It can be assumed that color occurs at elevated temperatures for instance due
to formation of a dimer of the amine and the ionic liquid or ionic liquid precursor
in which the amine is dissolved. As an alternative to the avoidance of such side reactions during synthesis, several approaches for decolorization have been developed,
ranging from recrystallization or adsorption to extraction. The prevalent method in
the literature is definitely adsorption. A number of attractive adsorbents have been
tested already, among which are activated carbons and aluminas, synthetic zeolites,
and silica gels.
Both batch contactors as well as chromatographic columns have been suggested
for decolorization [130]. Nevertheless, to remove color with a single adsorbent was
328 11 Technical Aspects
not always sufficient, neither with powdered carbon in a batch contactor [131] nor
with alumina or silica in a semi-continuous column [125]. However, with subsequently applied powdered carbon and alumina [123] or in a combined chromatographic column with granular carbon and silica gel, as described by Earle et al.
[132], the adsorption process provided even better results. However, losses of ionic
liquid were significant in such adsorption procedures [123].
Redox Processes. Among the most serious impurity problems for electrochemical
applications is the contamination of electrolytes with halides. Since they easily
react anodically they can be expected to reduce the size of the electrochemical
window drastically but the readiness of their anodic decomposition can be used
for a decontamination procedure. This was recently described by Li et al. [133] for
chloride impurities. They found that, in combination with a subsequent removal
of the gaseous product Cl2 by absorption, electrochemically pure ionic liquids can
be obtained. Ethylene was bubbled through the solution to absorb the chlorine gas.
Without such an absorption step, the soluble complex Cl3 − was formed which could
not be removed by vacuum distillation. Both formation and subsequent removal of
the complex Cl3 − can be easily followed spectrometrically due to a strong band of
this species at 302 nm.
The crucial parameter for the anodic decomposition of halides is the anodic
potential. This is simply due to the dilemma that a minimum potential for decomposition is needed but degradation of the ionic liquid cation is enhanced with
increasing potential. It was found for the [BMIM] cation that at a voltage about 20%
above the decomposition value the appearance turned gradually from colorless to
light yellow.
Another option could be photochemical decomposition of impurities. Yang and
Dionysiou [134] described a combined approach for treating solids or liquids which
contain environmentally important organic contaminants. They suggest using
room-temperature ionic liquids as solvent media and a subsequent photolytic degradation of the contaminants. The second step, the photolytic degradation, could, in
principle, also be used for regeneration. It can be assumed that photolytic degradation is capable of degrading components in ionic electrolytes to below the required
limit concentrations. The constraint here is that metabolites may be produced,
which accumulate instead of the primary compound, exceeding their own required
limits.
Mechanical Processes. For removing particulate matter from low viscous liquids
filtration generally is the technique of choice. Gan et al. [135] studied microfiltration
characteristics of room temperature ionic liquids. They found that due to the relatively high viscosity it was impossible to get the tested liquids permeated through
the microfiltration membranes with ease. They suggested mixing the ionic liquid
with 20 % volumetric proportion of diluting polar agents, preferably methanol or
ethanol, to drastically reduce viscosity. Alternatively, it can be assumed that at elevated temperatures it should be possible to receive comparable results at high
temperatures without addition of another solvent.
11.4 Towards Regeneration and Reuse of Ionic Liquids in Electroplating 329
The separation of non-volatile products from ionic liquid solutions using nanofiltration was suggested by Kröckel and Kragl [136]. It was shown for both bromophenol blue and lactose, each in ionic liquid, that the product was rejected while the
ionic liquid permeated. It should be noted that in such cases the products are
not isolated. Instead, concentrated ionic liquid solutions are produced. However,
depending on the solubility, phase separation might occur.
Another already mentioned application of membrane filtration is for the recovery of ionic liquids from wastewaters. Here the challenge is to find appropriate
membranes, since rejection values that have been reported to date [136] are too low
for industrial application. However, for similar ionic liquids we found a membrane
that shows rejection rates above 99% throughout at considerably high permeate
flow rates above 50 L m−2 h−1 in cross-flow filtration. Such numbers make washing
in combination with nanofiltration an interesting option.
11.4.3
Case Study
Every single regeneration problem has to be analysed individually, however, the
following case study demonstrates how a selection of separation techniques, extraction and phase separation, can successfully be applied to regenerate a spent
ionic liquid based electrolyte satisfactorily. As a case study the electrolyte 1-butyl1-methylpyrrolidinium bis(trifluoromethanesulfonyl)amide ([BMP]Tf2 N) was chosen, which is used for electrodeposition of aluminum as described in the literature
[137, 138].
The spent electrolyte was prepared by IoLiTec GmbH as follows: Dry AlCl3
(26.2 wt%) was dissolved in [BMP]Tf2 N, resulting in a solid at room temperature.
At the process temperature (100 ◦ C) the mixture formed two liquid phases. The
deposition took place in the upper phase at a voltage of –2 V (anode: Al plate,
cathode: gold plate). After 90 min of deposition (charge: 145 A s) a mixed black and
silver colored coating was received cathodically, and the electrolyte was collected
for regeneration (Figure 11.18(a)).
Starting the regeneration procedure a 10 mL sample of the spent ionic liquid/
AlCl3 mixture was heated to 75 ◦ C and stirred under nitrogen flow (Figure 11.18(b)).
Deionized water was added stepwise (in amounts of 1 mL) with a syringe to the
stirred two-phase liquid.
During addition of water gas evolution could be observed in the vials. This could
either be due to the strongly exothermic hydration (Hsolvation = –330 kJ mol−1 ) of the
AlCl3 (Eq. (11.14)) leading to generation of water vapor or to the thermal decomposition of the hexaaquaaluminum trichloride resulting in the liberation of HCl gas
(Eq. (11.15)).
AlCl3 + 6 H2 O → AlCl3 ·6 H2 O
(11.14)
AlCl3 ·6 H2 O → Al(OH)3 + 3 HCl + 3 H2 O
(11.15)
330 11 Technical Aspects
Fig. 11.18 Samples of spent [BMP]Tf2 N electrolyte (a) directly after
electrodeposition of Al and (b) stirred at 75 ◦ C under nitrogen atmosphere.
After total addition of 9 mL water, resulting in the weight fractions wIL = 0.45,
wAlCl3 = 0.16 and wH2 O = 0.39, the mixture was stirred for 15 min. Subsequently the
sample was shaken for an additional 10 min while cooling to ambient conditions.
At room temperature the mixture divided into two liquid phases. These phases
were separated (Figure 11.19) by centrifugation (20 min at 2460 g). The lower clear
and more viscous phase was presumed to be the IL phase and the upper liquid the
water phase. At the interface fine particles were collected.
Using a syringe the phases were carefully separated and transferred to different
vials. The ionic liquid phase was then submitted to an evaporation procedure
(rotary evaporator, 50 ◦ C, 10 mbar, 6 h). Shortly after connecting to vacuum bubble
generation could be observed.
Fig. 11.19 Samples after mixing with water, showing phase separation.
11.4 Towards Regeneration and Reuse of Ionic Liquids in Electroplating 331
Fig. 11.20 Recovered ionic liquid phase ([BMP]Tf2 N) after regeneration.
After evaporation the ionic liquid phase contained a dispersed precipitate. These
solid particles were concentrated at the bottom of the flask by centrifugation (20 min
at 2460 g) resulting in a very clear, slightly yellowish ionic liquid phase (Figure
11.20).
It should be stated here that on a technical scale washing requires a concept for
water reuse and recovery of ionic liquid from the wastewater. As already discussed,
nanofiltration is likely be a successful approach for the recovery task.
The regenerated ionic liquid phase was investigated electrochemically to determine its quality. Cyclic voltammetry was performed using a rotating platinum disk
electrode (500 rpm), a platinum counter electrode and a platinum wire as (quasi-)
reference electrode placed closed to the rotating disk.
In Figure 11.21 two ionic liquids are compared, a freshly synthesized [BMP]Tf2 N
received from Iolitec GmbH and the regenerated ionic liquid. It can be clearly seen
that current densities after regeneration are lower than for the fresh electrolyte
throughout the entire potential range. No additional signal can be recognized for
the regenerated ionic liquid. This indicates that none of the electrochemically
active additional ingredients, water and Al(III), remain. The regenerated ionic
liquid appears to be at least as pure as the originally synthesized ionic liquid. The
regeneration was successful.
In the fresh electrolyte a first anodic step starts at 1500 mV. This could be a hint
for chloride impurity. Since this signal almost vanished for the regenerated ionic
liquid, it can be assumed that the procedure presented is suitable also for purifying
fresh ionic liquids.
332 11 Technical Aspects
Fig. 11.21 Cyclic voltammograms of original and regenerated ionic
liquid [BMP]Tf2 N)). The potential was determined vs. Pt as quasireference electrode. Scan rate 5 mV s−1 .
The pros and cons of this approach are summarized as follows:
Ĺ Ease of process.
Ĺ Small amount of losses of ionic liquid, but aluminum salts are completely lost
after hydrolysis for the plating process.
Ĺ Pure ionic liquid as product.
Ĺ Only water as solvent necessary, which can be re-used in the regeneration process
to a certain extent.
Ĺ Before discharge of the wastewater dissolved amounts of ionic liquid need to be
recovered for environmental and economical reasons, e.g. by nanofiltration.
Ĺ In contrast to dead-end microfiltration, which could also be used to remove
solids from spent electrolytes producing (after addition of a solvent and at elevated
temperatures) an ionic liquid as residue, the residue in the extractive regeneration
is wet sludge only.
Despite the conspicuous advantages of the presented water-based regeneration
approach, it is still to be shown whether it can be transferred to other tasks and
whether the reuse of the regenerates in plating processes leads to surface qualities
similar to those received from fresh electrolytes.
11.4.4
Conclusions
A general approach towards both more economical and more environmentally
benign applications of ionic liquids is maximization of their lifetime. The measures to be applied in electroplating are recovery and regeneration, both to allow
11.4 Towards Regeneration and Reuse of Ionic Liquids in Electroplating 333
reuse. This study focuses mainly on regeneration, but also recovery of drag-out
is considered, which should be possible in conventional counter-current rinsing systems, albeit in combination with regeneration units such as membrane
concentrators.
This study focuses firstly on the transfer of regeneration principles as they have
been developed in the field of water-based electroplating and of purification options
for ionic liquids as they are experienced in other fields of ionic liquid application. A
number of purification procedures for fresh ionic liquids have already been tested
on the laboratory scale with respect to their finishing in downstream processing.
These include distillation, recrystallization, extraction, membrane filtration, batch
adsorption and semi-continuous chromatography. But little is known yet about
efficiency on the technical scale. Another important aspect discussed is the recovery
of ionic liquids from rinse or washing water.
However, the financial and environmental cost might be too high for a certain
approach. Hence the optimality of a solution is always subject to technical constraints and the technical bottleneck of each option has to be identified. For any
optimization approach it has to be considered that the demand for regeneration
is finally related to large scale applications. The mass flow that has to be treated
during regeneration will range typically from grams to kilograms per minute.
Since little is yet known about efficiency on the technical scale, future investigation should focus on (i) efficiency with respect to separation yield, energy demand
and amount of mass separation agents required, (ii) long-term re-use options of
auxiliary agents such as extractants or adsorbents and (iii) ease of scaling up. Furthermore, a crucial point for further development of regeneration will be to identify
the pollutants that disturb the main process as well as their critical concentration
levels in the electrodeposition process.
In a case study, the extraction of a spent, turbid electrolyte with water at elevated
temperature and subsequent phase separation is shown as an example. It could
be demonstrated that purification of ionic liquids for re-use is not necessarily as
difficult as suspected in the literature [113]. The case study is going to be continued
to demonstrate whether the application of the regenerates will lead to comparable
surface qualities. Accordingly, it is of future interest to see whether the roughness
of cathodic deposits using regenerated electrolytes shows similar dependences
on current density and temperature as does the roughness of deposits from fresh
electrolytes. Additionally, future investigations should consider gaseous impurities,
which could be dragged in easily due to the high solubility and capacity of many
ionic liquids for trace gases, especially sulfur compounds.
Acknowledgments
The authors wish to thank IoLiTec GmbH for providing ionic liquid and spent
electrolyte, and several partners of the BMBF project NEMESIS for fruitful discussions. Financial support by VDI/VDE-IT (project No. 16SV1970) is gratefully
acknowledged.
334 11 Technical Aspects
11.5
Impurities
Impurities are a concern in ionic liquids electrochemistry. Whereas even considerable amounts of impurities, like different metal ions, water or organic impurities,
might not disturb a technical process (e.g. extractive distillation, organic synthesis)
the wide electrochemical windows of an ionic liquid (∼± 3 V vs. NHE) allow the
electrodeposition of even reactive metals like lithium and potassium, as well as the
oxidation of halides to the respective gases. In the best case this codeposition only
leads to a low level of impurities, in the worst case fundamental physicochemical
studies are made impossible as the impurities are adsorbed onto the electrode surface and subsequently reduced. Furthermore, passivation or activation effects at
the counter electrode have to be expected.
In the last few years the different suppliers of ionic liquids have developed several
purity grades. Merck has introduced “synthesis”, “high purity” and “ultrapurity”
(see Chapter 1.2) and other suppliers also follow this purity scheme. In the past
much attention was focussed on impurity effects. However, as the different suppliers are about to establish large scale production lines where the costs for the educts
have to be quite low one has to be prepared that the problem of impurities may
return. As many groups (in part without any experience at all) have entered the
field of ionic liquids in recent years we would like to draw attention to the subject of
impurities. Impurities can be a concern but do not necessarily have to be a concern.
We could also imagine that for some processes impurities are beneficial but, as a
minimum, one should know their role in the respective process.
11.5.1
Origin of Impurities
11.5.1.1 Synthetic Impurities
The synthesis process represents a very significant source of impurities in ionic
liquids. Because of their typically low volatility, which makes distillation impractical,
and the lack of any straightforward crystallization method of purification, ionic
liquids are often delivered in a semi-impure state. Significant impurities include
starting materials, such as halides, and metal cations, such as lithium, sodium or
silver, and any impurities carried through from the synthesis of the organic cation,
in particular amines. Where halides such as bromide and iodide are present, some
oxidised species such as I3 − are often also present, generating color in the otherwise
colorless ionic liquid. Most of these can be quite difficult to remove; however, at
the 1% or below level, in many cases they can be tolerated as long as the impurity
level is consistent from batch to batch. Seddon [139] has discussed the impact of
low levels of impurities on physical properties including viscosity. Chloride ions
are particularly notable for their effect in lowering viscosity.
If analytical information is not available from the supplier of the ionic liquid it
is advisable to carry out analysis, using traditional AAS or ICP-MS methods for
the metals and halide-selective electrode analysis for the halides. Residual amine
11.5 Impurities 335
can be easily detected using a Cu(II) complex formation and UV/vis absorbance
measurements.
In some ionic liquids, acid (proton) impurities are significant. This is common
in the phosphonium cation family of ionic liquids and can also be the case with
nitrogen-based cations if the synthetic method involves a neutralization reaction. It
is relatively easy to deal with this situation. Acidity should be determined by a standard titration method and then the acidity neutralized by addition of an appropriate
base. Carbonates are particularly useful in this regard since they produce CO2 as
the product.
11.5.1.2 Water
Water presents a rather different problem in that its presence can originate from
the synthesis, or from handling and storage prior to (or even during) the electrodeposition. Notably, even ionic liquids such as [EMIM]Ntf2 ] that we think of as being
hydrophobic are nonetheless reasonably hygroscopic up to their saturation point,
so that storage and handling needs to involve an inert atmosphere. The presence of
this water is particularly significant in the potential region below –0.5 V (vs Ag/Ag+ )
where it produces reduction products directly and also may cause degradation of
the ionic liquid and/or a surface film to form on the deposited metal. Howlett et al.
[140] have suggested that [Ntf2 ] ionic liquids produce breakdown products of the
anion on metals such as lithium and magnesium in a reaction that is catalyzed by
reduction products of water such as the hydroxyl radical. These reduction products
may produce useful protective films in some cases, such as lithium, such that further reduction of the metal ion can take place via transport through the film, but
this is unlikely to be the situation in the case of more highly charged metal ions
such as Ti(II).
Water analysis can be routinely carried out by a Karl–Fischer analysis in which
the ionic liquid is diluted in methanol before analysis. A spiking approach can be
used to produce a calibration curve that allows for background effects. At very low
levels of water (<10 ppm) quite substantial sample sizes can be needed for this
method to be meaningful.
Other significant impurities can arise from the breakdown of the ionic liquid
during storage. This is particularly important in the case of the PF6 − ionic liquids.
It has been shown by a number of groups (sometimes with catastrophic results) that
small amounts of water in the ionic liquid (in some cases introduced by repeated
opening of a container of the IL in the laboratory atmosphere) can hydrolyze the
anion during storage to produce a variety of oxyfluoride species:
+
−
PF−
6 + H2 O = H + F + phosphorus oxyfluorides
The exact nature of the hydrolysis products is not well established, however, the
generation of HF is clear. Such ionic liquids have been observed to etch their glass
containers and, in one case, a small explosion occurred when there was a build up of
gas pressure in the bottle (presumably from the etch products). Notably, a number
of PF6 − ionic liquids are water immiscible; nonetheless they are able to dissolve
336 11 Technical Aspects
sufficient amounts of water for this process to be significant. For this reason the use
of PF6 − ionic liquids in electrowinning processes is not recommended, especially
for a continuous use electroreduction process, or any situation where the IL is
being recycled. While the build up of HF can be controlled and dealt with, the
additional analysis and process steps required mean that it is usually more effective
to consider an alternate ionic liquid. Therefore, liquids with PF6 − and BF4 − cannot
be recommended to the beginner!
11.5.1.3 Gaseous Impurities
Dissolved gases, in particular oxygen and nitrogen, are no less problematic in ionic
liquids than in other solvents. At very negative potentials, the presence of nitrogen
can be problematic in an unfamiliar way in that metals such as Li easily form
nitrides. There are also suggestions that the reduction products of oxygen may play
a catalytic role in ionic liquid degradation in some cases [140]. The solutions in
both cases are the familiar degassing strategies adopted in electrochemical work
whenever dissolved gases are an issue. The only additional aspect that one needs
to be aware of is that the higher viscosity of ionic liquids as compared to aqueous
or aprotic solvents, means that degassing methods generally need more time and
stirring. Increased temperatures during degassing facilitate the process and also
lower the gas solubility in a useful way.
11.5.1.4 Particulate Impurities
Given the unusual solvency properties of ionic liquids, especially towards ionic
materials, it is not surprising that there have been recent reports [141] concerning
the dissolved and nano-particulate impurities arising from absorbents used in
the synthetic process. Nano-particulate silica and alumina from this source can be
identified in many ionic liquids. Parts per million levels of Al and Si can be detected
by ICP-MS and there is evidence from one of our groups (F.E.) of this Al and Si
being electrodeposited (see below). Such impurities may, or may not be an issue,
depending on the metal being electrodeposited. They certainly may appear in the
electrodeposit as low level impurities, but they may also appear as surface and grain
boundary layers on the deposit.
11.5.2
Impurities in Deep Eutectic Solvents
Ionic liquids can be compared to any other liquid in that the reactivity of a species
will depend upon its relative activity in solution. To this end it is important to
consider the relative Lewis and Brønsted acids that can interact with the solutes to
affect their activity. It is also important to remember that ionic liquids with discrete
anions have wider potential windows and what we therefore hope to achieve with
them is more susceptible to the presence of reactive species. The influence of
impurities on the electrochemical behavior of an ionic liquid will depend upon the
relative Lewis acidity/basicity of the liquid and of the redox process in question.
Eutectic-based ionic liquids behave very differently from ionic liquids with discrete
11.5 Impurities 337
anions. The presence of a high concentration of both Lewis acid and base in the
eutectic mixtures makes them act like a buffer solution would in an aqueous
solution. Hence, the addition of a Lewis acid or base has minimal effect on the
properties of the liquid. An example of this is the addition of a simple halide salt
e.g. to a chloroaluminate ionic liquid. The presence of a large amount of AlCl3 in
the liquid means that the equilibrium:
−
a AlCl3 + b Cl− ↔ c AlCl−
4 + d Al2 Cl7
will be perturbed but the concentration of free chloride will still be negligible. The
ability of the ionic liquid to act as a buffer will depend upon the relative Lewis or
Brønsted acidities of the components. This is true of ionic liquids with discrete
anions, e.g. Tf2 N− will be approximately Lewis neutral whereas Tf − will be more
Lewis basic. These systems will be considerably more susceptible to the presence
of a Lewis Base such as Cl− .
The main contaminants in an ionic liquid will be introduced from the synthesis, absorbed from the atmosphere or produced as breakdown products through
electrolysis (see above). The main contaminants for eutectic-based ionic liquids
will be from the components. These will be simple amines (often trimethylamine
is present which gives the liquid a fishy smell) or alkyl halides. These do not interfere significantly with the electrochemical response of the liquids due to the
buffer behavior of the liquids. The contaminants can be effectively removed by
recrystallization of the components used to make the ionic liquids. For ionic liquids with discrete anions the major contaminants tend to be simple anions, such
as Li+ , K+ and Cl− , present from the metathesis technique used. These can give
significant difficulties for the deposition of reactive metals such as Al, W and Ti as
is demonstrated below with the in situ scanning tunnelling microscope.
The absorption of species from the atmosphere is common to all electrolyte
solutions and clearly the absorption of water is the biggest issue. This is not solely
confined to ionic liquids, however, as all electroplaters who deal with aqueous
solutions of acids know, if the solution is not heated then the tank will overflow
from absorption of atmospheric moisture after some time. In the aqueous acid
the inclusion of water is not a major issue as it does not significantly affect the
current efficiency or potential window of the solution. Water absorption is also
not such a serious issue with eutectic-based ionic liquids and the strong Lewis
acids and bases strongly coordinate the water molecules in solution. The presence
of up to 1 wt.% water can be tolerated by most eutectic-based systems. Far from
having a deleterious effect, water is often beneficial to eutectic-based liquids as it
decreases the viscosity, increases the conductivity and can improve the rate of the
anodic reaction allowing better surface finishes. Water can even be tolerated in the
chloroaluminate liquids to a certain extent [139] and it was recently shown that the
presence of trace HCl, that results from hydrolysis of the liquid, is beneficial for
the removal of oxide from the aluminum anode [140].
In ionic liquids with discrete anions the presence of water is often far from ideal.
The lack of a Lewis or Brønsted acid to coordinate the water results in a high activity
338 11 Technical Aspects
of water molecules in the liquid. Accordingly the presence of only traces of water
can seriously limit the potential window of the ionic liquid.
No concerted studies on the breakdown products of ionic liquids have been
carried out. It is unlikely that these will interfere with the metal species formed,
but tests need to be carried out at typical current densities that would be used for
commercial plating procedures.
11.5.3
Impact of Impurities on Electrochemistry
Figure 11.22 shows a typical cyclic voltammogram of ultrapure 1-butyl-1methylpyrrolidinium bis(trifluoromethylsulfonyl)amide ([Py1,4 ] TFSA) on Au(111)
[143]. Ultrapurity means that the supplier (in the present case: Merck KGaA/EMD)
guarantees that water and halide impurities are below the 10 ppm level. Routinely
the liquids are dried under vacuum and at elevated temperature to water contents
below 3 ppm (the detection limit of the Karl–Fischer method) prior to use in our
laboratory.
As shown in Figure 11.22, the electrochemical window of this liquid on Au(111)
can be determined by extrapolation of the rising cathodic and the rising anodic
currents to zero. The cathodic limit is mainly due to the irreversible reduction
of the [Py1,4 ] to N-methylpyrrolidine and butyl radicals which undergo further
Fig. 11.22 Cyclic voltammogram of ultrapure [Py1,4 ]TFSA on Au(111)
with v = 10 mV s−1 . The electrochemical window is 5.6 V. The reduction peaks C1–C3 are correlated with TFSA breakdown which may be
induced by ultralow amounts of water or other impurities.
11.5 Impurities 339
Fig. 11.23 The in situ STM picture evidences that the Au(111) surface
near the open circuit potential shows a wormlike surface pattern, likely
due to interference between the gold surface and the [Py1,4 ] cation.
decomposition to butene(s) and hydrogen. The oxidation wave A4 is directly correlated with the breakdown of the cation C4. The anodic limit is due to gold
disintegration and partly to irreversible anion oxidation. The potential regime in
between is the maximum available electrochemical window and can be determined
to be 5.6 V. But what about the peaks/waves C1–C3 and A3 ? Macroscopically there
is no surface modification visible. In the potential regime between +2 and –3 V vs.
ferrocene/ferrocinium (Fc/Fc+ ) gold looks – macroscopically – like gold should.
The quartz crystal microbalance does not show any mass effect in this potential
regime. One could argue about soluble organic or inorganic impurities, but the
liquids contain negligible inorganic ions and, furthermore, purified by chromatography to remove organic impurities, thus high impurity concentrations cannot be
the reason. Figure 11.23 shows the surface of Au(111) under [Py1,4 ] TFSA at the
open circuit potential, i.e. at around –0.5 V vs. Fc/Fc+ .
It is quite interesting that gold does not show here the typical surface known
in aqueous electrochemistry with flat terraces separated by steps, it is, in contrast,
strongly structured with a wormlike pattern. Interactions of the gold surface with
ions of the ionic liquid lead to such restructuring phenomena. If the electrode
potential is reduced successively to –1.7 V the wormlike pattern disappears slowly.
There is a potential regime where only vacancy islands are observed, Figure 11.24(a).
On this picture the transformation is not yet complete (still some vacancy islands
are present), but on a time-scale of 20–30 min these vacancy islands disappear
completely, i.e. in the potential regime of wave C1. Thus the peak C1 could, at
first glance, be correlated to the restructuring of the gold surface. But, on the one
hand, the respective oxidation peak is missing, on the other hand MacFarlane et al.
have described that in this potential regime the irreversible breakdown of the TFSA
ion starts [140]. New results from the Passerini group make it likely that quite low
amounts of water or other impurities, below the 10 ppm level, are involved in the
TFSA decomposition reaction [144].
340 11 Technical Aspects
Fig. 11.24 The quality of the STM picture is reduced strongly in the
potential regime of the reduction peaks C1–C3 in [Py1,4 ]TFSA. The
STM probes the breakdown of the TFSA, which may be induced by
water and/or other impurities in the ultralow concentration regime.
If we perform the same experiment with ultrapure 1-ethyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)amide ([EMIM] TFSA), we do not see the same restructuring with the in situ STM, although it is described that it is the anion which is
subject to irreversible breakdown in this potential regime. This led us to the conclusion that adsorption of the [Py1,4 ] cation (maybe together with TFSA + TFSA
breakdown products) is responsible for the wormlike pattern and the formation of
a flat surface finally. As shown in Chapter 7 the impact of pyrrolidinium ions on
the grain size of electrochemically made metals is evident and pyrrolidinium ions
might act as surface-active species which would open a novel concept of additives for
ionic liquids. Between –1.7 and –1.9 V vs. Fc/Fc+ a flat gold surface can be probed
and around –2 V (i.e. in the potential regime of wave C2) we observed routinely
that the picture quality got worse, see Figure 11.24(b). In the first experiments we
thought this would be due to a bad tip, a common problem the experimentalist has
to struggle with, but it was surprising that the noise shown in Figure 11.24(b) disappeared again when the electrode potential was set back to –1.7 V and reappeared
at about –2 V. Such a reversible and reproducible behavior excludes a “bad tip”. If
11.5 Impurities 341
in the in situ STM experiment the electrode potential is set to values between –2.2
and –2.7 V it is evident that the picture quality is dramatically reduced, as shown
in Figure 11.24(c). It should be mentioned clearly that this is definitely not due to
a bad tip, and there is still a tunnelling contact between tip and surface allowing
one to probe the surface. At –2.9 V (in the potential regime of wave C3), Figure
11.24(d), the surface is now obviously covered by a film which makes probing the
surface difficult. Nevertheless, the steps can still be identified. At lower electrode
potentials, i.e. in the regime of the cathodic breakdown C4, the tunnelling contact
is finally lost. Together with the results from the MacFarlane group it can be concluded that the TFSA is subject to a certain cathodic breakdown, possibly induced
by ultralow amounts of impurities [144] although this cannot be confirmed at the
moment. In any case, with the in situ STM this breakdown can be probed. It can
be concluded from these in situ STM experiments that the understanding of even
apparently simple electrochemical windows of ionic liquids can be a tough job. In
our experience it is not sufficient to have just a look at the cyclic voltammograms,
it may first be necessary to acquire fundamental local probe information of any
ionic liquid which is to be employed for fundamental studies at the electrode/ionic
liquid interface. Furthermore there are known combined anion/cation effects in
ionic liquids that affect the chemical and electrochemical processes so that it may
not be sufficient just to exchange the anion for a more stable one. The surface
behavior may be different. Fundamental local probe electrochemistry might first
require a detailed characterization of the potential-dependent interface effects. It
should be mentioned that inorganic impurities in ionic liquids (even in apparently ultrapure liquids) may lead to a complete misunderstanding of the surface
processes.
We have discussed in Ref. [143] that liquids, made by a metathesis reaction from
[Py1,4 ]Cl and Li-TFSA, can contain low amounts of Li ions. The results will be briefly
summarized here. The ionic liquid is apparently perfect on glassy carbon but there
is clear evidence for deposition of Li on Au(111), Figure 11.25(a). From the cyclic
voltammogram the redox process C5 implies electrodeposition in the UPD regime
whereas the peak C6 would imply bulk deposition. Furthermore, the irreversible
decomposition peaks C1–C3, discussed above, are not observed here. The in situ
investigation of Au(111) in this liquid shows, for an electrode potential of –1.6 V
quite a nice gold surface, Figure 11.25(b), whereas at –2.4 V, Figure 11.25(c), a
deposit is clearly observed. If we add LiTFSA to a [Py1,4 ]TFSA liquid made by the
metathesis routine we get more or less the same result, furthermore Li can be
deposited reproducibly (see Chapter 12).
Our clear recommendation is to have a critical look at any ionic liquid delivered by
any of the suppliers. Glassy carbon is a bad substrate to probe inorganic impurities
in ionic liquids. In our experience noble metal electrodes like gold or platinum
are better suited to detect low amounts of impurities. We have further examples in
Clausthal where we saw clearly, with the in situ STM, metal deposition in apparently
ultrapure liquids. In several liquids we found, with XPS and/or EDX, considerable
amounts of potassium. Upon our insistence, the supplier finally told us that the
synthesis route had been changed. For technical experiments such low amounts of
342 11 Technical Aspects
Fig. 11.25 (a) Cyclic voltammogram of [Py1,4 ]TFSA (with low amounts
of Li impurities) on Au(111): the processes C5 and C6 are typical for
deposition in the UPD and OPD regimes, respectively. (b) In situ STM
of Au(111) before C5: flat gold surface. (c) In situ STM of Au(111)
between C5 and C6: obviously Li deposition has occurred.
impurities maybe do not matter, but for fundamental studies such impurities can
be a nightmare, making progress a bit slow.
A further error in IL synthesis can originate from purification processes. In
order to remove the often yellowish color of ionic liquids after synthesis they are
commonly purified over silica or alumina powder (see above). Once we obtained
a liquid where the supplier invested a lot of effort to deliver “Endres-quality”.
[EMIM]TFSA was made with the best available educts in the acid–base routine
from diluted aqueous [EMIM]OH and H-TFSA. This approach excludes metal and
halide impurities. The supplier removed the slight yellowish color by purification
over silica. For this purpose the supplier used quite a fresh silica, which had
not been used in any purification process before. One has to bear in mind that
the dominant impurities, even in hiqh quality silica, are aluminum species. Figure
11.26 shows the 1st, the 7th and the 15th cycles of this liquid on Au(111). Apparently
11.5 Impurities 343
Fig. 11.26 Cyclic voltammograms (10 mV s−1 ) of [EMI]TFSA (purified
over silica) on Au(111), (a) 1st , (b) 7th , (c) 15th cycle.
perfect in the beginning, deposition occurs in the 15th cycle, C1. In the in situ STM
experiment it was quite surprising that we got a bad surface, Figure 11.27(a). The
less experienced STM-experimentalist would conclude that he had a bad tip. Quite
interestingly, the STM picture quality is improved tremendously by reducing the
electrode potential to –0.5 and subsequently to –1.8 V (Figure 11.27(b) and (c)). At
–1.8 V there is a nice terraced surface. From unpublished XPS data and from the
height probed by in situ STM we have no other explanation but that a tiny, only a
few nanometers high, aluminum layer was deposited, which would not be visible
with a less sensitive analytical tool. The only explanation that we have is that, from
the fresh silica, aluminum species were washed off into the ionic liquid, adsorbed
loosely at the electrode and subsequently electrodeposited. This assumption was
recently proved independently by one of us (D.R.M.). We have further results
that show that impurities can strongly alter chemical reactions. The hydrolysis of
TiCl4 by water in liquids with the TFSA anion only delivers nanocrystalline rutile
if the liquids are ultrapure, i.e. with impurities below the 10 ppm level. Already
impurity levels in the 100 ppm regime lead to a mixture of nano-anatase and
nano-rutile [145].
As a consequence, in our laboratories in Clausthal, any newly delivered ionic liquid is first tested by cyclic voltammetry and in situ STM on Au(111) thoroughly before it is used for fundamental studies. This approach is somewhat time-consuming
and in part frustrating for the students, on the other hand it is currently the only
chance to avoid misinterpretation of electrochemical experiments, especially with
the in situ STM. This is one of the challenges in ionic liquids electrochemistry.
344 11 Technical Aspects
Fig. 11.27 In situ STM pictures of [EMIM]TFSA (purified over silica)
on Au(111): obviously there is a surface film adsorbed initially which
can be reduced to a nice terrace-like deposit.
This chapter draws attention to impurities in the ultralow concentration regime.
We do not say that what we have described here will occur all the time in all
liquids, but it can occur and, therefore, in our opinion, care has to be taken not to
misinterpret results. The suppliers should take more care in the characterization
of ionic liquids.
Appendix A
Protocol for the Deposition of Zinc from a Type III Ionic Liquid
Preparation of Ionic Liquids
The eutectic mixture is formed by stirring the two components, 1 ChCl : 1.5 urea (for
example: 69.815 g ChCl and 45.045 g urea) at 75 ◦ C until a homogeneous colorless
liquid is formed. Once the liquid has formed add 34.07 g ZnCl2 and stir until it is
dissolved. The ratio of the three components is 1 ChCl : 1.5 urea : 0.5 ZnCl2 .
References
Electroplating Experiment
Method
Pour each of the ionic liquids into a 0.25 l beaker. Heat for 20 min with magnetic
stirring to 60 ◦ C. To one of the beakers add 5 wt.% LiCl, and to a second add 10 wt.%
ethanol.
Immerse copper workpiece in HCl for 2 min, then rinse with deionized water
and dry. Place in dichloromethane (to degrease the surface) for 2 min then remove
using tweezers and allow the dichloromethane to evaporate. Attach crocodile clips
to leads on the power supply. Position the two electrodes opposite each other in the
solution and clip them in place with crocodile clips. Attach the copper workpiece to
the positive terminal and use either glassy carbon or zinc attached to the negative
terminal. Perform an anodic etch at 1.0 V for 40 s. Then swap the leads on the
power supply to make the copper workpiece the cathode and the glassy carbon or
zinc the anode. Next begin plating by setting the current density to 20 mA cm−2 .
Leave for 1 h. When the experiment is finished, rinse the copper workpiece with
deionised water and allow to dry.
Safety Precautions
Ĺ Choline chloride (HOCH2 CH2 N(CH3 )3 Cl) – Incompatible with strong oxidizing
agents, moisture. Store under a dry atmosphere. May act as an irritant. Minimize
contact. Risk phrases: R36/37/38, irritating to eyes, respiratory system and skin.
Safety phrases: S26-36. In case of contact with eyes, rinse immediately with water
and seek medical advice. In case of contact with skin, wash immediately with
water. Wear safety glasses and protective clothing/gloves.
Ĺ Urea (H2 NC=ONH2 ) – Irritating to eyes, respiratory system and skin. In case
of contact with eyes, rinse immediately with plenty of water and seek medical
advice. Wear suitable protective clothing. Risk phrases: R36 R37 R38 R40. Safety
phrases: S26 S36.
Ĺ Zinc Chloride (ZnCl2 ) – Incompatible with potassium. Corrosive. Causes burns.
Harmful if swallowed or inhaled and in contact with skin. Irritant, risk of serious
damage to eyes. In case of contact with eyes, rinse immediately with plenty of
water and seek medical advice. Risk phrases: R34-50/53. Wear Safety glasses, use
with adequate ventilation and wear protective clothing. Safety phrases: S7/8-2628-36-45-60-61.
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Earle, M., Gordon, C., Plechkova, M.,
Seddon, K., and Welton, T. (2007)
Decolorization of ionic liquids for
spectroscopy. Anal. Chem., 79, 758–
764.
Li, Z., Du, Z., Gu, Y., Zhu, L., Zhang, X.,
and Deng, Y. (2006) Environmentally
friendly and effective removal of Br- and
Cl- impurities in hydrophilic ionic
liquids by electrolysis and reaction.
Electrochem. Commun., 8, 1270–1274.
Yang, Q. and Dionysiou, D. (2005) Room
temperature ionic liquids as solvent
media for the photolytic degradation of
environmentally important organic
contaminants. Ionic liquids IIB:
Fundamentals, Progress, Challenges
and Opportunities: Transformations and
Processes. ACS Symp. Ser., 902,
182–198.
Gan, Q., Xue, M., and Rooney, D. (2006)
A study of fluid properties and
microfiltration characteristics of room
temperature ionic liquids
[C10-min][NTf2] and N8881[NTf2] and
their polar solvent mixtures. Sep. Purif.
Technol., 51 (2), 185–192.
Kröckel, J. and Kragl, U. (2003)
Nanofiltration for the separation of
non-volatile products from solutions
References
containing ionic liquids. Chem. Eng.
Technol., 26, 1166–1168.
137 Moustafa, E., Zein El Abedin, S.,
Shkurankov, A., Zschippang, E., Saad,
A., Bund, A., and Endres, F. (2007)
Electrodeposition of Al in
1-Butyl-1-methylpyrrolidinium
Bis(trifluoromethylsulfonyl)amide and
1-Ethyl-3-methylimidazolium
Bis(trifluoromethylsulfonyl)amide Ionic
Liquids: In Situ STM and EQCM
Studies. J. Phys. Chem. B, 111,
4693–4704.
138 Zein El Abedin, S., Moustafa, E., Natter,
H., Hempelmann, R., and Endres, F.
(2005) Additive free electrodeposition of
nanocrystalline aluminium in a water
and air stable ionic liquid. Electrochem.
Commun., 7, 1116.
139 Seddon, Kenneth R., Stark, Annegret,
140
141
142
143
144
145
and Torres, M-J. (2000) Pure Appl.
Chem., 72 (12), 2275–2228.
Howlett, P.C., Izgorodina, E.I., Forsyth,
M., and MacFarlane, R. (2006) Z. Phys.
Chem., 220, 1483.
Abbott, A.P., Eardley, C.A., Farley, N.S.,
and Pratt, A. (1999) Trans. Inst. Metal
Finish., 77, 26.
Abbott, A.P., Qiu, F., and Ryder, K.S. (in
press).
Endres, F., Zein El Abedin, S., and
Borissenko, N. (2006) Z. Phys. Chem.,
220, 1377.
Passerini, S., lecture on the COIL-2
presymposium, 4 August 2007
Yokohama, Japan
Kaper, H., Endres, F., Djerdj, I.,
Antonietti, M., Smarsly, B., Maier, J.,
and Hu, Y.-S., Small, DOI:
10.1002/smll.200700138
351
353
12
Plating Protocols
Frank Endres, Sherif Zein El Abedin, Q. Liu, Douglas R. MacFarlane,
Karl S. Ryder, and Andrew P. Abbott
In this chapter we would like to present some plating protocols for the electrodeposition of aluminum, lithium, tantalum and zinc from different ionic liquids. These
“recipes” have been elaborated in our laboratories and should allow the beginner
to perform his first electrodeposition experiments. For aluminum we give four
different recipes in order to show that the ionic liquid itself can strongly influence
the deposition of metals. In the case of tantalum the deposition of the metallic
phase is not straightforward as, in unstirred solutions, the more nonstoichiometric
tantalum halides form the higher the current density for electrodeposition. Apart
from the zinc deposition all experiments should be performed at least under dry
air.
12.1
Electrodeposition of Al from 1-Ethyl-3-methylimidazolium chloride/AlCl3
In this protocol we describe an electroplating procedure for mild steel with an
adhesive aluminum layer in Lewis acidic ionic liquid 1-ethyl-3-methylimidazolium
chloride [EMIM]Cl containing AlCl3 . We aim to electroplate mild steel with dense,
adherent and uniform aluminum layers in the employed ionic liquids at room
temperature.
12.1.1
Experimental Set-up
In the designed electrochemical cell, Al sheets (Alfa, 99.999%) machined into a
cylinder configuration were used as reference and counter electrodes. Mild steel
sheets were employed as working electrodes. Prior to use, the mild steel sheets were
mechanically polished with emery paper, cleaned with acetone in an ultrasonic bath,
treated with dilute hydrochloric acid and rinsed with distilled water. The mild steel
sheets were always anodically polarized in the employed ionic liquid immediately
before the electrodeposition in order to remove as far as possible the inevitable
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
354 12 Plating Protocols
oxide layer. This pre-run anodic polarization, termed as in situ etching, has been
proved to be a vital prerequisite for the formation of adherent coatings. The cell was
thoroughly cleaned in boiling double-distilled water mixed with 10% (in volume
ratio) H2 O2 . The deposition experiments were performed in an argon-filled glove
box using a Parstat 2263 Potentiostat/Galvanostat (Princeton Applied Research).
12.1.2
Chemicals and Preparation
The commercial available 1-ethyl-3-methylimidazolium chloride [EMIM]Cl high
purity quality of was dried under high vacuum for 24 h at a temperature of 60 ◦ C
then transferred into an argon-filled glove box with water and oxygen below 1 ppm
(OMNI-LAB from Vacuum-Atmospheres). Anhydrous AlCl3 (Fluka, 99%) was used
without further purification as a source of aluminum. It is recommended strongly
that AlCl3 grains are used, as powders (even in 99.999% quality) only contain
low amounts of active AlCl3 , according to our experience (due to their high surface
powders rapidly absorb water, leading to a phalanx of different aluminumoxohalides
which are difficult to reduce).
The mixing of AlCl3 with [EMIM]Cl is a violent exothermic reaction accompanied
by a sharp temperature increase. The addition of AlCl3 to [EMIM]Cl should be
carried out carefully. The molar ratio of both chemicals is 2:3, with AlCl3 in excess.
The obtained liquid [EMIM]Cl/AlCl3 is yellowish and of satisfactory fluidity at room
temperature.
12.1.3
Results
The SEM micrograph of Figure 12.1 shows the surface morphology of an electroplated aluminum layer obtained at a current density of −20 mA cm−2 for 2 h in
Lewis acidic [EMIM]Cl/AlCl3 at room temperature on mild steel. As can be seen,
the obtained Al deposit consists of coarse crystallites forming rather a compact
layer without observable cracks.
Figure 12.2 presents the optical view of the polished cross-section of Al layers
made at −20 mA cm−2 on a mild steel substrate without performing in situ etching
prior to the electroplating. An interstice between the substrate and the electroplated
layer is seen due to a thin oxide layer which forms after the last pre-treatment step.
Figure 12.3 shows the optical view of the polished cross-section of Al layers made
at −20 mA cm−2 on a mild steel substrate. In contrast to the previous sample, an
anodic polarization at 1.0 V was applied to the working electrode for around 2 min
before electrodeposition. The layer adhesion is significantly improved.
The importance of an in situ etching process in making an adhesive Al layer in
ionic liquid on mild steel has been discussed in detail in one of our publications
[1].
12.1 Electrodeposition of Al from 1-Ethyl-3-methylimidazolium chloride/AlCl3
Fig. 12.1 SEM micrograph of an Al layer electroplated on mild steel
substrate at −20 mA cm−2 .
Fig. 12.2 The optical view of the cross-section for an Al layer electroplated on mild steel substrate at −20 mA cm−2 .
Fig. 12.3 The optical view of the cross-section for an Al layer electroplated on mild steel substrate at −20 mA cm−2 with significant adherence improvement by in situ electrochemical etching.
355
356 12 Plating Protocols
Fig. 12.4 An optical photo of deposits manufactured from the employed ionic liquid with uneven surface and screw geometry.
The employed liquid can also be used to electroplate Al on an uneven surface
or even a screw, as shown in Figure 12.4. The in situ etching leads to Al layers all
showing satisfactory adhesive qualities.
12.2
Electrodeposition of Al from 1-Butyl-3-methylimidazolium
chloride–AlCl3 – Toluene
12.2.1
Apparatus, Materials and Chemicals
Potentiostat, (typically, Echochemie, Autolab PGSTAT), Schlenk tube × 2, aluminum sheet, mild steel rods, P400 sand paper, anhydrous AlCl3 , 1-butyl-3-methyl
imidazolium chloride ([BMIM]Cl), toluene, acetone, dichloromethane, HCl, HNO3 ,
H3 PO4 , acetic acid, isopropanol.
12.2.2
Preparation of AlCl3 –[BMIM]Cl–Toluene Ionic Liquid Mixture ([2:1]:3)
Aluminum chloride (Aldrich >99%) and 1-butyl-3-methyl imidazolium chloride
([BMIM]Cl) (Aldrich >99%) were weighed at a 2:1 mole ratio into two separate
Schlenk tubes and dried on a vacuum line for 2 h prior to use. The two components were mixed by adding AlCl3 to the [BMIM]Cl tube with stirring at room
temperature, leading to a homogeneous, straw-brown liquid of “technical quality”. Finally, 3 mole equivalents of toluene (corresponding to 39 wt.%) were then
transferred into the 2:1 neat ionic liquid using a stainless steel cannula. A homogeneous mixture (dark green in color) was obtained by stirring the liquid
for 15 min. The liquid was maintained under a dry nitrogen atmosphere at all
times.
12.2 Electrodeposition of Al from 1-Butyl-3-methylimidazolium chloride–AlCl3 – Toluene 357
12.2.3
Pretreatments
To achieve a good adhesive coating and maintain the electrolyte stability, both
cathode and anode need to be treated properly before they are mantled for electroplating.
12.2.3.1 Cathode (Mild Steel Rods)
Ĺ Polished using P400 sand paper and cleaned using tissues.
Ĺ Degreased in acetone under ultrasonic conditions for 15 min.
Ĺ Activated chemically in 5 wt.% HCl for 2 min to remove possible oxide layer then
rinsed thoroughly using deionized water.
Ĺ Degreased in dichloromethane for 10 min to remove any organic impurities and
form a chloride layer which is resistant to oxide formation.
12.2.3.2 Anode (Al)
The anode was polished using P400 sand paper, and then activated by dipping in
(1% HNO3 , 65% H3 PO4 , 5% acetic acid and water) for 5 min, followed by rinsing
thoroughly with deionized water and degreasing in acetone for 5 min.
12.2.4
Electroplating and Morphology Analysis
Electroplating experiments were performed using a two-electrode set-up under N2
atmosphere in a Schlenk tube. The cathodes were mild steel rods with diameter
0.6 cm, and the anode was a cylindrical bucket of Al sheet of diameter 1.6 cm placed
around the cathode. Anodic etching of mild steel rods was performed by applying
+1 V for 30 s to remove any possible oxide layer prior to electroplating. All samples
were prepared by applying constant potentials for 60 min. Samples were rinsed
using toluene followed by isopropanol and then deionized water after removal
from the Schlenk tube. Surface analysis was carried out using scanning electron
microscopy (Philips XL30 ESEM) and energy dispersive analysis by X-rays (EDX).
12.2.5
Results
The surface morphologies of the deposits are highly dependent on the potential
applied between the anode and cathode. For lower voltages (<0.5 V), the deposits
tend to grow to a bigger crystal, which gives a dull finish; for higher voltages
(>0.5 V), a growth of nanocrystals dominates, leading to smooth bright and shining
samples. Figure 12.5 shows photos of the samples and Figure 12.6 SEM images of
the same samples.
358 12 Plating Protocols
Fig. 12.5 Photos showing (a) the dull finish at 0.5 V and (b) the bright finish at 1.0 V.
12.3
Electrodeposition of Al from 1-Ethyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)amide/AlCl3
In this protocol we describe the electroplating of mild steel with thick layers of
aluminum in the water and air stable ionic liquid 1-ethyl-3-methylimidazolium
bis(trifluoromethylsulfonyl) amide [EMIM]TFSA containing AlCl3 . We aim to electroplate mild steel with dense, adherent aluminum layers in the employed ionic
liquids.
12.3.1
Experimental Set-up
A quartz round flask was used as an electrochemical cell with three electrodes.
Al-wires (Alfa, 99.999%) were used as reference and counter electrodes. Mild steel
sheets were employed as working electrodes. The working electrodes were mechanically polished with emery paper, cleaned with acetone in an ultrasonic bath,
treated with dilute hydrochloric acid and rinsed with distilled water. Prior to the
electrodeposition process the electrodes were anodically polarized in the employed
ionic liquid to remove as far as possible the native oxide layer. Removal of the
Fig. 12.6 SEM images of (a) the dull finish at 0.5 V and (b) the bright finish at 1.0 V.
12.3 Electrodeposition of Al from 1-Ethyl-3-methylimidazolium 359
surface air-formed oxide layer is a prerequisite for achieving adherent coatings.
The cell was thoroughly cleaned in a mixture of 50/50 vol% H2 SO4 /H2 O2 followed
by refluxing in bi-distilled water. The deposition experiments were performed in
an argon-filled glove box using a Parstat 2263 Potentiostat/Galvanostat (Princeton
Applied Research).
12.3.2
Chemicals and Preparation
The ionic liquid [EMIM]TFSA was purchased from Merck KGaA(EMD) in the
highest available quality and was dried under vacuum for 12 h at a temperature
of 100 ◦ C then stored in an argon-filled glove box with water and oxygen below
1 ppm (OMNI-LAB from Vacuum-Atmospheres). Anhydrous AlCl3 (Fluka, 99%)
was used without further purification as a source of aluminum. It is important that
AlCl3 grains are employed, as powders (even in 99.999% quality) only contain low
amounts of active AlCl3 , according to our experience.
The ionic liquid [EMIM]TFSA shows, at room temperature, biphasic behavior
on addition of AlCl3 . AlCl3 dissolves well in [EMIM] TFSA up to a concentration of
about 2.5 mol L−1 , then a biphasic mixture is obtained on further addition of AlCl3 .
The upper phase of the mixture AlCl3 /[EMIM] TFSA is clear and colorless while the
lower one is pale and more viscous. On further addition of AlCl3 the viscosity of the
lower phase increases. It is worth noting that Al can only be electrodeposited from
the upper phase, the clear one, at AlCl3 concentrations ≥ 5 mol L−1 . Furthermore,
after a few days a precipitate which contains Al(TFSA)3 forms as a third phase.
12.3.3
Results
The SEM micrograph of Figure 12.7(a) shows the surface morphology of a deposited
aluminum layer obtained galvanostatically at a current density of −5 mA cm−2 for
2 h in the upper phase of the biphasic mixture [EMIM] TFSA/6 M AlCl3 at room
temperature. Prior to Al electrodeposition, the electrode was anodically polarized
at a potential of 1 V (vs. Al) for 2 min. As seen, the deposited Al layer is dense and
contains crystallites in the micrometer regime.
Figure 12.7(b) shows SEM micrographs of the cross-section of the deposited
aluminum layer on a mild steel substrate. As shown in the SEM micrograph
the deposited Al layer adheres well to the mild steel substrate and the layer is
homogeneous with a thickness of about 10 µm. Also, with higher magnification,
the Al layer exhibits a good adhesion without any splits between it and the substrate,
inset of Figure 12.7(b).
Figure 12.8 shows the photo of a deposited aluminum layer obtained potentiostatically on a mild steel substrate at −0.3 V (vs. Al) for 4 h in the upper phase of
the mixture [EMIM] TFSA/6 M AlCl3 . The substrate was electrochomically etched
at 1 V (vs. Al) for 2 min prior to electrodeposition. The aluminium layer adheres so
well that it can be mechanically polished to a mirror appearance.
360 12 Plating Protocols
Fig. 12.7 (a) SEM micrograph of an about 10 :m aluminum layer
electrodeposited galvanostatically on a mild steel substrate at
−5 mA cm−2 . Inset: SEM micrograph of higher magnification showing the excellence of the coating adhesion. (b) SEM micrograph of the
polished cross-section of the deposited aluminium layer [2].
12.4
Electrodeposition of Al from 1-Butyl-1-methylpyrrolidinium
bis(trifluoromethylsulfonyl)amide/AlCl3
In this protocol we describe the electrodeposition of nanocrystalline aluminum without additives in the water- and air-stable ionic liquid 1-butyl-1methylpyrrolidinium bis(trifluoromethylsulfonyl) amide [Py1,4 ]TFSA containing
AlCl3 .
12.4.1
Experimental Set-up
The experimental set-up used was as described in Section 12.3.1. Gold substrates
from Arrandee (gold films of 200–300 nm thickness deposited on chromiumcovered borosilicate glass) and glassy carbon (Alfa) and mild steel sheets were
used as working electrodes, respectively. Directly before use, the gold substrates
were heated in a hydrogen flame to slightly red glow for several minutes. The glassy
carbon substrate was degreased with acetone in an ultrasonic bath for 10 min. The
mild steel substrates were mechanically polished with emery paper, cleaned with
Fig. 12.8 An optical photo of a deposited Al layer made potentiostatically at −0.3 V (vs. Al) in the upper phase of the mixture [EMIM]
TFSA/6 M AlCl3 at room temperature.
12.4 Electrodeposition of Al from 1-Butyl-1-methylpyrrolidinium 361
acetone in an ultrasonic bath, treated with dilute hydrochloric acid and rinsed with
distilled water.
12.4.2
Chemicals and Preparation
The ionic liquid 1-butyl-1-methylpyrrolidinium bis (trifluoromethylsulfonyl)amide
was purchased from Merck KGaA(EMD) in the highest available quality and was
dried under vacuum for 12 h at a temperature of 100 ◦ C then stored in an argonfilled glove box with water and oxygen below 1 ppm (OMNI-LAB from VacuumAtmospheres). Anhydrous AlCl3 (Fluka, 99%) was used without further purification
as a source of aluminum.
Similar to the AlCl3 /[EMIM]TFSA mixture, the mixture of AlCl3 /[Py1,4 ] TFSA
shows biphasic behavior with increase in the concentration of AlCl3 up to 1.6 M.
In contrast to the AlCl3 /[EMIM]TFSA mixture, the lower phase is colorless while
the upper one is pale and more viscous. By adding more AlCl3 the volume of
the lower phase decreases till a concentration of 2.7 mol L−1 is reached, then only
one solid phase can be formed at room temperature. The biphasic mixture of
AlCl3 /[Py1,4 ]TFSA becomes monophasic by heating to a temperature of about 80 ◦ C.
The electrodeposition of aluminum occurs only from the upper phase at AlCl3
concentrations ≥ 1.6 mol L−1 .
12.4.3
Results
Nanocrystalline aluminum can be made in the employed ionic liquid without
additives, see Chapter 8. The SEM micrograph of Figure 12.9 shows the surface
morphology of a deposited aluminum layer obtained potentiostatically on mild steel
at −0.75 V (vs. Al) for 2 h in the upper phase of the biphasic mixture [Py1,4 ] Tf2 N/2 M
AlCl3 at 100 ◦ C. Prior to Al electrodeposition, the electrode was anodically polarized
at a potential of 1 V (vs. Al) for 2 min. The deposited layer is dense, shining and
adherent to the substrate with crystallites in the nanosize regime.
Figure 12.10 shows a high resolution SEM micrograph of an about 5 µm thick
layer of Al on gold substrate electrodeposited potentiostatically at 100 ◦ C at −0.45 V
(vs. Al) for 2 h in the upper phase of the mixture [Py1,4 ]TFSA/1.6 M AlCl3 . Generally,
the electrodeposited layer contains very fine crystallites in the nanometer regime.
Figure 12.11 shows the XRD patterns of a nanocrystalline Al film obtained at
a constant potential of −1.7 V for 2 h at 100 ◦ C in the ionic liquid [Py1,4 ] TFSA
containing 1.6 M AlCl3 on a glassy carbon substrate. The XRD patterns show the
characteristic diffraction patterns of crystalline Al, furthermore the peaks are rather
broad, indicating the small crystallite size of the electrodeposited Al. The grain size
of Al was determined using Scherrer’s equation to be 34 nm. For more information
on the electrodeposition of nanocrystalline aluminum in the employed ionic liquid
we refer to Refs. [3, 4].
362 12 Plating Protocols
Fig. 12.9 SEM micrograph of electrodeposited Al on mild steel made
potentiostatically at −0.75 V (vs. Al) for 2 h in the upper phase of the
mixture [Py1,4 ]TFSA/2 M AlCl3 at 100 ◦ C.
12.5
Electrodeposition of Li from 1-Butyl-1-methylpyrrolidinium
bis(trifluoromethylsulfonyl)amide/Lithium bis(trifluoromethylsulfonyl)amide
A solution of lithium bis(trifluoromethylsulfonyl)amide is made up at
∼0.5 mol kg−1 in the ionic liquid 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide. Both the salt and the ionic liquid are dried prior to use
at 100 ◦ C or above, under vacuum for 12 h or more, which gives water values at
least below 10 ppm. These materials must be handled only in an argon-filled dry
Fig. 12.10 SEM micrograph of electrodeposited Al on gold formed
potentiostatically at −0.45 V (vs. Al) for 2 h in the upper phase of the
mixture [Py1,4 ]TFSA/1.6 M AlCl3 ) at 100 ◦ C.
12.5 Electrodeposition of Li from 1-Butyl-1-methylpyrrolidinium 363
Fig. 12.11 XRD patterns of an electrodeposited Al layer obtained potentiostatically at −1.7 V for 2 h in the upper phase of the mixture
[Py1,4 ]TFSA/1.6 M AlCl3 at 100 ◦ C on a glassy carbon substrate.
box. Since lithium reacts rapidly with both oxygen and nitrogen this experiment
must be carried out under an argon, or other inert gas, atmosphere. To ensure
that the ionic liquid does not contain traces of nitrogen or oxygen the lithium salt
solution in the ionic liquid should be degassed by bubbling pure argon through
the solution overnight; the solution is held at an elevated temperature (>50 ◦ C)
during this process to reduce the viscosity. If available, the water content should be
determined at this stage by Karl–Fischer titration and should be < 30 ppm.
Plating of lithium can be performed galvanostatically with a current density
around 1–2 mA cm−2 on a range of substrates including nickel, copper, platinum and glassy carbon. Nickel and glassy carbon tend to require a greater overpotential to achieve Li deposition. Platinum requires less over-potential, but tends
to form alloys easily with lithium and therefore it is not preferred for a simple plating experiment. In all cases a passive film will form on the lithium deposit as it plates. This surface film that forms in 1-butyl-1-methylpyrrolidinium
bis(trifluoromethylsulfonyl)amide is approximately a few hundred nanometers
thick and is known to be highly conductive to lithium [5]. Lithium plating occurs through the film but at a rate which, to some extent, is limited by the film,
particularly at lower temperatures. Despite the presence of this film, the deposit
will remain bright and shiny.
364 12 Plating Protocols
The counter electrode is preferably lithium metal in order to provide a constant
lithium concentration in the electrolyte. Lithium is also a very useful reference electrode in this ionic liquid in the form of a strip of foil. For preliminary experiments
platinum is a suitable counter electrode.
The electrodeposition can be carried out at room temperature, but is more facile at
50 ◦ C or higher due to the resistance of the passive film. Typically about 50–100 mV
of overpotential vs. Li/Li+ is sufficient to obtain a deposit. It is important to limit
this overpotential to <150 mV because of the reductive instability of the ionic liquid
at more negative potentials. It is advisable therefore to plate under potentiostatic
conditions. The achievable current density is very much dependent on the temperature involved. At 50 ◦ C a good deposit can be obtained at 1–1.5 mA cm−2 . Initiation
of a good uniform film is often achieved by depositing initially at lower current
densities to allow the creation of the passive film before higher current densities
are applied.
Additives such as certain zwitterionic compounds [6] can allow an increase in
the current density achieved.
12.6
Electrodeposition of Ta from 1-Butyl-1-methylpyrrolidinium
bis(trifluoromethylsulfonyl)amide
Electrodeposition of tantalum thin layers in the water- and air-stable ionic liquid
1-butyl-1-methyl pyrrolidinium bis (trifluoromethylsulfonyl) amide at 200 ◦ C using
TaF5 as a source of tantalum is presented in this protocol. The electrodeposition of
Ta is not a straightforward process as, under the wrong experimental conditions.
Ta subhalides can be formed, see Chapter 4.4.
12.6.1
Electrodes
Platinum sheets of thickness 0.5 mm (Alfa, 99.99%) were used as a working electrode. Directly before use, the Pt substrate was cleaned for 10 min in an ultrasonic
bath in acetone then heated in a hydrogen flame to red glow for a few minutes.
Pt-wires (Alfa, 99.99%) were used as reference and counter electrodes, respectively.
A quartz round flask was used as an electrochemical cell. The electrodeposition
experiments were performed in an argon-filled glove box with water and oxygen
below 1 ppm.
12.6.2
Chemicals
The ionic liquid 1-butyl-1-methyl pyrrolidinium bis (trifluoromethylsulfonyl)amide
was purchased from Merck KGaA(EMD) in ultrapure quality. TaF5 (Alfa, 99.99%)
and LiF (Alfa, 98.5%) were used without further purification.
12.7 Electrodeposition of Zinc Coatings from a Choline Chloride 365
Fig. 12.12 (a) SEM micrograph of the electrodeposit formed potentiostatically on Pt in [Py1,4 ]Tf2 N containing 0.25 M TaF5 and 0.25 LiF
at a potential of −1.8 V for 1 h at 200 ◦ C. (b) XRD patterns of the deposited Ta layer.
12.6.3
Results
The electrodeposition of Ta was performed at 200 ◦ C. It was found that the mechanical quality and the adherence of the electrodeposits improved as the temperature
increased. Moreover, the quality and the adherence of the electrodeposit were found
to be improved upon addition of LiF to the electrolyte [7, 8].
During the electrodepsoition of Ta a non-stoichiometric layer of tantalum subhalide(s) on top of tantalum is also formed. This layer can be removed efficiently
by washing with isopropanol followed by boiling in water. After such treatment
only crystalline and elemental Ta can be detected at the electrode surface, about
1 µm in thickness. The SEM micrograph of the Ta electrodeposit, Figure 12.12(a),
made potentiostatically at −1.8 V in the ionic liquid [Py1,4 ]TFSA containing 0.25 M
TaF5 and 0.25 M LiF on Pt at 200 ◦ C for 1 h shows a coherent and dense layer. XRD
patterns of the electrodeposit clearly show the characteristic patterns of crystalline
tantalum, Figure 12.12(b).
12.7
Electrodeposition of Zinc Coatings from a Choline Chloride:
Ethylene Glycol-based Deep Eutectic Solvent
12.7.1
Experimental
Choline chloride (Aldrich >99%), ethylene glycol (Aldrich >99%), ZnCl2 (Aldrich
>99%) and ethylene diamine (Aldrich >99%) were used as obtained. The eutectic
mixture was formed by stirring a 1:2 molar ratio mixture of choline chloride and
366 12 Plating Protocols
ethylene glycol at 70 ◦ C until a homogeneous colorless liquid was formed. ZnCl2
was then dissolved in the liquid. As a further experiment two molar equivalents of
ethylene diamine were added as a complexing agent.
12.7.2
Pretreatment
To attain an adherent Zn Coating the pre-treatment protocol below was followed:
The cathode (mild steel) must be:
Ĺ
Ĺ
Ĺ
Ĺ
Polished using P400 sandpaper, rinsed in deionized water and dried.
Degreased in acetone for 5 min.
Chemically etched in 30% H2 SO4 for 30 s, and rinsed with deionized water.
Degreased in dichloromethane for 5 min to remove organic impurities, rinsed in
deionized water and dried with N2.
Anode (IrOx -coated Ti mesh) must be:
Ĺ Degreased in acetone for 5 min, rinsed thoroughly in deionized water and dried
with N2.
Homogeneous coatings were obtained by driving a constant current density of
5 mA cm−2 at 50 ◦ C, without stirring for 60 min.
12.7.3
Results
The deposition obtained from ZnCl2 dissolved in choline chloride: ethylene glycol
(Figure 12.13(a)) contains small crystals of a homogeneous size. This deposit has
a matte dark gray appearance. The addition of ethylene diamine (Figure 12.13(b))
leads to the growth of larger crystallites and a disperse silvery metallic finish.
Fig. 12.13 Scanning electron micrographs
showing the deposits gained from (a)
choline chloride: ethylene glycol (1:2) +
0.3 M ZnCl2 and (b) choline chloride: ethylene glycol (1:2) + 0.3 M ZnCl2 + 1 molar
equivalent ethylene diamine. Both experiments were carried out by applying a constant current density of 5 mA cm−2 at 50 ◦ C,
without stirring for 60 min.
References
References
1 Liu, Q.X., Zein El Abedin, S., and Endres,
F. (2006) Surf. Coat. Technol., 201 (3–4),
1352.
2 Zein El Abedin, S. (2006) Z. Phys. Chem.,
220, 1293.
3 Zein El Abedin, S., Moustafa, E.M.,
Hempelmann, R., Natter, H., and Endres,
F. (2006) Chem. Phys. Chem., 7,
1535.
4 Zein El Abedin, S., Moustafa, E.M.,
Hempelmann, R., Natter, H., and Endres,
F. (2005) Electrochem. Commun., 7,
1116.
5 Howlett, P.C., Brack, N., Hollenkamp,
A.F., Forsyth, M., and MacFarlane, D.R.
(2006) J. Electrochem. Soc., 153, A595.
6 Tiyapiboonchaiya, C., Pringle, J. M., Sun,
J.Z., Byrne, N., Howlett, P.C., Macfarlane,
D.R., and Forsyth, M. (2004) Nature
Mater., 3, 29.
7 Zein El Abedin, S., Farag, H.K., Moustafa,
E.M., Welz-Bierman, U., and Endres, F.
(2005) Phys. Chem. Chem. Phys., 7, 2333.
8 Zein El Abedin, S., Welz-Bierman, U., and
Endres, F. (2005) Electrochem. Commun., 7,
941.
367
369
13
Future Directions and Challenges
Frank Endres, Andrew P. Abbott, and Douglas MacFarlane
In this book the current state-of-the art of electrodeposition in ionic liquids has been
summarized. In Chapter 2 the key aspects of three types of ionic liquids, i.e. (i) first
generation ionic liquids based on AlCl3 , (ii) air- and water -stable ionic liquids and
(iii) systems based on choline chloride, were introduced. After an introduction to
the physical properties there are four chapters describing the more or less classical
electrodeposition of metals, alloys, semiconductors and conducting polymers. The
subsequent chapter describes rather novel aspects such as the electrodeposition
of nanocrystalline metals and alloys which seems to be quite easy in some ionic
liquids. The in situ scanning tunneling microscope gives direct insight into dynamic
nanoscale processes during electrodeposition and plasma electrochemistry allows
the preparation of suspensions of nanocrystalline metal particles quite simply by
discharging a plasma over the ionic liquid. In the following we discuss possible
future directions and some challenges.
13.1
Impurities
As briefly discussed in Chapters 4.4 and 11.5 ionic liquids can contain variable
amounts of organic and inorganic impurities. The organic impurities, which often
give a yellowish color to some liquids, arise either from impurities in the starting
material or formed during the synthesis by partial decomposition/oligomerization
of the cation and/or the anion. In our experience low levels of such organic impurities are not critical and even with the in situ STM, which is highly sensitive
towards impurities that adsorb at an electrode surface, there is quite good picture
quality, even in yellowish ionic liquids. Thus a low level of organic impurities might
be tolerable for an electrochemical application. Common inorganic impurities in
ionic liquids are water, metal ions and halide(s). Water is introduced during the
synthesis of most ionic liquids as they are typically made either by the acid–base
route from e.g. bis(trifluoromethylsulfonyl)amide-acid and diluted solutions of 1ethyl-3-methylimidazolium hydroxide in water, or via a metathesis reaction from
e.g. lithium bis(trifluoromethyl-sulfonyl)amide and 1-ethyl-3-methylimidazolium
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
370 13 Future Directions and Challenges
chloride in aqueous solution. Water can be easily removed from such liquids simply by stirring them at elevated temperature (about 100 ◦ C) under vacuum. Water
levels of 3 ppm and below are easily achieved.
Metal ion and halide impurities are an issue in ionic liquids with discrete anions.
As we have demonstrated in Chapter 11.5 Li+ (and K+ ) are common cationic impurities, especially in the bis(trifluoromethylsulfonyl)amides which typically contain
100 ppm of these ions from the metathesis reaction. Although Li and K are only
electrodeposited in the bulk phase at electrode potentials close to the decomposition potential of the pyrrolidinium ions, there is evidence for the underpotential
deposition of Li and K on gold and on other rather noble metals. For a technical
process to deposit nickel or cobalt from ionic liquids the codeposition of Li and/or
K, even in the underpotential deposition regime, has to be expected.
Halide impurities can alter the complex chemistry in ionic liquids and can lead
to unexpected oxidation reactions at the counter electrode. Furthermore even low
amounts of e.g. chlorine can be formed, leading to some side reactions.
When ILs were first commercially available, the quality of most samples was
questionable as they contained numerous organic and inorganic impurities. More
recently different quality levels have been introduced (for synthesis, high purity,
ultrapurity). Ultrapure ionic liquids usually contain water, halide and metal ion
impurities below 10 ppm and they are currently the best choice for fundamental
physicochemical studies.
There might be two distinct approaches to the purity issue in future: on the
one hand ultrapure ionic liquids (i.e. impurity levels below 10 ppm) should be
used for fundamental electrochemical studies to understand the electrochemical
reactions alone, which can be quite complicated in ionic liquids. On the other
hand a deep(er) understanding of the mentioned impurities might allow the use of
lower quality ionic liquids for technical electrochemistry or electroplating. Water
impurities might be less critical (if not beneficial) if an element like nickel or
cobalt is deposited. Halide impurities might not be critical for semiconductor
electrodeposition which can be achieved easily from halides. An understanding of
the influence of impurities on the electrochemical processes and information on
the levels that can be tolerated for a reaction would help in the design of technical
processes.
Impurities are a lot less problematic for eutectic-based ionic liquids. The strong
acid–base nature of these systems leads to predominantly halometallate species
which tend to be unaffected by simple salts or other impurities such as water. The
strong Lewis acids and bases coordinate well to water and even in the chloroaluminate systems low amounts of water do not significantly affect voltammetric
behavior or have a deleterious effect on deposit morphology.
13.2
Counter Electrodes/Compartments
In Chapter 11.1 some aspects of counter electrode reactions and metal dissolution were discussed. An interesting aspect in ionic liquid electrochemistry is
that some reactive metals are quite noble. Aluminum, for example, is easily
13.3 Ionic Liquids for Reactive (Nano-)materials 371
oxidized electrochemically in first generation ionic liquids based on AlCl3 , giving a reversible counter electrode. In air- and water-stable ionic liquids with the
bis(trifluoromethylsulfonyl)amide anion, aluminum behaves as a passive electrode.
On the other hand gold is oxidized in both types of ionic liquids. This is not
too surprising: ionic liquids can have wide anodic decomposition potentials (up
to 3 V vs. NHE), wide enough to allow the oxidation of almost all elements. In
contrast to aluminum, gold can be present as “naked” Au+ ions which seems
to facilitate an electrochemical oxidation in the mentioned liquid. In some ionic
liquids platinum (especially in the presence of halide) can be oxidized and deposited on the working electrode if cathodic and anodic compartments are not
separated.
Counter electrode reactions have so far been more or less neglected in ionic liquid
electrochemistry. As there can be unusual reactions, more effort should be invested
in studying these processes. There are suggestions that the counter electrode can
also influence the morphology of deposits at the cathode. In haloaluminate liquids,
for example, although aluminum dissolves, the rate is limited by the diffusion of
AlCl4 − to the electrode surface. The competition of generated Al3+ with AlCl3 for the
halide anion is controlled by the relative Lewis acidity of the ionic liquid components
or, more accurately, of the components in the double layer close to the electrode
surface. Hence, in Lewis acidic ionic liquids the rate of aluminum dissolution is
slower than the rate of deposition and under constant potential the rate is limited
by anodic dissolution. Preliminary results have shown that the increased rate of
deposition and improved quality of the deposit brought about by the addition of
toluene is due primarily to the increase in the rate of the anodic process.
The limited reversibility of some electrode reactions might require consideration
of consumable (cheap) ionic liquids in the anode compartment for technical applications and commercial electroplating. For example, the electrochemical oxidation
of oxalate delivers carbon dioxide, hydride could be oxidized to hydrogen, halides
to the halogen or trihalide salt in the case of iodide ionic liquids and so on . Since
ionic liquids can readily form biphasic systems an alternative may be to have the
anodic reaction in an immiscible solvent. In that case a common ion would be
needed that can be transferred from one phase to the other.
13.3
Ionic Liquids for Reactive (Nano-)materials
The electrodeposition of reactive elements like Al, Si, Ge, Ta and a few others is
possible. As discussed in Chapter 4.4 the successful electrodeposition of Ti, Mg,
Mo and many others in relevant layer thicknesses has not yet been described,
though attempts have been made in some cases. Apart from the availability of
suitable precursors there is at least one other issue to consider: ionic liquids can
be reactive. It was found that magnesium and its alloys can form passivating
films in ionic liquids with the bis(trifluoromethylsulfonyl)amide (Tf2 N) anion, especially in the presence of water. It was found by two of our groups (Endres,
MacFarlane) that, under certain circumstances, the Tf2 N ion is subject to cathodic
372 13 Future Directions and Challenges
breakdown. It is likely that water in the liquid plays an important role in the breakdown reaction. Attempts to deposit magnesium from Mg(Tf2 N)2 have not yet been
successful since, in the presence of water, the Tf2 N is subject to reduction, producing a variety of decomposition products. The IL designers and synthesizers
should cooperate more intensively with fundamental electrochemists and theoreticians to develop new ILs which do not exhibit such undesired electrochemical side
reactions.
A further important aspect is how to handle reactive elements? It was found in
the Clausthal group that nanocrystalline aluminum and nanoscale silicon made in
1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide react even with
the comparably low level of oxygen (<1 ppm) in an inert gas glove box. Under air
the deposit can be oxidized on the time scale of a few days. Maybe in situ passivation
methods will have to be developed. One could think about deposition of a reactive
element in an ionic liquid, washing off the ionic liquid, followed by passivation in
a different liquid.
13.4
Nanomaterials/Nanoparticles
Usually there is a lot of effort required to make nanomaterials by electrochemical means. In aqueous solutions the electrodeposition of nanocrystalline metals
requires pulsed electrodeposition and the addition of additives whose reaction
mechanism hitherto has only been partly understood (see Chapter 8). A further
shortcoming is that usually a compact bulk material is obtained instead of isolated
particles. The chemical synthesis of metal or metal oxide nanoparticles in aqueous
or organic solutions by colloidal chemistry, for example, also requires additives and
often the desired product is only obtained under quite limited chemical conditions.
Changing one parameter can lead to a different product.
In ionic liquids the situation seems to be totally different. It was surprising to us that the electrodeposition of metals and semiconductors in 1-butyl1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide delivers nanocrystalline
deposits with grain sizes varying from 10 to 200 nm for the different materials,
like Si, Al, Cu, Ag and In, investigated to date. It was quite surprising in the
case of Al deposition that temperature did not play a tremendous role. Between
25 and 125 ◦ C we always got nanocrystalline Al with similar grain sizes. Similar
results were obtained if the deposition was performed in tri-hexyl- tetradecylphosphonium bis(trifluoromethylsulfonyl)amide. Maybe liquids with saturated nonaromatic cations deliver preferentially nanomaterials; this is an aspect which, in
our opinion, deserves further fundamental studies.
It was quite surprising (but also quite pleasant) to find that by plasma electrochemistry (see Chapter 10) isolated nanoparticles could be made in an ionic
liquid. The physical mechanisms are not yet fully understood but it is likely that
the particle size can be influenced by the ionic liquid itself, the metal salt concentration and the temperature. As the particles coagulate on a long time scale, at
13.6 Polymers for Batteries and Solar Cells 373
the present state-of-the-art one might think about tailor-made ionic liquids which
form a surface layer on the particles, thus protecting them against agglomeration.
Catalysis studies would be of interest to investigate the catalytic efficiency of such
nanoparticles.
13.5
Cation/Anion Effects
As pointed out above there are unexpected cation effects on the electrodeposition of
metals in ionic liquids leading, in one liquid, to nanocrystalline metals, in another
liquid to microcrystalline metals. Viscosity effects alone are excluded. Furthermore,
it is known that the addition of toluene or benzene to first generation ionic liquids
based on AlCl3 can lead to strongly improved deposits, in part with shiny appearance. In our opinion it is worth investigating to what extent side chains in a cation
or an anion can influence the quality, structure and grain size of electrochemically
made deposits. For example, what happens if in 1,3-dialkyl-imidazolium cations
one side chain is modified, for example a short ethyl group is replaced by a comparatively long tetradecyl group with a long hydrophobic chain? One could also
consider the introduction of aromatic groups in the cation (e.g. a benzyl group)
given the effect of aromatic additives noted above; potentially the effect of toluene
and benzene on Al deposition could thus be realized by non-volatile, and thus environmentally friendly, side groups. It might also be of interest to investigate the
electrodeposition of metals from mixed ionic liquids.
13.6
Polymers for Batteries and Solar Cells
Electrodeposition of conducting polymer materials from ionic liquids (see Chapter
7) clearly has important potential to generate a new and wider range of conducting polymer materials and morphologies. The morphology aspect is particularly
important in applications such as batteries and photoelectrochemical solar cells
where the internal, electrochemically accessible surface area of the material is a
critical parameter. Thus there is scope for development of a range of novel conducting polymer films for these devices. On the nanometer length scale there is
also the scope to produce conducting polymer nanoparticles and nanofibres via
electropolymerization in the ionic liquid or at the interface between an ionic liquid and another phase. Similarly the first steps are emerging that will allow the
preparation of metal nanoparticle composites in conducting polymer materials.
Without doubt, as has been shown in the case of the thiophene oligomers and benzene, the greater potential window of the ionic liquids will allow the electropolymerization of monomers/oligomers which cannot be polymerized by chemical
means.
374 13 Future Directions and Challenges
13.7
Variable Temperature Studies
Hitherto almost all electrochemical studies in ionic liquids have been performed at
moderate temperature, often at room temperature. This motivation may be based
on aqueous electroplating processes that are mostly performed between 30 and
70 ◦ C but one should not neglect the wide thermal window of ionic liquids: 1butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide as an example can
be heated to 200 ◦ C without considerable decomposition. In comparison to an
aqueous electroplating process this temperature is maybe high, however, in comparison to a “real” molten salt process at temperatures of 500 ◦ C and more, this
is quite a low temperature. Technical tantalum electrodeposition is carried out
around 600–700 ◦ C in molten salts. In comparison to this, a low temperature process at 200–250 ◦ C with comparable quality would be a milestone. The aqueous
electrodeposition of selenium is limited by the fact that, even at 100 ◦ C considerable amounts of insulating black and red selenium are formed, although the phase
transition temperature from black and red selenium to gray metallic selenium is at
80 ◦ C. If electrodeposition is not focussed on low temperatures, further benefits of
ionic liquids arise: the electrodeposition of alloys and compound semiconductors
that often require considerable activation energy might be facilitated at elevated
temperature. Thus it is likely that ionic liquids are the missing link in terms of
temperature regimes between aqueous/organic electrochemistry and molten salt
electrochemistry. In our opinion, this somewhat neglected region of temperature
should be considered more seriously.
13.8
Intrinsic Process Safety
The toxicological properties of some ionic liquids are only now being quantified.
Ionic liquids with alternative cations such as those derived from biodegradable imidazoles, lactams, amino acids and choline have been prepared although it is only
the last of these which has been used for metal deposition. Liquids that are less
toxic tend to have narrow potential windows. It should, however, be appreciated
that the most toxic component of any ionic liquid is still likely to be the metal salt,
which is naturally the same for aqueous electroplating. Some efforts have been
made to substitute high oxidation state metal salts with other less toxic alternatives
e.g. replacing CrO3 with CrCl3 . Ionic liquids do not represent a safety hazard in
terms of their flammability as most will not burn, even upon contact with a naked
flame. The extremely low vapor pressure makes them easier to handle and generally
circumvents the necessity for air extraction. Many liquids are, however, sensitive to
moisture content and therefore may have to be handled under a controlled atmosphere. There is still a large amount of fundamental optimization that needs to be
carried out before the overall green credentials of ionic liquids can be ascertained.
The liquids will have to be recycled to make them economically viable and so it
13.10 Which Liquid to Start With? 375
should only be trace levels that will be emitted to the environment through material rinsing. Methodologies will have to be developed to minimize these, possibly
through an initial rinse with non-aqueous solvents allowing the bulk of the drag-out
to be separated and recycled.
13.9
Economics (Price, Recycling)
There is naturally a significant difference between the current retail price for most
ionic liquids and the current cost of aqueous electroplating solutions. It is difficult
to imagine that many ionic liquids will ever approach the desirable $20 / kg level.
This is due fundamentally to the synthetic complexity involved in producing ionic
liquids. It is, however, essential to comprehend that the cost of the liquids need not
necessarily be an issue, as the key driver will be the running costs of the process, of
which the capital outlay for the liquid may only be a small component. The overall
cost will be made up of:
1.
2.
3.
4.
5.
6.
7.
pre-treatment costs
equipment costs
the cost of the liquid
power consumption
post-treatment costs
disposal/recycling costs
labour costs.
Points 1, 2, 5 and 7 will be effectively the same in both aqueous and ionic liquids.
Point 3 will naturally be higher in an ionic liquid than an aqueous solution although
some liquids, particularly the eutectic-based systems, are approaching the costs
of current aqueous solutions. The power consumption for ionic liquid processes
should be less than for water-based systems due to higher current efficiencies and
hence the overall economic viability of ionic liquids will depend upon the balance
of the disposal and/or recycling costs. The key issue to address is the longevity of
the ionic liquids which, in principle, should be long if soluble anodes are used, but
depends upon the breakdown products of the ionic liquid. No sustained studies
of this issue have been undertaken and it is particularly important that these are
carried out under high current density conditions.
13.10
Which Liquid to Start With?
We have often been asked this question by beginners. To be honest, there is no
common answer to the question. On the one hand there is an incredibly high number of theoretically possible liquids (of which around 300–500 are commercially
376 13 Future Directions and Challenges
available in 2008), on the other hand one liquid does not make it all. In Chapter 12
we have given some plating protocols that can be easily performed by beginners in
the field. Apart from these recipes, we generally recommend to employ chemically
stable ionic liquids. Liquids with PF6 − or BF4 − anions cannot be recommended.
Maybe they are cheap, but under the wrong conditons they are subject to strong
degradation liberating highly toxic HF gas. In the best case just the electrochemistry
gets fairly reproducible. From our point of view liquids with the CF3 SO3 − anion
are good candidates for the beginner, as this anion is both chemically and electrochemically pretty stable. A suitable cation would be 1-ethyl-3-methylimidazolium.
We would not like to recommend liquids with pyrrolidinium ions for the beginner
as they tend to lead to nanostructured deposits. The beginner might misinterpret
the results and conclude that ionic liquids only deliver bad deposits. This liquid
and others are available in different qualities, as mentioned briefly in Chapter 2.2.
Despite the possibly somewhat higher price we would like to recommend purchasing these liquids in ultrapure quality, which means that the impurity level is
below 10 ppm. This would in any case exclude tremendous impurity effects, allowing reproducible electrochemical experiments. This liquid is not sensitive to
water, although it absorbs water under environmental conditions, and water can be
removed by evacuating the liquid under stirring and at an elevated temperature of
about 100 ◦ C. With such a high quality liquid, especially if the experiments can be
performed under controlled inert gas atmosphere, the beginner would get a feeling
of how to handle these liquids, thus helping him to perform further experiments.
13.11
Fundamental Knowledge Gaps
Apart from the above mentioned points, ionic liquids are fascinating liquids for
fundamental physicochemical studies. Unfortunately, there are less than a handful
of groups investigating the local processes at the interface electrode/electrolyte in
ionic liquids with in situ scanning probe microscopy. One shortcoming is maybe
that this requires ultrapure liquids with an extremely low level of impurities, thus
inert gas conditions are strongly required. Nevertheless, there are unprecedented
effects being observed that have not yet been described in aqueous solutions and
there are hints that the double layer is relatively thick, interfering with the electrode
surface, thus making atomic resolution, at the very least, difficult.
The nature of the double layer in ionic liquids is a fundamental issue that
is important to many applications but is little understood. The fact that double
layer charging can produce only a limited range of concentration changes near
the electrode means that the double layer is probably much thicker in an ionic
liquid than it is in a solution-based electrolyte. This requires both theoretical and
experimental investigation.
Speciation of metal ions and complex ions in ionic liquid solution prior to electrodeposition is also clearly an important issue in understanding and developing
electrowinning processes. Little is known about the stoichiometry and structure of
13.11 Fundamental Knowledge Gaps 377
complex metal ion species, even in the most studied aluminium-based systems.
When one considers the permutations of all the metal species of interest, the possible metal salt precursors and the possible ionic liquid solvents, there is plainly an
enormous body of work needing to be done in this field.
These issues impact on the most basic of knowledge gaps that one is confronted
with in this field: transport and thermodynamic property (e.g. redox potential) data.
Again, there is an enormous body of work needed in this area and we hope that
this book will serve to stimulate a new generation of researchers to undertake this
important task.
379
Subject Index
a
acetylacetonatotetramethylethyldiaminecopper
(II) 62
additive 118
– organic 216ff.
adsorbent 323
adsorption process 327
1-alkyl-3-methylimidazolium
tetrafluoroborate 50
1-alkylimidazole 17f.
alloy
– deposition 7
– electrochemical alloying 137
– electrodeposition 125ff.
– nanocrystalline 8, 222
– nanostructured 213ff.
aluminium
– alloy 88, 128ff., 222, 316
– anode 357
– antimonide 150
– 1-butyl-3-methylimidazolium chloride
356
– 1-butyl-3-methylpyrrolidinium
bis(trifluoromethylsulfonyl)amide 360
– chromium 132
– electrodeposition 88ff., 128, 245ff., 329,
353ff.
– 1-ethyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)amide 358f.
– 1-ethyl-3-methylimidazolium chloride
127f., 353f.
– film 229
– halide 16
– Hall-Héroult process 289
– magnesium 131
– manganese 131, 222
– microcrystalline 118
– molybdenum 128ff.
– nanocrystalline 372
– nanostructured 222f.
– nickel 132
– titanium 127
– zirconium 129
aluminium chloride 337
aluminium chloride based ionic liquid
– electrodeposition 84
amine 71
ammonium cation
– quaternary 42, 110
ammonium halide
– quaternary 33
anion
– cation effect 119, 373
anode material 11
anodic dissolution 91
anomalous codeposition (ACD) 219
antimony
– electrodeposition 91ff.
attenuated total reflection (ATR) 198
autosolvolysis 20
b
BASIL process 4
battery 373
[BBIM]+ 22
benzoic acid 223f.
benzyltrimethyl ammonium [BTMA]+
chloride 130
1-benzyl-3-methylimidazolium [BZMIM]+
60
betaine dye 59
bipolaron 200
bis(trifluoromethylsulfonyl)amide [NTF]− ,
[NTf2 ]− , [Tf2 N]− , [TFSI]− , [TFSA]− 27,
48, 54ff., 93, 98ff., 140, 180, 243
bithiophene 185
brightener 11, 221, 315
bromoalkane 18
Brønsted acidity 104, 336
Electrodeposition from Ionic Liquids. Edited by F. Endres, D. MacFarlane, A. Abbott
C 2008 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim
Copyright
ISBN: 978-3-527-31565-9
380
Subject Index
1-butyl-2,3-dimethylimidazolium
[BMMIM]+ , [b-dimIM]+ 56ff.
– chloride 36
1-butyl-3-methylimidazolium [BMIM]+ ,
[C4 mim]+ 20ff., 50ff., 181ff.
– bis(trifluoromethylsulfonyl)amide 102,
189
– chloride 150, 222, 356
– chloride/aluminium chloride 356
– decomposition pathway 272f.
– hexafluorophosphate 96, 151f., 181,
194ff.
– tetrafluoroborate 96, 200, 273
– trifluoromethylsulfonate 96
1-butyl-1-methylpyrrolidinium [BMP]+ ,
[C4 mpyr]+ , [Py1,4 ]+ 98ff., 114, 155, 229,
241, 252
– bis(trifluoromethylsulfonyl)amide [Tf2 N]−
94ff., 114, 155ff., 200, 227, 241, 252, 271,
329, 339, 360f.
– decomposition pathway 271f.
– trifluoromethylsulfonate [TFO]− 99
butylpyridinium [BP]+ , [bpyr]+ 56,
134
– zinc chloride 134
c
[C2 mim] 168ff.
– hexafluorophosphate 168
– tetrafluoroborate 168
– trifluoromethanesulfonate 176
cadmium
– alloy 133
– electrodeposition 95, 133
– removal 323
– semiconductor 147
– telluride 151
– zinc 133
carbon
– activated 322
cathode 357
cation
– anion effect 119, 373
– structure 10
charge mobility characteristic 5
charge trapping 192
chloroaluminate 10
– ionic liquid 15, 222
– synthesis 19
chloroaluminate liquid
– colorless 17
chlorometallate 39
– ionic liquid 222ff.
chlorozincate 105ff., 132ff.
– anion 35
choline 38
– chloride 34, 108, 232, 365
chromium 38
– alloy 132
– aluminium 132
– electrodeposition 95, 132
– film 109
– metallic-looking coating 108
– nanocrystalline film 95
– removal 323
chronoamperometry 106
cobalt
– alloy 134
– electrodeposition 134
– zinc 134
component ion 73
concentration 323
conducting polymer 167ff.
– air- and water-stable ionic liquid 179ff.
– chloroaluminate ionic liquid 177ff.
– electrochemical characterization 191
– morphological characterization 194
– spectroscopic characterization 198
– synthesis 177ff.
conductivity 117, 313f.
– ionic 70
– ionic liquid 313
– modelling 40
– molar 5
copper 112f.
– (II)
acetylacetonatotetramethylethyldiamine
62
– alloy 133
– electrodeposition 94, 133ff., 229
– foil 89
– indiumselenide (CIS) 161
– nanocrystalline 94, 229
– plasma electrochemical deposition
(PECD) 278f.
– removal 323
– rotating disk electrode (RDE) 127
– tin 142
– trifluoromethanesulfonate [TFO] 231
– zinc 133
corrosion
– resistance 2, 143
counter electrode 370
– reaction 287ff., 317
crystallite size 216ff.
CVD 311
Subject Index
cyano-based ionic liquid 28
2-(cyclohexylamino)ethanesulfate [CHES]−
60
d
de-alloying 137
– electrochemical 137
1-decyl-3-methylimidazolium cation
[DMIM]+ , [decyl-MIM]+ 57ff.
deep eutectic solvent (DES) 39, 336, 365
density 56f.
deposition
– mechanism 216
– metal from non-chloroaluminate eutectic
mixture 103
– nanometal 217
– reactive element 114
1,3-dialkylimidazolium tetrachloroaluminate
20
N,N-dialkylpiperidinium 27
N,N-dialkylpyrrolidinium 27
– based ionic liquid 114
dicyanamide anion [DCA]− 29
1,2-diethyl-3,4-dimethylimidazolium
[DEDMIM]+ 72
1,3-diethyl-5-methylimidazolium 186, 193
N,N-diethyl-3-methylpyrazolium [DEMPZ]+
72
N,N-diethyl-4-nitroaniline 59ff.
diffusion coefficient 73f.
diluent 10, 111
R process 4
Dimersol
1,2-dimethyl-3-(n-propyl)imidazolium
[DMPIM]+ 72
dimethylimidazolium [MMIM]+ 22
dipolarity 59
donor number (DN) 62
dopant
– chiral 204
– Raman-active 198
doping 187ff.
double layer structure 10
dye-sensitized solar cell (DSSC) 29
e
ED(A)X, see energy
electrochemical alloying 137
electrochemical application 47ff.
electrochemical atomic layer epitaxy
(ECALE) 148
electrochemical cycling 193
electrochemical method 172
electrochemical property 66
electrochemical quartz crystal microbalance
(EQCM) 102, 193
electrochemical synthesis
– conducting polymer 175ff.
electrochemical window 270
electrochemistry
– impurity 338
electrode 173f., 261
– blocking 69
– first kind 298
– redox 298
– reference, see reference electrode
– second kind 298
electrode-free discharge 265
electrodeposition
– air- and water-stable ionic liquid 92
– alloy 125ff.
– aluminium 245ff., 353
– aluminiumchloride-based ionic liquid
84, 353
– antimony 91
– chloroaluminate ionic liquid 126ff.
– chromium 95
– eutectic-based ionic liquid 104
– gallium 91
– indium 90
– ionic liquid 1ff., 84ff., 147ff.
– lithium 84
– metal 83ff.
– nanocrystalline 8
– nanometer scale 239ff.
– palladium 96
– platinum 96
– pulsed, see PED
– semiconductor 147ff.
– silver 96
– sodium 86
– tantalum 250
– tellurium 92
– tin 91
– zinc 92
electrodialysis 322
electrolyte 261ff.
electrolytic solution 297
electromotive force 260f.
electropickling 7
electroplating 6, 319ff., 345
electropolishing 7, 293ff.
– stainless steel 293
electropolymerization 171ff.
– solution-surface 181
electrostenolysis 266
Endres-quality 342
381
382
Subject Index
energy
– dispersive X-ray (EDX, EDAX) 109ff.,
128
1-ethyl-2,3-dimethylimidazolium
[e-diMIM]+ 67, 193
1-ethyl-3H-imidazolium
– trifluoroacetate 204
1-ethyl-3-methylimidazolium [EMIM]+ ,
[C2 mim] 29ff., 48ff., 69, 118, 168ff.
– bis(trifluoromethylsulfonyl)amide 98,
118, 176ff., 247, 340ff.
– chloride 18f., 85ff., 107, 134ff., 353
– chloride/aluminium trichloride ionic
liquid 85, 353
– ethylsulfate 31
– tetrafluoroborate 96
– trifluoromethanesulfonate 176
– zinc(II)chloride 95, 105
ethylene glycol 111, 365
eutectic
– choline chloride based 111
– eutectic-based ionic liquid 31ff., 104
– impurity in deep eutectic solvent 336
– mixture 103
– point 32
– solvent, see also deep eutectic solvent
336, 365
– type I 33ff., 103ff.
– type II 38, 103ff.
– type III 38f., 103ff.
extraction process 327
f
FAB MS, see mass spectrometry
ferric chloride 36, 107
ferrocene (Fc) 66, 305, 339
ferrocenium (Fc+ ) 66, 339
fluidity 5
Fourier transform infrared spectroscopy
(FTIR) 198
fragility parameter 69
g
gallium
– arsenide 107f., 149
– electrodeposition 91
germanium 151ff.
– electrodeposition 151, 232
– Ge(111) 231
Gibbs enthalpy 216
glass transition temperature 51
gold
– Au(111) 241ff., 338ff.
– electrodeposition 140
– nano 217
– porous 139
grain refiner 216ff.
grain size 217
greenness 319
h
haloaluminate
– anion 16
– ionic liquid 21ff.
HBD, see hydrogen bond donor
Helmholtz layer 10
heteronuclear Overhauser effect
spectroscopy (HOESY) 39
hexafluorophosphate [PF6 ]−
– ionic liquid 24
1-n-hexyl-3-methylimidazolium [HMIM]+
26ff., 253
– tris(pentafluoroethyl)trifluorophosphate
[FAP] 253ff.
1-hexyl-1-methylpyrrolidinium [HMPL]+
30
N-hexylpyridinium [HPYR]+ 30
high resolution scanning electron
microscopy (HRSEM) 277ff.
high resolution transmission electron
microscopy (HRTEM) 277ff.
highly oriented pyrolytic graphite (HOPG)
157
hydrogen bond
– acidity 59
– basicity 59
– donor (HBD) 39, 110
hydrophobicity 28ff.
2-hydroxylethyl ammonium formiate
205
1-hydroxyethyl-3-methylimidazolium
[OH-EMIM]+ 60
i
imidazolium 33, 51ff.
– ionic liquid 51
– salt 49
impurity 117, 334, 369f.
– gaseous 336
– particulate 336
– synthetic 334
indium 90, 140
– antimonide 108, 149
– electrodeposition 90, 140
– palladium 140
– tin-oxide (ITO) 188ff.
inert gas condensation (IGC) 214
Subject Index
ion conduction 75
– selective 76
ion conductor
– anisotropic 53
ion radius 48
ionic conductivity 68ff.
ionic fluid 3
ionic liquid (IL)
– air- and water stable 21ff., 140ff., 227
– biodegradable 374
– chiral 203f.
– chlorometallate based 222ff.
– chlorozincate 105ff., 222
– definition 4
– distillable 204
– electrodeposition 92ff.
– electrolyte 220f
– eutectic-based 104
– ferromagnetic 36
– haloaluminate-based 17ff.
– hexafluorophosphate 24
– moisture-stable 23
– nanocrystalline metal 227
– nanomaterial 371
– physical property 47ff.
– physicochemical property 29
– plasma electrochemical metal deposition
274
– plasma electrochemistry 259ff.
– protic 204
– recycling 375
– regeneration 319ff.
– reuse 319
– specific ion 75
– stability 269
– synthesis 15ff.
– task-specific, see TSIL
– technical potential 6
– tetrafluoroborate 24
– type III 344
– water content 174
– water-stable 21ff., 140ff., 227
iron
– (II) chloride 36, 107
– (III) chloride 36, 107
– alloy 135, 218f.
– electrodeposition 135
– nanostructured 226
– zinc 135
ITO (indium-tin-oxide), see indium
k
Kamlet-Taft parameter
59ff.
l
Lewis acid 16ff.
– aluminium chloride 89
– ionic liquid 35
Lewis acidity 104ff., 222, 336
Lewis base 10
life cycle analysis (LCA) 9
liquid
– crystallinity 53
– glycol-based 40
– junction 300
– junction potential (LJP) 300
liquor
– ionic liquid 324
– water-based 321ff.
lithium
– bis(trifluoromethylsulfonyl)amide [NTF]−
27, 188, 341, 362
– 1-butyl-1-methylpyrrolidinium [BMMIM]
bis(trifluoromethylsulfonyl)imide/
lithium bis(trifluoromethylsulfonyl)
imide 362
– electrodeposition 24, 85, 100
– [FAP]− 28
– salt 72ff.
m
macroporous structure 139
magnesium
– alloy 131ff.
– aluminium 131
– deposition 100, 114
– electrodeposition 136
– zinc 136
manganese
– alloy 131
– aluminium 131
– electrodeposition 143
– zinc 143
mass spectrometry (MS)
– fast atom bombardment (FAB MS) 34,
105
mass transport 104
material compatibility 312
mechanical process 328
melting point 47ff.
mesoporous structure 139
metal
– alloy 126
– air- and water stable ionic liquid 227
– deposit 218
– dissolution process 287
– group I 84
383
384
Subject Index
metal (continued)
– group II 88
– group III 88
– group IV 91
– group V 91f.
– group VI 92
– halide cluster 115
– less reactive 93
– nanocrystalline 8, 222ff.
– nanostructured 213ff.
– non-chloroaluminate eutectic mixture
103
– oxide 37
– porous surface 137
– reactive metal 97
– salt 64
– water-stable ionic liquid 227
1-methyl-3-methylimidazolium [MMIM]+
22
N-methylpyrrolidinium formiate 54
microporous structure 139
[MMIM], see dimethylimidazolium
mobility 41
molybdenum
– aluminium-alloy 128ff.
– electrodeposition 128ff.
– 1-ethyl-3-methylimidazolium chloride
128
montmorillonite 323
n
nanoalloy 218
– deposition 218
nanocomposite coating 8
nanomaterial 213, 371f.
– ionic liquid 371
– tailor-made 214
nanoparticle 8, 372
nanoporous structure 139
nanostructured metal
– electrodeposition 215
nanotube 205
nanowire 205
Nernst equation 297
Nernst-Einstein equation 41, 73
nickel
– alloy 132, 218
– aluminium 132
– chromium 132
– electroless plating 322
– nano 214
– recovery 323
– titanium alloy 102
nicotinic acid 316
nile red 64
niobium
– alloy 139
– electrodeposition 139
– ionic liquid 37
– tin 139
4-nitroaniline 59ff.
non-chloroaluminate eutectic mixture 103
nuclear magnetic resonance (NMR)
– pulse-field-gradient (PFG NMR) 73ff.
– solid-state 201
o
1-octyl-3-methylimidazolium [OMIM]+
55ff.
onium cation 48
p
palladium
– deposition 316
– electrodeposition 96, 140
– gold 140
– indium 140
– plasma electrochemical deposition
(PECD) 280
– silver 140
PED (pulsed electrodeposition) 214f.
– aqueous electrolyte 215
PEDOT, see
poly(3,4-ethylenedioxythiophene)
PFG NMR, see nuclear magnetic resonance
phosphorus oxyfluoride 335
plasma
– electrochemical cell 266
– electrochemical deposition (PECD)
274ff.
– electrochemical metal deposition in ionic
liquid 274
– electrolysis 267f.
– electrolyte interface 262
– experiment 269
– low-temperature 261
– reactor 264
– type 264
plasma electrochemistry 259ff.
plating protocol 353ff.
platinum
– (II) tetraamino 130
– (IV) bis(acetylacetonato) 130
– alloy 130
– aluminium 130
– electrodeposition 96, 130ff.
– plasma electrochemical deposition
(PECD) 280
– zinc 136ff.
Subject Index
PMIM, see 1-n-propyl-3-methylimidazolium
polarity 58ff.
polarizability 59
polarization 262
polaron 200
poly(aniline) 168, 179, 193ff.
poly(bithiophene) 183
poly(3,4-ethylenedioxythiophene) (PEDOT)
168, 188ff., 200
poly(fluorene) 178
poly(3-(4-fluorophenyl)thiophene) 186
poly(N-methylpyrrole) 194
poly(para-phenylene) (PPP) 178, 191,
252ff.
– electrodeposition 252ff.
poly(pyrrole) 168ff., 179, 200ff.
poly(terthiophene) 176ff., 205
poly(thiophene) 168ff., 183ff.
– nano 205
– paradox 183
polyethylene glycol (PEG) 322
polymer
– coating 7
– conducting, see conducting polymer
– heterocyclic 171
– nano-dimensional 205
polymerization
– chemical 205
– electrochemical, see electropolymerization
post-treatment protocol 317
potential difference 297
potential energy
– interionic 35
potential window 66
pretreatment protocol 289f., 313
probe molecule
– solvatochromic 61
process liquor
– regeneration 321
– water based 321
process safety
– intrinsic 374
1-n-propyl-3-methylimidazolium [PMIM]+
22
propylene carbonate/tetrabutylammonium
– hexafluorophosphate 168
PVD 311
pyridine 17
pyridinium 33
q
quartz crystal
– metal-coated 193
quaternization reaction 17f.
r
recovery 320
– electrolyte 320
recycling 318, 375
redox process 328
reference electrode 296ff.
– characteristics 298f.
– pseudo/quasi 299ff.
– room temperature ionic liquid 296
refractive index 56
regeneration
– electrolyte 320
– ionic liquid 319ff.
Reichardt’s betaine dye 59ff.
reuse
– electrolyte 320
– ionic liquid 319
room temperature ionic liquid (RTIL) 24,
300ff.
– structure 303
rotating disk electrode (RDE) 127
– copper 127
s
scanning tunnelling microscope (STM)
239ff.
– in situ 239ff.
selenium 160f.
– CIS 161
– gray 160f.
semiconductor 147ff.
side chain effect 49
SIGAL process 3, 97
silicon 155ff.
silver 308f.
– de-alloying 137
– deposition 309
– electrode 138
– electrodeposition 96, 140
– palladium 140
– plasma electrochemical deposition
(PECD) 274ff.
– reference electrode 308
– surface diffusion 138
solar cell 373
– copperindiumselenide (CIS) 161
– dye-sensitized (DSSC) 29
solubility 64
solution
– aqueous 2
– non-aqueous 3
solvatochromism 58
solvent
– polarity 59
385
386
Subject Index
stability
– electrochemical 270
– ionic liquid 269
Stokes-Einstein equation 41
surface energy 213
t
tantalum 114
– 1-butyl-1-methylpyrrolidinium
[BMP]bis(trifluoromethylsulfonyl)imide
364
– deposition 116
– electrodeposition 100, 250ff.
– ionic liquid 37
– XRD amorphous 101
tellurium
– electrodeposition 92
temperature 9, 172
– variable 374
tetraalkylammonium 27
tetraalkylphosphonium 27
tetracyanoborate anion [TCB]− 29
tetraethylammonium cation [N2222 ]+ 48
tetrafluoroborate [BF4 ]− 69
– ionic liquid 24
tetramethylammonium cation [N1111 ]+ 48
N,N,N′ ,N′ -tetramethylphenylenediamine
radical cation/N,N,N′ ,N′ -tetramethylphenylenediamine couple (TMPD·+ /
TMPD) 305f.
[TFSI]− , see
bis(trifluoromethylsulfonyl)amide
thermal conductivity 53
thermal decomposition temperature 52
thermal degradation temperature 47
thermal unit operation 325
R
Therminol
53
thiocyanate 29
thiophene 183
tin
– alloy 133ff.
– chloride 107
– copper 142
– deposition 111
– electrodeposition 91, 133ff.
– indium 142
– niobium 139
– zinc 133
titanium 102
– chloride 102
– electrodeposition 102, 114
– 1-ethyl-3-methylimidazolium chloride
127
toxicity 8
transition
– solid-solid 53
tri-1-butylmethylammonium [TBMA]+
143
– bis((trifluoromethyl)sulfonyl)amide 143
tricyanomethide anion [TCM]− 29
triflate [OTf]− , [TfO]− , [Tf]− 25, 93, 180
– ionic liquid 25
trifluoroacetate [ATF]−
– ionic liquid 25
trifluoromethylsulfonate, see triflate
trimethyl-n-hexylammonium [TMHA]+
140
– bis (trifluoromethylsulfonyl)amide 94,
140
trihexyl-tetradecyl phosphonium [P14,6,6,6 ]+ ,
[P6,6,6,14 ]+ 98, 229
– bis(trifluoromethylsulfonyl)amide [Tf2 N]
98, 202f.
2,4,6-triphenylpyridinium-N-4-(2,6diphenylphenoxide)betaine
59
tris(trifluoromethylsulfonyl)methide 93
trispentafluoroethyltrifluorophosphate
[FAP]− 28
TSIL (task specific ionic liquid) 65
u
underpotential deposition (UPD) 105, 128,
151, 224f.
urea 111ff.
v
vanadium
– electrodeposition 114
vapour pressure 54
viscosity 40ff., 104ff.
– modelling 40
Vogel-Tamman-Fulcher (VFT) relationship
40, 69
voltammetry 305f.
– quantitative 308
w
Walden product 69
Walden rule 5f., 41f.
Warren-Averbach technique 217
waste treatment 317
water 335
– content 174
Wilkes 23
Wiliamson Hall procedure 217
Subject Index 387
x
X-ray powder diffraction (XRD) 127
– pattern 101
X-ray photoelectron spectra (XPS) analysis
128, 199
z
zinc
– (II) chloride 36, 94
– (II) chloride-1-ethyl-3-methylimidazolium
chloride 95, 105
– cadmium 133
– chlorozincate 105ff., 132ff.
– coating 365
– cobalt 134
– copper 133f.
– deposition 111, 344
– electrodeposition 93, 132ff., 365
– iron 135
– magnesium 136
– manganese 143
– nickel 136
– platinum 136
– tin 133
zinc telluride
– electrodeposition 150
zirconium
– aluminium alloy 129f.
– electrodeposition 129f.
zwitterionic salt 77