Fluorinated Materials for Energy Conversion

Fluorinated Materials for Energy Conversion

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Fluorinated Materials for Energy Conversion Edited by

Tsuyoshi Nakajima Department of Applied Chemistry Aichi Institute of Technology Toyota-shi, Japan and

Henri Groult Pierre and Marie Curie University CNRS Paris, France

2005

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Preface Since the discovery of fluorine by Henri Moissan at the end of the 19th century, fluorine chemistry has developed many applications in energy conversion, medicine, biology, agriculture, telecommunication and so on. However, fluorine chemistry is not widely spread probably because of the difficulty in the fluorination techniques and handling of fluorinating reagents and fluorides. Elemental fluorine is a typical fluorinating gas with high reactivity arising from its small dissociation energy. Efficient production of elemental fluorine by molten salt electrolysis is still one of the important research subjects in fluorine chemistry notably because of industrial applications in nuclear energy field. A new and promising aspect in fluorine chemistry is the applications of fluorination reactions and various fluorides to energy conversion materials for lithium batteries, fuel cells, solar cells etc. Many examples regarding the introduction of fluorine into lithium battery materials, that is, fluorination of carbonaceous anodes and oxide cathodes, synthesis of new fluorine containing electrolytes, fluorination of organic solvents and so on, were recently reported and revealed the importance of fluorine chemistry in this field. It was also shown that the fluorinated materials had important roles in fuel cells and solar cells. This summarizes the recent advances on these topics. All authors are specialists actively working in fluorine chemistry, electrochemistry, polymer chemistry and solid state chemistry. We hope that the book offers new aspects of fluorine chemistry to readers in the various fields. Tsuyoshi Nakajima and Henri Groult

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Contributors Takeshi Abe Department of Energy and Hydrocarbon Chemistry, Kyoto University Japan

Guy Campet Institut de Chimie de la Matière Condensée de Bordeaux (ICMCBCNRS) France

Fannie Alloin Laboratoire d'Electrochimie et de Physico-Chimie des Matériaux et des Interfaces, INPG-CNRS UMR 5631 France

Vittorio Canevari IMEM-CNR Institute Italy

Bruno Améduri Laboratoire de Chimie Macromoléculaire, ENSCM-CNRS UMR 5076 France François Blanchard Faculté des Sciences, Université F. Rabelais France Alessio Bosio Department of Physics, University of Parma Italy Gérard Bosser Faculté des Sciences, Université F. Rabelais France Hubert Cachet Laboratoire des Interfaces et Systèmes Electrochimiques (UPR15-CNRS)–Université Pierre et Marie Curie France Magali Caillon-Caravanier Faculté des Sciences, Université F. Rabelais France

Bernard Carré Faculté des Sciences, Université F. Rabelais France Alexandre Chagnes Faculté des Sciences, Université F. Rabelais France Stephen E. Creager Depeartment of Chemistry, Clemson University U.S.A. Darryl D. DesMarteau Depeartment of Chemistry, Clemson University U.S.A. Marc Dubois Laboratoire des Matériaux Inorganiques, Université Blaise Pascal de Clermont-Ferrand France Olt E. Geiculescu Depeartment of Chemistry, Clemson University U.S.A. Henri Groult Laboratoire LI2C- CNRS UMR 7612, Université Pierre et Marie Curie France

viii

Contributors

Katia Guérin Laboratoire des Matériaux Inorganiques, Université Blaise Pascal de Clermont-Ferrand France

Kazuhiko Matsumoto Department of Fundamental Energy Science, Kyoto University Japan

Rika Hagiwara Depeartment of Fundamental Energy Science, Kyoto University Japan

Yoshiaki Matsuo Department of Materials Science and Chemistry, University of Hyogo Japan

André Hamwi Laboratoire des Matériaux Inorganiques, Université Blaise Pascal de Clermont-Ferrand France

Bengt-Erik Mellander Physics Engineering Physics, Chalmers University of Technology Sweden

Bernard Jousseaume Laboratoire de Chimie Organique et Organométallique, Université Bordeaux I France Kiyoshi Kanamura Department Applied Chemistry, Tokyo Metropolitan University Japan HanSu Kim Samsung Advanced Institute of Technology (SAIT) South Korea Chai-Won Kwon Samsung Corning Precision Glass South Korea Frédéric Lantelme Laboratoire LI2C- CNRS UMR 7612, Université Pierre et Marie Curie France Daniel Lemordant Faculté des Sciences, Université F. Rabelais France

Bénédicte Montigny Faculté des Sciences, Université F. Rabelais France Régine Naejus Faculté des Sciences, Université F. Rabelais France Tsuyoshi Nakajima Department of Applied Chemistry, Aichi Institute of Technology Japan Benjamin G. Nolan Department of Chemistry, Colorado State University U.S.A. Madeleine Odgaard IRD Fuel Cells A/S Denmark Zempachi Ogumi Department of Energy and Hydrocarbon Chemistry, Kyoto University Japan

Contributors

Yoshimi Ohzawa Department of Applied Chemistry, Aichi Institute of Technology Japan

Hidekazu Touhara Department of Chemistry, Shinshu University Japan

Nicola Romeo Department of Physics, University of Parma Italy

Thierry Toupance Laboratoire de Chimie Organique et Organométallique, Université Bordeaux I South Korea

Jean-Yves Sanchez Laboratoire d'Electrochimie et de Physico-Chimie des Matériaux et des Interfaces, INPG-CNRS UMR 5631 France Yukio Sasaki Department of Nanochemistry, Tokyo Polytechnic University Japan Johanna Saunier EA 401 Faculté de pharmacie France Christian Simon Laboratoire LI2C- CNRS UMR 7612, Université Pierre et Marie Curie France Renaud Souzy Laboratoire de Chimie Macromoléculaire, ENSCM-CNRS UMR 5076 France Steven H. Strauss Department of Chemistry, Colorado State University U.S.A. Masayuki Takashima Department Material Science and Engineering, Fukui University Japan

Shoichi Tsujioka Chemical Research Center, Central Glass Co., Ltd. Japan Pierre Turq Laboratoire LI2C- CNRS UMR 7612, Université Pierre et Marie Curie France Jan Uhlir Nuclear Research Institute Rez plc Czech Republic Jun-ichi Yamaki Institute of Material and Chemical Engineering, Kyushu University Japan Rachid Yazami Laboratoire d'Electrochimie et de Physico-Chimie des Matériaux et des Interfaces, INPG-CNRS UMR 5631 France Susumu Yonezawa Department of Material Science and Technology, Fukui University Japan Bin Zhu Chemical Engineering Technology/Chemical Reaction Engineering, Royal Institute of Technology (KTH) Sweden

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TABLE OF CONTENTS Preface Contributors Chapter 1: Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF H. Groult, C. Simon, A. Mantoux, F. Lantelme, and P. Turq

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1

Chapter 2: Applications of fluorinated carbon materials to primary and secondary lithium batteries T. Nakajima

31

Chapter 3: Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration Y. Ohzawa

61

Chapter 4: Electrochemical properties of fluorinated carbon nanotubes H. Touhara

89

Chapter 5: Fluorine-doped tin oxide electrodes for lithium batteries C.W. Kwon, H. Kim, T. Toupance, B. Jousseaume, and G. Campet

103

Chapter 6: Synthesis of fluorinated cathodes and fluoride electrolytes for lithium-ion batteries S. Yonezawa and M. Takashima

125

Chapter 7: Physicochemical properties of fluorine-containing electrolytes for lithium batteries D. Lemordant, F. Blanchard, G. Bosser, M. Caillon-Caravanier, B. Carré, A. Chagnes, B. Montigny and R. Naejus 137 Chapter 8: Fluorinated anions and electrode/electrolyte stability in lithium batteries R. Yazami and A. Martinent

173

Chapter 9: Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy) aluminate superweak anions B.G. Nolan, S. Tsujioka, and S.H. Strauss

195

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Table of contents

Chapter 10: Fluorinated electrolytes based on lithium salts of strong Brønsted acids O.E. Geiculescu, S.E. Creager, and D.D. DesMarteau

223

Chapter 11: Electrolytes for lithium batteries K. Kanamura

253

Chapter 12: Thermally stable fluoro-organic solvents for lithium ion battery Jun-ichi Yamaki

267

Chapter 13: Physical and electrochemical properties and application to lithium batteries of fluorinated organic solvents Y. Sasaki

285

Chapter 14: PVdF-based polymers for lithium batteries J.-Y. Sanchez, F. Alloin, and J. Saunier

305

Chapter 15: Lithium-ion-conductive polymer electrolytes exhibit a high lithium-ion transference number with the incorporation of fluorine atoms T. Abe and Z. Ogumi

335

Chapter 16: Room-temperature molten salts as new electrolytes R. Hagiwara and K. Matsumoto

349

Chapter 17: Fluorine-intercalated graphite for lithium batteries A. Hamwi, K. Guérin, and M. Dubois

369

Chapter 18: Battery application of graphite intercalation compounds Y. Matsuo

397

Chapter 19: Fluoride-based electrolytes and their applications for intermediate temperature ceramic fuel cells B. Zhu and B.-E. Mellander

419

Chapter 20: The use of Nafion® as electrolyte in fuel cells M. Odgaard

439

Chapter 21: Functional fluoropolymers for fuel cell membranes R. Souzy and B. Ameduri

469

Chapter 22: Films and powders of fluorine-doped tin dioxide H. Cachet

513

Table of contents

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Chapter 23: Doped transparent conducting oxides suitable for the fabrication of high efficiency thin film solar cells A. Bosio, N. Romeo, and V. Canevari

535

Chapter 24: Fluoride technologies application within the Molten-Salt Reactors fuel cycle J. Uhlir

549

Subject Index

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

1

Chapter 1

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF H. Groult, C. Simon, A. Mantoux, F. Lantelme, and P. Turq Laboratoire LI2C, CNRS UMR 7612, Université Pierre & Marie Curie- Case courrier 51, 4 place Jussieu, 75252 Paris cedex 05, France 1. PROPERTIES AND INDUSTRIAL USES OF FLUORINE GAS Fluorine gas is a yellowish, poisonous and highly corrosive gas which reacts with practically all organic and inorganic substances [1–6]; it reacts with all elements except helium, neon and argon to form ionic or covalent fluorides. Until World War II, there was no commercial production of elemental fluorine. However, due to the development of the atomic bomb and nuclear energy applications, the production of large amounts of elemental fluorine became necessary. Now, F2 gas can be considered as a necessary intermediate in uranium isotopic enrichment: separation of the isotopes of natural uranium is carried out by a diffusion process involving gaseous UF6. Uranium tetrafluoride (UF4) is first produced by the reduction of the oxide H2 or NH3 and fluorination by HF at 400–600°C according to UO3  H2 → UO2  H2O

(1)

UO2  4HF → UF4  2H2O

(2)

UF6 is then prepared by the reaction of fluorine gas and UF4 at a high temperature ( 1000°C) according to UF4  F2 → UF6

(3)

Uranium enrichment performed by gaseous diffusion or ultracentrifugation consists of increasing the content of natural uranium’s in isotope 235 to levels up to 3–5%. About 90% of the nuclear reactors currently in operation use this type of enriched uranium.

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In the coming decades, energy consumption is expected to rise significantly, due to dramatic increases in world population, coupled with the industrialisation of developing countries, notably in Asia and Latin America. Since the beginning of the 20th century, the emission of greenhouse gases has risen drastically because of the use of fossil fuels in industry and transportation. This emission is expected to increase considerably in the near future. Thus, to respect the climatic and environmental equilibrium, intense research programmes are being devoted to the development of renewable energies (solar, wind, etc.); however, the latter have, at the present time, a limited potential. Therefore, nuclear power, which could be considered as an environmentally safe form of energy since it does not produce harmful gases contributing to the greenhouse effect, should play an important role in the production not only of electricity, but also of other energy sources (hydrogen, etc.). For example, it generates 110 times less CO2 than natural gas and close to 240 times less CO2 than coal for electricity production. In the European Union, nuclear energy accounts for 35% of electricity production, thereby avoiding the emission of 300 million tonnes of CO2. In 2001, the installed capacity of all types of nuclear power plants in the world amounted to 358,000 MW, about a quarter in the United States (over 100,000 MW) and about 17% in France (almost 60,000 MW). Thirty-two reactors are under construction worldwide: 22 in Asia and 10 in central and eastern Europe. Fluorine gas is not only devoted to the synthesis of UF6, but is also widely used for the preparation of various fluorinated compounds involved in different industrial processes: WF6 for depositing tungsten on insulating or conducting substrates by CVD, NF3 for etching semiconductors, graphite fluorides (CFx, 0.5  x  1.24) for use as cathodes in primary lithium batteries and as lubricating agents, SF6 as insulating gas in electric devices, ClF3 to clean semiconductor fabrication vessels in the computer chip industry, CoF3 as solid fluorine carriers and F2–N2 mixture to strengthen the surface properties of plastics (impermeability, chemical resistance, barrier effect, etc.) or to control the fluorination of molecules in organic chemistry. Therefore, it seems to be of prime importance to study in detail the fluorine evolution reaction (FER) in order to optimise the process and to satisfy the increasing industrial requirement for this gas. The purpose of this review paper is to provide a brief overview of both experimental and theoretical aspects of the FER on carbon anode in KF–2HF. The process has been analysed of by correlating results deduced from electrochemical tests in KF–2HF, ex situ surface characterisations, notably by AFM and STM, and numerical calculations. Owing to molecular dynamics simulations of molten KF–2HF, the constituents of KF–nHF melt have been identified, depending on the temperature and the HF content in KF–nHF. The origin of the strong adhesion of fluorine bubbles on the surface is also discussed. Finally, the particular shape of fluorine bubbles generated on horizontal carbon anodes in KF–2HF

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

3

is studied from a theoretical point of view taking into account capillary forces between the electrode surface and the gas–liquid interface. 2. ANALYSIS OF THE PREPARATION PROCESS 2.1. Generalities

In 1886, the French scientist Henri Moissan prepared, for the first time, fluorine by electrolysis of anhydrous hydrogen fluoride containing a small amount of potassium fluoride in an electrochemical cell with platinum–iridium electrodes. KF was used to render HF conducting. The description of industrial cells is widely reported in the literature [1,2,7]; briefly, the cells operate at 6 kA and contain molten KF–2HF (40.8 wt% HF) with about 24 plate carbon anodes and steel or iron cathodes. Carbon anodes are used to avoid dissolution occurring with most metals in parallel with the evolution of fluorine; in addition, graphite must be avoided since exfoliation takes place due to co-intercalation of ionic species and fluorine gas between the lamellar graphene sheets. Monel skirts are also required to separate the hydrogen and fluorine gases formed at the cathode and the anode, respectively, and to avoid their explosive recombination. The global reaction involves HF decomposition: 2HF(liq) → F2 (g)  H2 (g)

(4)

The two corresponding half-cell reactions are supposed to involve the HF2 electrochemical species: HF2 →

1  2

F2  HF  e

(5)

at the anode, and 2 HF  e →

1  2

H2  HF2

(6)

at the cathode. First, molecular dynamic simulation of KF–nHF was investigated depending on the HF ratio and the temperature of the melt. 2.2. Molecular dynamics model for KF–nHF electrolytes

Within the framework of improvement of fluorine generation process, the model developed for KF–nHF electrolytes was intended to provide not only structural properties of the liquid (e.g. thermodynamical data, speciation, etc.) but also dynamical properties, especially transport coefficients such as self-diffusion coefficients, electrical conductivity and viscosity. This would make it possible, in principle, to address most of the previously cited problems on the microscopic scale.

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Among the available simulation techniques, ab initio molecular dynamics (AIMD) was the only one previously used to study KF–2HF [8]. While it is accurate, it needs very few hypotheses, and is computationally extremely intensive, making it unsuitable for the long simulations needed for transport properties computation. Lighter techniques such as Monte-Carlo simulations or integral equations based methods were also rejected because they were unable to predict transport properties. The only theoretical method left to investigate KF–HF electrolytes was therefore classical molecular dynamics (MD) [9]. This statistical mechanics-based simulation method is fast and provides all the desired results as soon as the model used therein is adequately defined. Transport properties have been examined with the herein presented iono-molecular model [10], but we will focus only on the structural features of the liquid. Molecular dynamics requires the description of the interaction forces between the liquid constituents. In pure molecular liquids, for example HF, this means defining the forces applied to each molecule by their surrounding molecular neighbours. In high-temperature molten salts, such as KF, this means defining the forces acting on each fluorine and potassium ion. But in KF–2HF, the microscopic nature of the constituents is still unknown: in electrochemistry, HF2 is often postulated, but there are no hints on how much HF are turned in that ionic form, how much in “neutral” HF. Oligomers are known to form in pure HF [11]. Polyfluorides, both centred F(HF)n and chained HnFn1 , have also been revealed in several condensed phases [12–15]. In fact, the question is whether KF–2HF is a molecular liquid or an ionic liquid. Is KF–2HF closer to pure HF or to molten KF? The point of view adopted is crucial for acidity definition. Up to now, the only attempt was based on Brönsted acidity [16], but why not try with Lewis fluoroacidity pF? This only depends on the nature of the species on the microscopic scale, which is in general known from direct observation. The paucity of available physical-chemistry data was an important obstacle to the definition of the model: to our knowledge, there are still no data on infrared, Raman spectroscopies, neutron scattering, and NMR spectroscopy (numerous corresponding data exist for the solid, but were not available for the liquid case). Therefore, the nature of the components was unknown, and a fortiori the forces acting between them. The only direct information about the liquid structure was given by AIMD [8], showing essentially centred polyfluorides and transitionally chained ones. We therefore decided to build an iono-molecular model to retrieve these F(HF)n, and their proportions, as a function of n. We denoted our model as iono-molecular since it describes, on the one hand HF molecules and K and F ions on the other. For alkali halides, the models developed by Fumi and Tosi [17] are indeed well tested and used to describe accurately the solid phases, pure molten salts, their mixtures and even aqueous solutions of the salts. For KF, we used the interactions given in Table 1.

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

5

The centred polyfluorides essentially consist of a fluorine ion solvated by HF molecules. In these F(HF)n complexes, the fluoride–hydrogen distance dHFranges from 1.35 to 1.7 Å depending on n, and is very similar to an ordinary H-bond. For pure HF, several models have been developed. Only one of them explicitly describes the H-bond between HF molecules (H-bond is implicit in the others and results from Coulomb or dipole interactions): the HF3 model derived by Klein and McDonald [18] (see Table 2). We therefore decided to use this model for our HF molecules, and to apply the H-bonding Morse potential (VHF) between HF molecules and fluoride. In addition to the attractive coulombic interaction between positively charged H and fluoride anions, this was intended to yield the right dHF shorter than the H-bond in pure HF (1.8 Å). The system was simulated at different temperatures ranging from 330 to 410 K (56.85 to 136.85°C). The initial simulation cell consisted of 384 atoms (64 times KF–2HF units) randomly arranged to get a density extrapolated, for each simulation temperature, from experimental data [19]. The coulombic forces were computed with an ordinary Ewald algorithm, the HF molecules were handled Table 1 Interaction potential for K and F ions Parameter







cij

1.25

1.00

0.75

σij (Å)

2.926

2.642

2.358

Cij (10-79 Jm6)

24.3

19.5

18.6

Dij(10-99 Jm8)

24.0

21.0

22.0

Note: Analytical form is Vij  ZiZje2\rbcij exp[B(σijr)] Cijr6Dijr8 with b  3.38 × 1019 J and B  2.96 × 106 m1.

Table 2 Interaction potential for HF molecules Pair

Potential

H-F

2(exp(10.6(r1.6))exp(5.6(r1.6)))

H-H

600exp(3.34r)

F-F

2  105exp(4.25r)VDISP

Note: Distances are in Å and energies in kcal mol1 with V  f(r)(220\r6  400\r8  4500\r10), where f(r)  exp((4.7/r1)2) if r 4.7 and f(r)  1 if r 4.7. DISP

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with a constraints algorithm, and the Verlet algorithm was used with a time step of 1.0 fs [10]. Simulation cells were equilibrated in 20,000 equilibration steps (i.e. 20 ps) to yield stationary averages for temperature, potential energy, and constant radial distribution functions. Simulations were then conducted in the micro-canonical ensemble. Visual examination of the cell after equilibration reveals the formation of polyfluorides (Fig. 1). This is a spontaneous formation of the complexes, since the initial configuration was randomly disordered. To investigate how many complexes F(HF)n of each n were formed during simulations, we had to choose a criterion to on which to base a decision on whether an HF molecule is bonded to a fluoride ion or not. This criterion is naturally derived from the radial distribution functions (RDF). In Fig. 2, the RDF for the hydrogen–fluoride pairs is plotted as a function of distance r. The first neighbour peak at 1.54 Å is perfectly in the range expected, clearly shorter than H-bonds in pure HF. Furthermore, the RDF decreases down to zero beyond this peak, between r  2.1 and 2.6 Å. This shows clearly that HF molecules, where the H atom is closer than 2.1 Å to a fluoride ion (or any cut-off distance up to 2.6 Å), are bonded to it. This allows one fluoride ions to determine how many molecules are bonded to each fluoride, and

Fig. 1. Instantaneous configuration of 2HF–KF liquid at 366 K. Grey spheres represent fluorine atoms (both molecular fluorine and fluoride), white spheres hydrogen atoms and black spheres potassium ions. Radii are set to the standard van der Waals radii.

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

7

Fig. 2. Radial distribution g(r) function for the H–F (fluorine ions) pair, in the liquid 2HF–KF at 366 K. The narrow peak at 1.54 Å results from the very strong bonding between fluorine and H (belonging to an HF molecule). g(r) decreases down to zero in the range 2.1–2.6 Å: closer pairs are bonded, more distant pairs are non-bonded.

to determine how many polyfluorides of each n exist. The results are reported in the next section. Following the above procedure, we obtained, at 366 K, for a simulation length of 2.6 ns, the average number of complexes F(HF)n for each n. The data obtained are displayed in Fig. 2, together with results from previous AIMD studies [8]. The results are very similar: both theoretical techniques reveal F(HF)n complexes up to n  4; the more common complex is F(HF)2 and not the simplest [FHF] . Fig. 3 shows only time-averaged results, but the complex populations undergo fluctuations during the simulation. It has been demonstrated that this is due to exchanges of HF molecules between complexes. This rapid exchange (on the simulation timescale) explains why the final observed populations do not depend on the starting random configuration: the system really reaches a dynamical chemical equilibrium. Fig. 4 illustrates this process: [F(HF)n…HF]   F(HF)m  F(HF)n  [FH…F(HF)m]  for each pair n,m.

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Fig. 3. Centred polyfluoride F(HF)n populations, depending on n. Black bars: results from our classical molecular dynamics simulations at 366 K. White bars: results from ab initio molecular dynamics by von Rosenvinge et al. Features of both distributions are the same.

Fig. 4. Exchange of HF molecules between two polyfluoride ions. One central fluorine is denoted by FA and the other by FB. The graphs show the distance between the transferred H atom (and therefore HF molecule) and the two centres, as a function of time.

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

9

To go beyond these results, other simulations were performed at higher temperatures up to 410 K and showed insignificant differences in complex populations. The absence of cell-size effect was also checked with simulation cell containing up to 1926 atoms (for details see Ref. [20]). Another factor potentially influencing the population of complexes is liquid composition, which will be discussed in the next section. Two other compositions of the liquid. KF–1.8HF and KF–2.4HF, were simulated using the same model from the RDF analysis, it appeared that the structure of these liquids is very similar to the one obtained for KF–2HF. It was therefore possible to apply the same procedure to determine populations of the different polyfluorides. Results are presented in Fig. 5. For the KF–1.8HF mixture, [FHF]  and F(HF)2 proportions are almost equal (but F(HF)2 is still predominant) and proportions of heavier polyfluorides are smaller than in KF–2HF. For the 2.4KF–HF mixture, F(HF)2 is also predominant but competes with F(HF)3 instead of [FHF] . In general, the relative proportions of heavier polyfluorides increase with increasing proportions of HF. This result seems quite intuitive but: ●

●

●

At the time the model was built such an effect was never observed and not even suggested. It was directly observed only later, in very different media such as tBu4NF/HF in solution and in cold freon mixtures [12]. In more similar media, such as (CH3)4NF/mHF, this behaviour is only suggested by more recent experiments.

Fig. 5. Centred polyfluoride F(HF)n populations, depending on n, for two different compositions. Black bars: 1.8 HF/KF; white bars 2.4 HF/KF. Distribution is shifted from low weight (less solvated fluorine) to higher weight (more solvated fluorine) while HF proportion is increased in the liquid.

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These results finally demonstrate the flexibility of the model and its ability to predict consistent and relevant structural results. To conclude, in spite of converging recent results about similar liquids, there is still no evidence for the existence of the polyfluorides in KF–2HF. We therefore tried to use NMR to address this problem, but HF exchange is too fast to have enough resolution on the spectra with current methods. Both theoretical and experimental work on similar systems with various cations is in progress to confirm the validity of the model and to get a more in depth understanding of the structure of these electrolytes. 2.3. Surface characterisations of carbon anodes fluorinated in KF–2HF by STM and AFM

The fluorine evolution process is characterised by a high current efficiency approaching 0.95 and also by a poor energy efficiency of about 0.3 [3–7, 21–23]. As a consequence, a large quantity of heat must be removed by cooling to maintain a constant temperature in the cell. The thermodynamic potential of HF decomposition is ca. 2.9 V [21,24], and not 1.8 V as reported in Ref. [25]. In fact, the total voltage in industrial cells is composed of five parts: the reversible decomposition voltage (2.9 V), the ohmic drop in the electrolyte, the ohmic drop in the electrodes, the cathode overvoltage, and the anode overvoltage. The high ohmic drop value is mainly due to the fact that a distance of several centimetres between anodes and cathodes is necessary to avoid explosive recombination of hydrogen and fluorine, in contrast to the NaOH-Cl2 preparation process in which the distance between anodes and cathodes is only a few millimetres. The high anode overvoltage (2.5 V) is commonly ascribed to the formation of a solid carbon–fluorine layer on carbon anodes during fluorine production [1–7, 21–23]. The inhibition of the FER is partly explained by the low surface energy of the film, which repels the electrolyte from the electrode. The contact angle is in the range of 120–160° [3]: fluorine bubbles have a lenticular form and are strongly adherent to the surface of the carbon anodes. This induces a significant decrease in the electroactive surface area of the electrode. Qualitative evidence of the formation of a passivating layer on the carbon anodes is given by cyclic voltammetry studies [7,26,27] and X-ray photoelectron spectroscopy (XPS) [3,5,6,27–30]. Indeed, the first voltammogram performed in KF–2HF with a new carbon anode exhibits an anodic passivation peak (Fig. 6) between 2.5 and 3.0 V vs. Pt–H2, which corresponds to the formation of a solid carbon–fluorine film (denoted C–F henceforth) on the anode surface. Many authors have concluded that, after reaching a high potential value in KF–2HF, the C–F film is composed of insulating graphite fluorides (denoted CFx) and it was assumed that the electron transfer occurs by tunnel effect through the passive CFx [31], the latter acting as an inhibiting barrier for the electron transfer during the FER. It leads to very low values of transfer coefficient, in agreement with a mechanism

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

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Fig. 6. I–E curves (v  0.4 V s1) obtained in KF–2HF with a graphite electrode.

involving electron tunnelling through a passive film. For such a mechanism, the probability of electron transfer depends on the thickness and the height of the potential barrier. In fact, XPS investigations performed with carbon anodes fluorinated in KF–2HF have shown that the position of the F1s and C1s peaks indicates the presence of ionic and semi-ionic C–F bonds [3,5]. Impedance measurements performed under “dry conditions”, i.e.without electrolyte (with a carbon/C–F/mercury structure) and in the presence of an aqueous solution containing a redox couple have revealed that the C–F film can be considered as an electronic conductor and thus cannot constitute a high-energy barrier for the electron transfer in FER. STM measurements on HOPG samples fluorinated at 6 V in KF–2HF [6,32] have confirmed these conclusions in revealing the presence of conducting GIC compounds, usually denoted CxF, with ionic and semiionic C–F bonds. In fact, two kinds of images (Fig. 7) were obtained with this technique: (i) In the first image, the same hexagonal symmetry as in pure HOPG (Fig. 7a) was observed, with a periodicity of 0.244 nm (Fig. 7b). Only half of the carbon atoms of a graphene layer exhibits high electronic density, due to the nonequivalence of the atomic sites resulting from the ABAB stacking of graphene layers. (ii) In the second image, all of the carbon atoms of the hexagonal rings were observed; this is due to the presence of fluorinated intercalants between two graphene sheets, which induce an increase in the distance between these two layers; as a consequence, the non-equivalence of the atomic sites resulting

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Fig. 7. Three-diamentional scanning tunnelling microscopy images. (a) Pure HOPG showing no overlap; z range, 15 nA/div. (b) Fluorinated HOPG (6 V) showing an overlap of electronic densities; z range: 5 nA/div. Experimental conditions : bias voltage 20 mV; constant height mode; Pt/Ir tip. (Images are reprinted from H. Groult et al., Electrochim. Acta, 44 (1999) 2793–2803.)

from the ABAB stacking of graphene layers is lost and all the carbon atoms of graphene layers are clearly visible. The spacing between two neighbouring atoms deduced from the corrugation amplitudes, is about 0.154 nm. In contrast two classical STM images observed in the case of pure HOPG, an overlap of the electronic densities between two neighbouring atoms was observed in the case of HOPG samples fluorinated in KF–2HF. The fluorination of the surface, coupled with the intercalation of fluorinated species between two graphene layers, induces the formation of carbon–fluorine bonds that modify the electronic density of each carbon atom. Therefore, the overlap of the electronic densities is due to two neighbouring F atoms. A schematic in-plane structural model corresponding to the image observed by STM in Fig. 7b is given in Ref. [32], showing the centred hexagonal lattice commensurate with the graphite lattice. STM observations on fluorinated HOPG have also revealed that in many parts of the electrode, no current can be detected even for very high bias value, i.e., no image can be recorded. This was attributed to the local presence of insulating compounds such as graphite fluorides (CFx). Finally, the C–F solid film generated at the surface of the carbon anode during the electrolysis of molten KF–2HF was supposed to be composed of conducting compounds belonging to the graphite intercalation compounds (GICs) family in which the C–F bonds are ionic and semi-ionic and insulating CFx with covalent C–F bonds. In spite of the presence of graphite fluorides, the fluorine evolution mechanism does not obey a mechanism involving electron tunnelling through a passive film as reported previously [31]: the electrons can be easily transferred from the electrolyte to the electrode. Nevertheless, no evidence was reported that proved effectively the presence of CFx in the C–F layer. This is why AFM measurements in contact

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

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mode were performed with HOPG samples fluorinated at 6 V in KF–2HF. In addition to the AFM characterisations, local electrical measurements were performed in ambient air with an original laboratory-made device derived from an AFM; this apparatus has been developed by Houzé [33–35] and allows one to cover nine decades of tip/sample resistance value (from 100 Ω to 100 GΩ). The fluorinated HOPG samples previously studied by STM were analysed on 5 μm  5 μm areas (Fig. 8). The electrical and topographical images are represented in Fig. 8a and b, respectively; the corresponding distribution of the electrical resistance measured is given in Fig. 8c. As shown in Fig. 8c, the resistance distribution is centred on 3  104 and 6  1010 Ω average values; for the former, most of the contact resistance values are measured between 8  103 and 2  105 Ω, associated with the red colours on the electrical images (Fig. 8b); for the resistance distribution centred on 8  1010 Ω, most of the contact resistance values are measured between 5  1010 and 2  1011 Ω, associated with the purple colours on the electrical images (Fig. 8b). Notice that the histogram also

(a)

(b)

(c)

Fig. 8. AFM investigations on HOPG fluorinated in KF-2HF at 6 V: (a) topographical image; (b) electrical image; (c) distribution of the measured resistance deduced from (b).

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reveals a continuous variation of the local resistance with two small shoulders around 107 and 4  108 Ω. These images clearly show that the surface of fluorinated HOPG samples does not appear to conduct uniformly. Also note that the presence of less conducting zones does not correspond to transitions between the two graphene sheets. As reported in the literature, graphite fluorides and chemically prepared graphite intercalation compounds [36–43] present a very large difference in terms of electrical resistivity. For instance, graphite fluoride with a composition of CF0.465 exhibits an electrical resistivity of about 107 Ω cm (about 104 Ω cm for the starting graphite material), whereas this value can reach about 105 Ω cm in the case of fluorine-GICs [36]. Based on this fact, we interpret the differences of resistivity observed on AFM images with HOPG samples fluorinated in KF–2HF (Fig. 8c) in the presence of GICs (with a resistance distribution centred on 3  104 Ω), graphite fluorides (with a resistance distribution centred on 6  1010 Ω), and intermediate compounds for which the composition varies from CFx to GICs and which give rise to intermediate colours (Fig. 8b) from red to purple. Therefore, these AFM measurements coupled with local electrical determinations confirm our previous assumptions deduced from STM measurements about the heterogeneity of composition of the C–F layer on carbon anodes fluorinated during electrolysis in molten KF–2HF. 2.4. Origin of the strong adhesion of F2 bubbles on the carbon surface

The CFx compounds in the C–F film a strongly influence the wettability of the electrode by KF–2HF and the kinetics rates of the FER. For example, we have shown [6] that in the case of carbon anode chemically fluorinated at high temperature under fluorine gas atmosphere prior to its introduction into the electrochemical cell, leading to a high amount of graphite fluorides on the surface, the fluorine evolution reaction is completely inhibited in KF–2HF. Here, we compare fluorine bubble evolution on carbon (Fig. 9a) and nickel (Fig. 9b) ring

Fig. 9. Evolution of fluorine in KF–2HF on (a) horizontal carbon and (b) nickel electrodes facing the top of the electrochemical cell (E  4 V vs. Cu/CuF2).

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

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disk electrodes; in both cases, only the upper part of the surface was electroactive, the electrodes being embedded into an insulating Teflon holder. As shown in Fig. 9a, the large fluorine bubble formed on the carbon anode has a lenticular form and is strongly adherent to the surface; almost the entire surface is covered by the gas and the contact angle is about 150°. A parallel drawn can be made with the electrolytic preparation of aluminium: when the alumina content of the cryolite melt decreases, a concomitant increase in the contact angle is seen. For a low amount of alumina, a rapid change in the wetting properties occurs and the contact angle jumps to around 180°. This results in a full coverage of the electrode by the gas phase followed by a dramatic increase in the overvoltage and the appearance of sparks around the electrode. This phenomenon, called “anode effect,” is due to the presence of solid fluorocarbon CFx compounds at the surface of the carbon electrode; the solid–liquid interfacial energy is very low and the liquid cryolite does not wet the anode any more. As similar explanation can be given in terms of wettability during the electrolysis of molten KF–2HF with carbon anodes, but no “anode effect” is seen. In contrast, in the case of a nickel electrode used for F2 evolution in KF–2HF (Fig. 9b), small, spherical fluorine bubbles with a zero contact angle form at the centre of the electrode; nevertheless, an important modification occurs on the lateral part of the electrode at the nickel/Teflon interface. Indeed, in this area, due to the presence of a perfluorinated carbon compound such as Teflon, the fluorine bubble that forms on nickel spreads over the adjacent coating (Teflon) and produces large and non-symmetric fluorine bubbles. In other words, the juxtaposition of two compounds with very different interfacial properties drastically modifies the shape of the fluorine bubble and its formation. In the case of fluorinated carbon anodes, STM and AFM measurements have revealed the presence of GIC and CFx; therefore, an analogy can be made with the phenomenon observed at the nickel/Teflon interface. The fluorine bubbles should form easily in the part of the electrode covered by GICs, but since these zones are located near the zones covered with CFx compounds, the fluorine bubbles spread over the surface covered by CFx and give rise to a large and strongly attached fluorine bubble. 2.5. Origin of the CFx compounds in the C–F surface film

Graphite fluorides used in primary lithium batteries or as lubricants are usually prepared by the chemical fluorination of carbon at high temperatures ( 350°C) using F2 gas. Depending on the operating conditions, the x value in CFx can vary from ~ 0.5 to ~ 1.3. As reported above, their electrical conductivity drastically decreases with increasing x values. In the case of the electrolysis of KF–2HF, CFx can also be generated from the electrofluorination of graphite oxides initially present on carbon surface. Indeed, these oxides are known to be easily fluorinated at low temperatures because of their high instability. For increasing potentials applied to the carbon electrodes in KF–2HF, oxygen might

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be easily exchanged with fluorine, to give rise to large amounts of insulating graphite fluoride [44,45]. For example, graphite oxide compounds such as CzO(OH)y could lead to the formation of CFx according to the reaction CzO(OH)y  (2  xz  3y) F → zCFx  (1  y) F2O  yHF  (2  xz  3y)e

(7)

C–O bonds could result to the presence of water in the electrolyte. In contrast, in the classical schematic representation of the carbon surface (Fig. 10a), several kinds of C–O bonds are present at the edges of graphene sheets and can make bridges between two stacks of graphene sheets. F atoms can also be easily exchanged with O atoms to form CF, CF2 and CF3 groups with covalent C–F bonds, as illustrated in Fig. 10b. In this case, very large insulating areas can be formed, explaining why very large insulating domains are observed by AFM with HOPG samples fluorinated in KF–2HF (Fig. 8c). Physicochemical characterisations of starting carbon materials have been made to determine if the model used for the surface representation (Fig. 10a) and our assumptions given above for the formation of large insulating domains are valid. The XRD pattern of the carbon used as anode is presented in Fig. 11. The bump generally observed around 20° due to the presence of amorphous phases is

Fig. 10. Schematic representation of carbon surface (a) before and (b) after electrochemical fluorination in KF–2HF.

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

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(002)

40 counts

(100) et (101)

0

20

40

(004)

60

80

100

120

140

2θ / deg.

Fig. 11. XRD pattern of carbon material.

not evidenced in the XRD pattern; in addition, the (002) diffraction line belonging to graphite is clearly visible, indicating a preferential orientation along the c-axis. Additionally, three minor contributions are pointed out due to the (100), (101) and (004) diffraction lines of graphite, but with a low intensity for each one. As a consequence, one may conclude that the carbon used for the preparation is similar to graphite but with a lower crystallinity. The distance, d002 between two graphene sheets, was determined to be from 0.34 nm the Bragg relation taking into account the position of the (002) diffraction line: the value very slightly higher than that usually observed for pure graphite (0.3354 nm). The crystallite size Lc along the c-axis was calculated from the broadness of the (002) diffraction line using Scherrer’s equation: 0.89λ1 Lc   B cos θ

(8)

where λ1 is the wavelength of the Kα1Co beam, θ the Bragg angle and B the angular full-width at half-maximum (FWHM) of the (002) diffraction line. It was found that Lc was close to 40 nm. Similar information has been provided by transmission electron microscopy analysis; as shown in Fig. 12, the carbon material is mainly composed of graphitised phases without clear orientation. Finally, Raman spectroscopy analysis was performed; this technique provides information about functional groups or chemical bonds in molecules. In a Raman spectrum, each line has a characteristic polarisation, and therefore, polarisation data provide information about the molecular structure. Graphitic carbons are strong Raman scatterers in spite of their intense optical absorption.

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Fig. 12. TEM image of carbon material.

Because of the weak interlayer bonding, disorder along the c-axis can occur in the graphite crystal, while at the same time, the strong interlayer C–C bonding maintains a high degree of order within the individual carbon sheets. Therefore, Raman spectra of carbon materials are usually characterised by a pair of bands called G- and D-bands. The G-band is assigned to the E2g2 carbon–carbon stretching mode, whereas the D-band is due to an A1g vibration mode in the disordered region of carbon materials or edge plane of powder carbon [46,47]. The ratio, R, of intensity of D-band, ID, to that of G-band, IG (R  ID/IG) depends on the structure of the carbon and indicates the degree of disordering of the surface of the carbon materials. In the presence of amorphous phases, an additional peak should be observed at around 1530 cm1. Fig. 13 shows the Raman spectrum of the starting carbon material. The G- and D-bands are observed at around 1600 and 1355 cm1. The R value obtained from our experimental spectra is close to 1.7. This value is very large compared with that usually observed for other kinds of carbonaceous materials: 0.07 for pyrolitic graphite, 0.20 for polycrystalline graphite and 0.76 for coke-type carbon. It indicates a high disordering at the surface of our carbon materials. It is also important to note that the experimental curve can be fitted by considering only the two contributions due to the G- and D-bands; in other words, it is not necessary to take into account an additional contribution at around 1530 cm1 due to the presence of amorphous phases to obtain a very good fitting of the experimental spectra. These conclusions are in good agreement with the results deduced from XRD analysis (no evidence of amorphous phases).

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

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Fig. 13. Raman spectra of carbon electrode.

To conclude, XRD and TEM analyses have pointed out the presence of graphitised phases without a clear orientation; these phases are necessary for the formation of GICs during the electrolysis of molten KF–2HF. However, Raman spectroscopy has shown that the carbon materials used for fluorine production are characterised by a high surface disordering, which means that according to the classical representation of the carbon surface described in Fig. 10a, many C–O bonds are initially present on the surface prior to electrolysis. These C–O bonds lead to the formation of insulating CFx compounds during electrolysis. The latter are not wetted by the melt and the electrode surface area involved in the FER is drastically limited. Some possibilities are available to limit the formation of CFx [5,7,48–54]. Typical examples are described briefly in the next section. 2.6. Improvements in the fluorine evolution process

For limiting the influence of CFx and improving the kinetics of the FER, surface treatments can be proposed. We present here one specific example that consists in performing an electrochemical activation of the carbon anode in KF–2HF at very high potential (40 V) during a short time period (1 min). The aim of this treatment is to burn vigorously the C–O groups and thereby avoid the formation of large amounts of CFx. The I–E curves recorded in KF–2HF (v  0.2 V s1) with a crude and an activated carbon electrode are shown in Fig. 14: if one considers a current density of 12 A dm2, the anodic overvoltage is decreased to 0.4 V in the case of activated carbon. One must notice that activation has two effects on the surface behaviour: first, the decrease in the amount of CFx on the surface leads to an increase in the electroactive surface area. Second, the activation of

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480

2

420

j (mA/cm²)

360 300

1

240 180 120 60 0 3.5

4.0

4.5

5.0

5.5

6.0

E (V) vs Cu/CuF2

Fig. 14. I–E curves (v  0.4 V s1) obtained in KF–2HF. (1) carbon after electrochemical passivation in KF–2HF; (2) carbon after electrochemical activation in KF–2HF at 40 V during 1 min.

the carbon anode in KF–2HF induces a smoothing and cleaning of the surface as evidenced recently by SEM [55]. Consequently, the porosity of the surface is modified and the influence of pores on the kinetics rate is decreased. Indeed, as described in a previous paper [55,56], fluorine is produced on the horizontal carbon surface and in the pores. In the latter, the walls stop the lateral growth of the bubble and the phenomenon that occurs is comparable to capillary rise with convex meniscus. The volume of the gaseous cap increases along the z-axis concomitantly with the decrease in the electroactive surface area. When the pore is filled with fluorine, the gas bubble spreads on the horizontal surface all around the pore; the same phenomenon occurs in the nearest pores and recombination of the gaseous bubbles is observed until a total coverage of the electrode is reached. Activation at high potential makes it possible to limit the influence of pores on the kinetics rate of the FER. The positive effect of activation procedures on the kinetics rate of the FER is clearly evidenced by faradaic impedance measurements performed in KF–2HF [57] at 4.2 V with crude and activated carbon anodes (Fig. 15): as shown in this figure that presents the impedance diagrams obtained in the Nyquist representation, the charge transfer resistance (⬇ diameter of the semi-circle) measured at medium frequency in the case of an activated carbon electrode (R2) is about nine times lower than that observed in the case of a crude carbon electrode (R1), indicating a faster kinetics in the case of activated carbon electrodes. Activation of laboratory-scale electrodes is a good, realisable procedure; however, such a procedure cannot be achieved on a large scale in industrial cells because: (i) the procedure is risky, since sparks are observed, and produces a high

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

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- Im(Z) (ohm cm2 )

60 50 40 30

1

20 10

2

0 0

20 R2

40

60

Re (Z) (ohm

cm2)

80

100 R1

Fig. 15. Impedance diagram in the Nyquist representation obtained in KF–2HF with carbon after electrochemical (1) passivation in KF–2HF; (2) activation in KF–2HF at 40 V for 1 min; E  4.2 V vs. Cu/CuF2.

quantity of heat that should be eliminated, and (ii) a special electrolyser is needed to perform such a pre-treatment of carbon anodes. 2.7. Interfacial properties and gas bubble formation

Generally, bubble evolution obeys a classic mechanism: nucleation, growth, coalescence, detachment and rise of bubbles, and is also related to their size and adherence to the electrode surface [58–60]. Practically, these processes occur concurrently, so, it is virtually impossible to separate them experimentally. As shown above, the interfacial properties also play an important role in the present electrochemical reaction. The size and adherence of bubbles depend on the properties of the liquid–gas interface. For many gas evolution reactions, the contact angle is nearly zero and the weakly adherent bubbles have a spherical shape. The electrode surface is not strongly modified in the presence of bubbles and most of the works deal with the influence of the bubble on the fluid motion in the vicinity of the electrode [61]. As shown in Fig. 9a, the fluorine bubble generated on a horizontal carbon surface electrode facing the top of the electrochemical cell has a particular shape; the fluorine gas coverage does not hinder the current flow even if a complete coating of the electrode by a fluorine gaseous film is observed and the volume of the gas bubble continues to increase. To explain this phenomenon and also to understand exactly the origin of the high anodic overvoltage that characterises this process, the influence of the mass transfer on the kinetics rate has been studied recently by impedance measurements using a rotating disk electrode [55]. The interpretation of the impedance diagrams obtained in the Nyquist representation allows us to propose a new model for the representation of the electrode/electrolyte interface as illustrated in Fig. 16, including the presence of an intermediate layer sandwiched between the

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Fig. 16. Schematic representation of the carbon/KF–2HF interface. Cross section of an horizontal carbon electrode facing the top of the electrochemical cell. A “fluidized” layer in which liquid KF–2HF and fluorine gas co-exist is sandwiched between the C–F film and F2 bubble. A gradient of concentration of fluorine in the solid C–F layer up to the surface of the carbon electrode is observed.

C–F surface layer and the gaseous phase. The current passes through this conductive layer at the periphery of the bubble where the gas layer is very thin. It should be pointed out that this hypothesis presents some similarity to the explanation of Brandon and Kelsall [62] in the case of the bubble departure radii of H2, Cl2, and O2 evolution using microelectrodes. They have proposed that a thin liquid film of electrolyte separates the gas and solid phases. Jennings et al. [63] have mentioned the existence of a mixed phase at the electrolyte/electrode interface composed of electrolyte and gas. Nevertheless, compared with fluorine bubbles, the shape of these three gas bubbles is completely different and their detachment easier. In the case of fluorine evolution, the intermediate conducting layer is supposed to be composed of a mixed phase comprising liquid KF–2HF and fluorine. Owing to this model, we are also able to propose an explanation for the origin of the high anodic overvoltage: it is due to the C–F film and to the intermediate conducting layer composed of a mixture of KF–2HF and nascent fluorine, giving rise to ηC–F (activation overvoltage for the FER) and ηfluid (ohmic drop in the conducting layer), respectively; ηT  ηC–F  ηfluid. For instance, it has been shown that, for a resistivity of the fluidised layer ρ ⬇ 10 Ω cm, a temperature of the melt θ  95°C, and a radius of electrode r0  0.8 cm, ηfluid is close to 1 V, comparable with ηC–F ⬇ 1.5 V. Therefore, the contribution of the “fluidised” layer must be taken into account for a good understanding of the fluorine evolution process.

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

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Let us now consider the mechanism of bubble growth and detachment to understand the origin of the shape of the F2 bubble formed on the horizontal surface facing upward (Fig. 9a). Visual examination [55,56] shows that a few small gas bubbles remain at the electrode surface just after the detachment of a fluorine bubble. These bubbles very rapidly coalesce, and a unique bubble grows and covers the electrode surface. Then, the bubble swells. A bulge appears in its centre and gives rise to a spherical bubble, which detaches through the electrolyte. A quantitative analysis is carried out considering that the current arising from the uncovered part of the electrode is calculated from the current density. The current arising from the conducting layer beneath the bubble is calculated as described in a previous paper [55]. The areas of the covered and uncovered surfaces depend on the gas volume and on the shape of the bubble. The amount of gas is readily calculated from the electrolysis current. At first, we have assumed that the curvature radius, R, at the gas–liquid interface is constant. Consequently, the bubble shape is a spherical cap. In situ observations have shown that the bubble is not exactly a spherical cap: the periphery of the bubble is flattened and the adherent bubble has the shape of a flying saucer (Fig. 9a). Therefore, to obtain a more realistic bubble shape, the model has been improved by considering that such a bubble shape is due to the presence of capillary forces between the electrode surface and the gas–liquid interface [64]. The increment of pressure, pint, due to the interface curvature, R, obeys the following equation [65]:

γ GL pint  ppub pext  2  R

(9)

where pbub is the pressure inside the bubble, pext the external pressure and γGL the gas–liquid surface tension. A phenomenological approach is used taking into account a simplified assumption: the capillary pressure at the gas–liquid interface is supposed to be proportional to the reverse of distance, hcap (Fig. 17): K pcap   hcap

(10)

where hcap  h  δ. The thickness, δ, of the conducting layer has been estimated from in situ observations [66] coupled with the exploitation of impedance diagrams [55]; δ  0.3 mm. K is the proportionality constant, which is homogenous to surface tension. In order to obtain an accurate description of the phenomenon, the hydrostatic pressure, phyd, was also introduced into the model phyd  hhyd dg

(11)

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Fig. 17. Scheme for the pressure balance at the gas–liquid interface.

where d is the density of the electrolyte and g the acceleration due to gravity. The distance, hhyd, to the electrolyte surface is determined from the distance, hsur, between the surface of the conductive layer and electrolyte surface: hhyd  hsur  h. The external pressure is pext  patm  phyd

(12)

where patm is the atmospheric pressure. The pressure inside the bubble obeys the equation pbub  patm  phyd  pcap  pint

(13)

In the present situation, R varies all along the gas–liquid interface. At a distance close to the electrode, the term pcap is large and the value of pint can be negative, which corresponds to a concave curve. In this case, according to Eq. (9), R is negative. At time t the volume of the bubble obeys nF2 RT V  pbub

(14)

with nF being the amount of gas generated by the electrolysis (in mol). 2 However, the gas bubble no longer has the shape of a spherical cap. To solve the problem, a step-by-step process was used. The interface profile was divided into small distance intervals, Δs. The coordinates of a point i1 were deduced from the coordinates xi and hi of the point i. At point i, the curvature radius, Ri, was determined from the internal pressure, pint (Eq. (13)).

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

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As an example, the profile for a bubble containing 1.8 μmol of fluorine is shown in Fig. 18a. The following parameters were used: patm  1 atm, hsur  1 cm, d  1.98 g cm3, g  981 cm s2, γGL  750 dyn cm1, K  740 dyn cm1 and pbub  1.0086 atm. Then, the bubble growth during electrolysis was studied using a stepby-step procedure. The time was divided into small time intervals, Δt, and the current at time t was deduced from the current arising from the covered and uncovered areas of the electrodes [64]. However, when the radius of the covered area, rcov, tends to the electrode radius, rD, the current still passes in the form of a very thin layer at the periphery of the bubble. The additional resistance, Radd, depends on the thickness of the layer and is assumed to be proportional to tan θ1. The angle θ1 is the value of θi in the external ring and is equal to the contact angle. In the model [64], the following empirical equation is used:

λρ Radd   tan θ1 2πr

(15)

where λ is a proportionality constant. Since the electrode surface is fully covered, the fluorine gas generated by the electrolysis induces a swelling of the bubble. The pressure inside the bubble and the contact angle increase, so that the calculated volume of the bubble fits the experimental volume deduced from the amount of gas. The bubble growth leads to the formation of a spherical excrescence, which soon detaches from the electrode surface under the action of the hydrostatic pressure. During that last phase, due to the formation of the evolving bubble, the pressure remains nearly constant. It results in a current plateau just before the bubble detachment.

1

d

h / cm

c 0.5

b a

0 0.3

0.6

0.9

1.2

x / cm

Fig. 18. Calculated profile for bubbles of various volumes: (a) 55 mm3 , (b) 220 mm3, (c) 460 mm3, (d) 700 mm3.

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The various parameters introduced in the model are adjusted in order to obtain simulated curves, which fit the experimental curves. The shape of the fluorine bubble at various stages of electrolysis is shown in Fig. 18. In contrast to the flat bubble obtained at the beginning of the growth (Fig. 18a), the bubble height just before the detachment (Fig. 18d) is now greater than the radius of the contact disk (bubble electrode). 3. CONCLUSIONS The uses of fluorine gas in various industrial fields confer an attractive aspect to this compound. In this work, a brief overview of the FER in molten KF–2HF is given in considering both a theoretical and experimental approach. First, the simulation of KF–nHF using molecular dynamics was presented to identify the species present in the electrolyte depending on the HF/KF ratio and the temperature of the melt. It has been shown that, under our experimental conditions (composition: KF–2HF; θ  95°C), the major species are presum-ably F(HF)2. Nevertheless, direct experimental evidence to confirm the validity of the model and to provide an deeper understanding of the structure of this electrolyte is welcome. Further comparisons with similar systems (such as EMIMF-2.3HF) both on the theoretical and experimental levels are in progress. Then, AFM investigations coupled with local electrical measurements in ambient air on HOPG fluorinated in KF–2HF allowed us to validate previous assumptions on the presence of both conducting GICs and insulating CFx on the carbon surface. Owing to these observations, coupled with physicochemical characterisations by XRD, TEM and Raman spectroscopy, the origin of the strong adhesion of F2 bubble on carbon surface was elucidated. Graphitised phases randomly oriented and C–O bonds are present at their edges; these C–O bonds are easily changed during electrolysis of molten KF–2HF to form insulating and non-wetted CFx compounds (CF, CF2 and CF3 groups). One example for improving the kinetics of the fluorine evolution process is presented; it consists of the activation at very high potential in KF–2HF, the aim is to limit the generation of CFx to enhance the wettability of the electrode, thereby increasing the electroactive surface area. Finally, the particular shape of the fluorine bubble evolving on a horizontal electrode surface upward of the electrochemical cell was discussed; in that frame, an original approach was made in considering the characteristic features of the process: (i) strong adherence of fluorine gas at the electrode surface, (ii) existence of an electrolytic current even when the electrode is fully covered with the gas film, and (iii) flat bubble with a nearly null contact angle. The experimental observation pointed to the existence of a thin conducting layer beneath the gas bubble. The variable

Experimental and theoretical aspects of the fluorine evolution reaction on carbon anodes in molten KF–2HF

27

curvature radius of the bubble is the consequence of a skin effect due to the influence of capillary forces between the electrode surface and the gas–liquid interface. ACKNOWLEDGEMENTS The authors thank Pr. D. Devilliers, Drs. S. Durand-Vidal, F. Nicolas, B. Morel, J.-P. Caire, C. Belhomme, F. Houzé, R. Baddour-Hadjean, F. Warmon and MM. A. Colisson, M. Combel, and M. Vogler for helpful and fruitful discussions and assistance in the experimental work. We also acknowledge the Comurhex-Cogema Company (Pierrelatte, France) for their joint support of this research project. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13]

[14] [15] [16] [17] [18] [19]

A.J. Rudge, “Production of elemental fluorine by electrolysis” in Industrial Electrochemical Processes, A.T. Kuhn (Ed.), Amsterdam, 1971, pp. 1–69, Chap. 1. P.T. Hough, W.V. Childs, and T. Fuchigami (Eds.), The ECS Proceeding Series, Pennington, 1997, PV 97-15, p. 113. N. Watanabe, T. Nakajima, and H. Touhara, Graphite Fluorides, Vol. 8, Elsevier, Amsterdam, 1988, pp. 1–22, Chap. 1. T. Nakajima, Fluorine-Carbon, and Fluoride-Carbon Materials, T. Nakajima (Ed.), M. Dekker, New York, 1995, pp. 1–31, Chap. 1. H. Groult, D. Devilliers, and M. Vogler, Current Topics in Electrochemistry, Research Trends, Vol. 4, Poojapura, Trivandrum, India, 1997, pp. 23–39. H. Groult, J. Fluorine Chem., 119 (2003) 173. D. Devilliers, M. Chemla, and T. Nakajima (Ed.), Fluorine-Carbon and Fluoride-Carbon Materials, M. Dekker, New York, 1995, pp. 283–331,Chap. 8. T. von Rosenvinge, M. Parrinello, and M.L. Klein, J. Chem. Phys., 107(19) (1997) 8012. M.P. Allen and D.J. Tildesley, Computer Simulation of Liquids, Oxford University Press, Oxford, 1987. C. Simon, T. Cartailler, and P. Turq, Phys.-Chem. Chem.-Phys., 3 (2001) 3119. D. Deraman, J.C. Dore, J.G. Powles, J.H. Holloway, and P. Chieux, Mol. Phys., 55(6) (1985) 1351. J.D. Forrester, M.E. Senko, A. Zalkin, and D.H. Templeton, Acta Crystallogr., 16 (1963) 58. I.G. Shenderovich, S.N. Smirnov, G.S. Denisov, V.A. Gindin, N.S. Golubev, A. Dunger, R. Reibke, S. Kirpekar, O.L. Malinka, and H.-H. Limbach, Ber. Bunsenges. Phys. Chem., 102 (1998) 422. R. Hagiwara, T. Hirashige, T. Tsuda, and Y. Ito, J. Electrochem. Soc., 149(1) (2002) D1. A. Tasaka, Y. Shodai, S. Kohara, and M. Inaba, Proceedings of Third French–Japanese Seminar on Fluorine in Inorganic Chemistry and Electrochemistry, Paris, 2003, p. 9. H. Dumont, S.Y. Qian, and B.E. Conway, J. Appl. Electrochem., 27 (1997) 267. F.G. Fumi and M.P. Tosi, J. Phys. Chem. Solids, 25 (1964) 31. M.L. Klein and I.R. McDonald, J. Chem. Phys., 71(1) (1979) 298. A.I. Semerikova and A.F. Alabyshev, Russian J. Phys. Chem., 35 (1961) 2791.

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H. Groult et al. C. Simon, T. Cartailler, and P. Turq, J. Chem. Phys., 117(8) (2002) 3772. O.R. Brown, Electrochim. Acta, 25 (1980) 367. H. Imoto, T. Nakajima, and N. Watanabe, Bull. Chem. Soc. Jpn., 48 (1975) 1633. L. Bai and B.E. Conway, J. Appl. Electrochem., 18 (1988) 839. D. Devilliers, F. Lantelme, and M. Chemla, J. Chim. Phys., 76 (1979) 428. H. Wendt and G. Kreysa, Electrochemical Engineering, Science and Technology in Chemical and Other Industries, Springer, Berlin, 1999, pp. 290–369. Chap. 11. D.M. Novak and P.T. Hough, J. Electroanal. Chem., 144 (1983) 121. O.R. Brown, B.M. Ikeda, and M.J. Wilmott, Electrochim. Acta, 32 (1987) 1163. H. Groult, D. Devilliers, M. Vogler, C. Hinnen, P. Marcus, and F. Nicolas, Electrochim. Acta, 38 (1993) 2413. L. Bai and B.E. Conway, J. Appl. Electrochem., 20 (1990) 916. P. Cadman, J.D. Scott, and J.M. Thomas, Carbon, 15 (1977) 75. M. Chemla and D. Devilliers, J. Electrochem. Soc., 136 (1989) 87. H. Groult, D. Devilliers, S. Durand-Vidal, F. Nicolas, and M. Combel, Electrochim. Acta, 44 (1999) 2793. F. Houzé, R. Meyer, O. Schneegans, and L. Boyer, Appl. Phys. Lett., 6(13) (1996) 1975. J.P. Kleider, C. Longeaud, R. Brüggemann, and F. Houzé, Thin Solid Films, 383 (2001) 57. S. Guessab, L. Boyer, F. Houzé, S. Noël, and O. Schneegans, Synth. Metals, 118 (2001) 121. N. Watanabe, T. Nakajima, and H. Touhara, Graphite Fluorides, Vol. 8, Elsevier, Amsterdam, 1988, pp. 240–261, Chap. 8. T. Mallouk and N. Bartlett, J. Chem. Soc. Chem. Commun., (3) (1983) 103. R. Hagiwara, M. Lerner, and N. Bartlett, J. Chem. Soc. Chem. Commun., (9) (1989) 573. Y. Sato, T. Kume, R. Hagiwara, and Y. Ito, Carbon, 41 (2003) 351. A. Hamwi, M. Daoud, and J.C. Cousseins, Synth. Metals, 30 (1989) 23. R. Yazami, P. Hany, P. Masset, and A. Hamwi, Mol. Cryst. Liq. Cryst., 310 (1998) 397. T. Nakajima, Y. Matsuo, B. Cemva, and A. Jesih, Carbon, 34 (1996) 1595. A. Tressaud, F. Moguet, S. Flandrois, M. Chambon, G. Guimon, G. Nanse, E. Papirer, V. Gupta, and O.P. Bahl, J. Phys. Chem. Solids, 57(6–8) (1996) 745. T. Nakajima and N. Watanabe, Graphite Fluorides and Carbon-Fluorine Compounds, CRC Press, Boca Raton, 1991, pp. 155–171, Chap. 7. T. Nakajima and M. Touma, J. Fluorine Chem., 57 (1992) 83. F. Tunistra and J.L. Koenig, J. Chem. Phys., 53 (1970) 1126. D.S. Knight and W.B. White, J. Mater. Res., 4 (1989) 385. T. Nakajima, T. Ogawa, and N. Watanabe, J. Electrochem. Soc., 134 (1987) 8. D. Devilliers, B. Teisseyre, and M. Chemla, Electrochim. Acta, 35 (1990) 153. P.T. Hough and D.M. Novak-Antoniou, US Patent 4 602 985, 1986. N. Watanabe, M. Inoue, and S. Yoshizawa, J. Electrochem. Soc. Jpn., 31 (1963) 113. T. Tojo and T. Nakajima (Ed.), Fluorine-Carbon and Fluoride-Carbon Materials, M. Dekker, New York, 1995, pp. 333–354, Chap. 9. Asahi Glass Co., JP-Kokai Patent 58 81 981, 1983. O.R. Brown and M.J. Wilmott, European Patent 255 225, 1988. H. Groult, D. Devilliers, F. Lantelme, J.–P. Caire, F. Nicolas, and M. Combel, J. Electrochem. Soc., 149 (2002) E485. H. Groult and F. Lantelme, J. Electrochem. Soc., 148 (2001) E13. F. Nicolas, H. Groult, D. Devilliers, and M. Chemla, Electrochim. Acta, 41 (1996) 911. S. Lubetkin, Electrochim. Acta, 48 (2002) 357.

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H. Vogt, Electrochim. Acta, 42 (1997) 2695. F. Lantelme, D. Diamanacos, and J. Chevalet, Electrochim. Acta, 23 (1978) 717. G. Kreysa and M. Kuhn, J. Appl. Electrochem., 15 (1985) 517. N.P. Brandon and G.H. Kelsall, J. Appl. Electrochem., 5 (1984) 475. D. Jennings, A.T. Kuhn, J. Stepanek, and R. Whitehead, Electrochim. Acta, 20 (1975) 903. [64] F. Lantelme and H. Groult, J. Electrochem. Soc., 151(5) (2004) D121. [65] J.T. Davies and E.K. Rideal, Interfacial Phenomena, Academic Press, San Diego, 1961. [65] H. Roustan, Ph.D. Thesis, ENSEEG-INPG, Saint Martin d’Hères, France, 1998.

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 2

Applications of fluorinated carbon materials to primary and secondary lithium batteries Tsuyoshi Nakajima Department of Applied Chemistry, Aichi Institute of Technology, Yakusa-cho, Toyota-shi 470-0392, Japan E-mail: [email protected] 1. INTRODUCTION Reaction of carbon materials with elemental fluorine yields two kinds of intercalation compounds: graphite fluorides, (CF)n and (C2F)n, and fluorine – graphite intercalation compound, Cx F [1–4]. Graphite fluorides with puckered (sp3) graphene layers to which fluorine atoms are covalently bonded are synthesized by the fluorination of various carbon materials at high temperatures of 300 to 600°C. For example, (CF)n is prepared from petroleum coke at 300–600°C and from natural or synthetic graphite at ca. 600°C. (C2F)n is obtained from high crystalline graphite in a limited temperature range of 350–400°C. Fluorination of a graphite between 400°C and ca. 550°C yields a mixture of (CF)n and (C2F)n. In a lowtemperature range less than ca. 100°C, fluorine – graphite intercalation compound, Cx F is synthesized in the presence of Lewis acid(s) such as HF. Cx F has planar (sp2) graphene layers with ionic or semi-ionic (semi-covalent) C–F bond. Synthesis of Cx F is usually performed under a fluorine atmosphere in the presence of Lewis acid(s) or in anhydrous liquid HF (aHF) with elemental fluorine. In most cases, Cx F is prepared at room temperature. (C2F)n and (CF)n are black and graywhitish in color, respectively, and both are electric insulators because of their C–F covalent bonds. The black color of (C2F)n may be due to a trace amount of Cx F type sp2 carbon. On the other hand, Cx F is black and an electric conductor because intercalated fluorine atoms are mobile at stage 2 or higher stages and somewhat mobile even at stage 1 in which semi-ionic (semi-covalent) C–F bond exists. Graphite fluoride, (CF)n, was used as the cathode material of the primary lithium battery, as solid lubricant, as water repellent and so on. Among them, the most important application is the use as a cathode material in primary lithium battery with metallic lithium anode and organic solvents [1,2]. Li/(CF)n primary

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Tsuyoshi Nakajima

battery has been practically used for many years. It has a high discharge voltage between 2 and 3 V, high discharge capacity (⬃800 mAh/g), high-energy density, high safety and long shelf life. Cx F is also a new candidate as cathode for primary lithium battery, having a higher discharge potential than that of (CF)n because fluorine atom semi-ionically bonded to graphene layer has a higher activity than covalently bonded one [2–10]. A disadvantage of Cx F is that its discharge capacity is less than that of (CF)n due to the lower fluorine content of Cx F. Cx F with the highest fluorine content (x ⬇ 2) is usually prepared in aHF with elemental fluorine [2,3,11–15] or using elemental fluorine in the presence of gaseous HF and IF5 [2–4,7–10]. Cx F with a higher fluorine content than C2F was recently synthesized by the fluorination of graphite using high oxidation state transition metal complex fluorides and elemental fluorine under pressure in aHF [16–18]. Recent research interest is mainly on the materials for secondary (rechargeable) lithium batteries, which use carbonaceous anodes, transition metal oxide cathodes and organic electrolyte solutions. Application of fluorination techniques and various fluorides to secondary lithium batteries is an interesting research subject. Polyvinylidenefluoride (PVDF) and LiPF6 are currently used as a binder and electrolyte for practical lithium ion secondary batteries, respectively. Imide and methide salts containing CF2/CF3 groups are new candidates as electrolytes, because of their high thermal and electrochemical stability although they cause corrosion of aluminum cathode current collector [19,20]. Light fluorination is one of the effective methods of surface modification for transition metal oxide cathodes [21] and carbonaceous anodes [22–27]. Surface treatment of lithium cobalt oxide by elemental fluorine increased the capacity and improved the cycleability [21]. Surface fluorination of natural graphite samples with different particle sizes increased the capacities without any decrease in the first coulombic efficiencies [22–25]. In case of petroleum cokes, first coulombic efficiencies of graphitized petroleum cokes were improved [26,27]. Corrosion of aluminum cathode current collector occurs in the solvents containing fluoro-organic electrolytes as mentioned above. The corrosion of aluminum is inhibited by fluorination [19,28]. The present chapter deals with recent results on the synthesis, structures and electrochemical behavior of highly fluorinated graphite as a cathode of primary lithium battery and those of surface-fluorinated graphites and petroleum cokes as anodes of secondary lithium battery. 2. INFLUENCE OF COINTERCALATED HF ON THE DISCHARGE BEHAVIOR OF HIGHLY FLUORINATED GRAPHITE AS A CATHODE OF PRIMARY LITHIUM BATTERY Fluorine – graphite intercalation compound, Cx F is usually synthesized by several different methods at room temperature in the presence of Lewis acid(s) [2–4]. Typical synthetic methods are a gas/solid reaction using graphite and

Applications of fluorinated carbon materials to primary and secondary lithium batteries

33

elemental fluorine in the presence of a small amount of gaseous HF [2–4,29] or in the presence of gaseous HF and IF5 [2–4,7–10], and a reaction of graphite with elemental fluorine in aHF [2–4,11–15]. The gas/solid reaction using graphite and elemental fluorine gives stage 1 and higher stage compounds with composition of ⬃C3F. The reaction of graphite with elemental fluorine in aHF or with HF and IF5 yields mainly stage 1 compounds with the higher fluorine contents of C4F–C2F. The reaction of graphite with high oxidation state transition metal complex fluoride in aHF also provides stage 1 Cx F with by-products being insoluble in aHF [30]. Stage 1 compounds with higher fluorine contents than C2F are synthesized by the reaction of graphite with high oxidation transition metal complex fluoride and elemental fluorine under pressure in aHF at room temperature [16–18]. In the present section, the synthesis, structure and discharge behavior of highly fluorinated stage 1 Cx F compounds are described [16–18]. 2.1. Synthesis and structure of highly fluorinated graphite, Cx F

When Cx F is synthesized in aHF, some amount of HF is cointercalated into graphite because the fluorination reaction proceeds via CxHF2, which is formed at the beginning of intercalation reaction of fluorine into graphite in aHF (Eq. (1), step 1). Mobile HF2  easily diffuses into graphite, and HF is gradually desolvated along with the formation of semi-ionic C–F bond at stage 1 (Eq. (1), step 2). Cointercalated HF thus remains in graphene layers after the formation of semi-ionic C–F bond. HF remaining in graphene layers is removed from Cx F by pumping. However, the complete removal of HF is usually difficult (Eq. (2)). Stage 1 Cx F sometimes contains stage 2 and 3 phases as minor components, where C–F bond is nearly ionic and the main intercalated species are HF2 [2–4,11–15]: xC  1/2F2  yHF → CxHF2(HF)(y1) → Cx F(HF)y

(1)

Cx F(HF)y → Cx F(HF)z  (yz)HF (by pumping)

(2)

Highly fluorinated graphite, Cx F, was synthesized by the reaction of graphite with high oxidation state transition metal complex fluoride (K2PdF6, K2MnF6, K2NiF6 or KAgF4) and elemental fluorine under pressure ((3.9–11.8)  105 Pa) in aHF at room temperature [16–18]. The raw products contained by-products such as PdF2/PdF3, MnF2/MnF3, NiF2 or AgF2 insoluble in aHF. PdF2/PdF3 or MnF2/MnF3 were removed from Cx F by adding KF, fresh HF and elemental fluorine to the raw product and exposing the mixture to UV light for 2 days and washing away as K2MF6 (M, Pd or Mn) soluble in aHF with an excess of KF in the form of yellow solution [18]. Solid by-product, NiF2 or AgF2, was separated from Cx F by adding aHF and AsF5 to the raw product and washing off as M(AsF6)2 (M: Ni or Ag), soluble in aHF in the form of yellow solution [16].

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Tsuyoshi Nakajima

Composition and X-ray diffraction data of Cx F samples are summarized in Table 1 [16–18]. The C/F ratio was in the range of 3.7–1.2, decreasing with increasing fluorine pressure from 3.9  105 to 11.8  105 Pa, and it was usually the lowest by the use of KAgF4 as a fluorinating agent. Hydrogen (H) was Table 1 X-ray diffraction data and composition of Cx F samples prepared by high oxidation state complex fluorides and elemental fluorine in anhydrous HF at room temperature Sample

Composition

F2

Fluoride

5

X-ray diffraction data a

(x10 Pa)

Stage

Ic (nm)

a0 (nm)

1

C3.7F(HF)1.2

5.9

K2PdF6

1

0.629

0.245

2

C2.6F(HF)0.68

5.9

K2PdF6

1 (2)

0.582 0.925

0.246

3

C2.7F(HF)0.41

7.9

K2MnF6

1

0.599,0.571

0.245

4

C2.9F(HF)0.65

7.9

K2MnF6

1

0.597

0.246

5

C2.3F(HF)0.31

3.9

K2NiF6

1 2 3

0.634 0.929 1.236

0.245

1 (2) (3)

0.678 0.927 1.245

0.245

1 (2) (3)

0.644 0.933 1.231

0.246

1 (2) (3)

0.673 0.921 1.226

0.245

6

7

8

C1.7F(HF)0.12

C1.9F(HF)0.58

C1.5F

7.9

7.9

7.9

K2NiF6

K2NiF6

K2NiF6

9

C1.6F

11.8

K2NiF6

1

0.647

0.247

10

C1.6F

11.8

K2NiF6

1

0.680

0.247

11

C1.5F(HF)0.35

11.8

K2NiF6

1 (2) (3)

0.639 0.942 1.240

0.247

12

C1.2F

11.8

KAgF4

1

0.631

0.247

13

C1.3F

11.8

KAgF4

1

0.626

0.247

14

C1.4F

11.8

KAgF4

1

0.628

0.248

a

( ) represents minor phase.

Applications of fluorinated carbon materials to primary and secondary lithium batteries

35

detected in some samples by elemental analysis, mainly existing as HF in them. As shown by IR spectra later, all the samples contained small amounts of HF. Fluorination of graphite by elemental fluorine in aHF usually yields Cx F with a high fluorine content. The maximum C/F ratio is approximately 2 [11–15]. The use of high oxidation state transition metal complex fluoride and elemental fluorine under pressure provides highly fluorinated graphite with C/F ratio less than 2 although H was detected in some samples by elemental analysis. Among Cx F samples in Table 1, those prepared using KAgF4 (samples 12–14) had the highest fluorine contents. Their composition would be in the range of 1 x 2, except traces of HF cointercalated in them. All the Cx F samples were composed of stage 1, or stage 1 with stage 2 and 3 phases as minor components, as given in Table 1. The repeat distances along the c-axis (Ic) of the stage 1 phases were in the range 0.57–0.68 nm. Most of the Ic values were larger than 0.60 nm, as given in Table 1. The smallest Ic value of stage 1 Cx F was reported to be 0.47 nm for C6F [12], which means that semi-ionically bonded fluorine atoms form a single intercalated layer between two graphene layers due to the low in-plane density of fluorine. With increasing in-plane density of fluorine, the single fluorine-intercalated layer gradually changes to the double rows where intercalated fluorine atoms are in contact with each other between two graphene layers. The Ic of stage 1 thus increases to a value higher than 0.6 nm with increasing fluorine content. Cointercalation of HF into graphite also contributes to the increase in the Ic values. The lattice parameters along the a-axis (a0) were in the range of 0.245–0.248 nm. The a0 values were 0.245 nm for the Cx F samples containing stage 2 and 3 phases and stage 1 compounds with relatively low fluorine contents. Pure stage 1 compounds (samples 9, 10, and 12–14) and C1.1F (sample 11) prepared under high fluorine pressure (11.8  105 Pa) had greater a0 values of 0.247–0.248 nm than that of graphite (0.246 nm). The lattice parameter a0 is directly correlated with the carbon – carbon bond length, i.e. it is proportional to C–C bond length. When the carbon – fluorine bonding of Cx F is ionic or nearly ionic (usually stage 2 or higher stage), the C–C bond length is slightly shorter than that of graphite itself due to the electron transfer from graphite to intercalated fluorine [3]. This is usually observed in a bond length of an acceptor-type graphite intercalation compound. However, the C–C bond is longer than 0.1421 nm of graphite lattice in a highly fluorinated stage 1 Cx F due to an increase in the covalent nature of C–F bond [3]. The formation of semi-ionic C–F bond at stage 1 causes the localization of electrons and may slightly change the sp2 nature of graphene layers. It was thought that stage 1 Cx F with a composition of ⬃C2F kept planar sp2 graphene layers. However, the increase in the C–C bond length and partial puckering of graphene layers may occur in a highly fluorinated phase in which the fluorine content is C2F or higher than C2F (C2F–C1F). The nature of C–F bonding of Cx F samples prepared using K2NiF6 and KAgF4 was evaluated by X-ray photoelectron spectroscopy (XPS) [16]. The

36

Tsuyoshi Nakajima

peaks for C1s and F1s electrons were observed at high binding energies, probably due to the charging effect caused by electron localization. For Cx F samples prepared using K2NiF6, C1s spectra showed two strong peaks at 287.7 and 290.2 eV, a medium peak at 291.8 eV and a weak one at 294.0 eV on average. The C1s peak at 287.7 eV corresponds to carbon atom unbound to fluorine. The peak at 290.2 eV is attributed to nearly semi-ionic C–F bond. Those observed at 291.8 and 294.0 eV are covalent C–F bond and –CF2 group. Corresponding to these C1s peaks, two F1s peaks were observed at 689.2 and 690.9 eV, indicating nearly semi-ionic and covalent C–F bonds, respectively. The Cx F samples prepared using KAgF4 had two strong C1s peaks at 287.7 and 290.2 eV and a strong F1s peak at 689.2 eV. In addition to this peak, a very weak F1s peak was present at 686.4 eV, indicating nearly ionic C–F bond. The lack of C1s peaks at 291.8 and 294.0 eV and F1s peak at 690.9 eV shows that fluorination degree is weaker in the Cx F prepared using KAgF4 than in the Cx F prepared with K2NiF6, which coincides with the fact that K2NiF6 is a stronger fluorinating agent than KAgF4. Highly fluorinated stage 1 compounds consist of several different phases as shown by IR spectra (Fig. 1) [18]. Table 2 summarizes the assignment of IR absorption peaks, C–C bond of graphene layers and possible Cx F phases [18]. The absorptions observed at 1084 cm1 and between 1123 and 1134 cm1 are assigned to the stretching vibration of semi-ionic C–F bond [13], probably arising from stage 1 C3F and C2F phases with sp2–sp3 graphene layers, respectively, in comparison with the composition in Table 1. The absorptions at 1225 and 1230 cm1 are due to C–F stretching vibration of covalent C–F bond with sp3 hybridized orbital, being observed for graphite fluoride, (CF)n [31,32]. A new absorption at 1196 cm1 may be assigned to the stretching vibration of nearly Table 2 IR absorption peaks (cm1) and assignments for Cx F samples Samplea

Assignment

4

7

9

11

14

1084

–

–

–

–

1123 –

Semi-ionic C–F

Graphene

Possible

layer

Cx F phase

sp2–sp3 2

C3F

1125

1126

1131

1134

Semi-ionic C–F

sp –sp

C2F

–

–

1196

1196

Nearly covalent C–F

sp2–sp3

C1F

1225

1225

1230

1230

1257

1250

1257

1256

1256

sp2

1524

1524

1525

1524

1523

sp2

1570

1575

1570

–

–

Sample number is the same as given in Table 1.

Covalent C–F

A2u

sp

3

1225

a

3

sp2

(CF)n

Applications of fluorinated carbon materials to primary and secondary lithium batteries

37

Fig. 1. IR absorption spectra of Cx F samples. 4, 7, 9, 11 and 14 in the figure: sample numbers in Table 1.

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covalent C–F bond in somewhat puckered graphene layers with the composition of C1F because this absorption has a larger vibration energy than that of C2F. It was recently reported that the absorption observed in the range 1250–1257 cm1 were due to C–C stretching in sp2 graphene layers of Cx F containing HF [14]. This absorption disappeared and a new absorption appeared at 1220 cm1 when HF was completely removed from Cx F [14]. It suggests that complete desolvation of HF from Cx F gives rise to the puckering of graphene layers by shortening the C–F bond length, i.e. the formation of graphite fluoride, (CF)n, with puckered sp3 graphene layers. The absorptions between 1250 and 1257 cm1 thus show that all the Cx F samples in Table 1 contain trace amounts of HF. In addition to above absorptions, two absorptions were observed in the range of 1570–1575 and 1523–1525 cm1. The absorptions at 1570 and 1575 cm1 are located at slightly lower wave numbers than those of graphite, 1580 cm1, being attributed to A2u mode indicating C–C stretching of graphene layers [33,34]. Since the absorptions have slightly weaker vibration energies than those of graphite itself, they may be derived from Cx F phase with relatively lower fluorine content such as C3F–C2F. The absorptions at 1523–1525 cm1 possess more weaker vibration energies corresponding to longer C–C bond length in sp2 graphene layers, probably arising from Cx F phase with a higher fluorine content, x ⬇ 2. The Cx F with higher fluorine content than C2F was prepared by the fluorination using high oxidation state transition metal complex fluorides and elemental fluorine under pressure in aHF. The IR absorption data indicate that the Cx F samples in Table 1 are composed of several different phases. Since graphite is polycrystal, the fluorination degree may not be uniform in crystallites constituting a graphite particle. If C3F or C2F phase with planar graphene layers exists in the same crystallite with graphite fluoride (CF)n having puckered graphene layers, such a structure may be unstable due to a high structural strain. Graphite fluoride (CF)n and stage 1 Cx F are both stable compounds even at high temperatures [1–4,35]. It is suggested therefore that CxF and (CF)n with planar and puckered graphene layers consist of different crystallites from each other, although some transient states from sp2 to sp3 structure may exist. The lattice parameter (a0)values in Table 1 were obtained by X-ray diffraction as average values for several different phases. 2.2. Influence of cointercalated HF on the discharge behavior of highly fluorinated graphite, Cx F

Discharge curves of Cx F samples were obtained in 1 mol/dm3 LiClO4–propylene carbonate (PC) solution at 25°C. Figs. 2–4 show the discharge curves obtained at a current density of 10 mA/g [16,18]. The discharge potential and capacity are governed by fluorine content, the nature of C–F bond and the amount of HF cointercalated in graphite. The discharge potential was at 3.1–3.2 V vs Li/Li at first and moved to a plateau at ca. 1.5 V as shown in Fig. 2. Since the electrochemical reduction of PC occurs at around 1 V, two plateaus observed at

Applications of fluorinated carbon materials to primary and secondary lithium batteries

39

3.1–3.2 and 1.5 V indicate the reduction of two different fluorine species. Sample 11 in Fig. 2 seems to have a larger amount of HF than other Cx F samples in Figs. 2–4 because it had the larger amount of H and stage 2 and 3 phases. Sample 11 had a long plateau at about 1.5 V, finally approaching 1 V. The same and short plateaus were also observed in samples 5 and 6 in Fig. 3. Cx F samples contain the semi-ionic, covalent and ionic C–F bonds as already mentioned. The ionic fluorine, HF2 , is in a more reduced state as anion than the semi-ionic and covalent fluorine atoms. It suggests that the equilibrium potential of F(semi-ionic and covalent)/F couple is higher than that of HF2 /F. Therefore, the discharge potential

Fig. 2. Discharge curves of Cx F samples, obtained at 10 mA/g in 1 mol/dm3 LiClO4-PC at 25°C. 9 and 11 in the figure: sample numbers in Table 1.

Fig. 3. Discharge curves of Cx F samples, obtained at 10 mA/g in 1 mol/dm3 LiClO4-PC at 25°C. 5, 6, 8 and 10 in the figure: sample numbers in Table 1.

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at 3.1–3.2 V can be assigned to the reduction of semi-ionic and covalent fluorine atoms and that at 1.5 V to the reduction of ionic fluorine, HF2 . The Cx F samples having H and stage 2 and 3 phases contained larger amounts of HF than others, and IR absorption spectra indicated that HF was present even in the samples in which no H was detected by elemental analysis. Samples 5, 6, 8 and 11 with small amounts of HF and stage 2 and 3 phases showed gradual decrease in the potentials with discharge. Stage 1 samples 9 and 10 exhibited the similar discharge curves. The discharge capacities of samples 5, 6 and 8–11 were in the range 260 to 580 mAh/g at a cut-off potential of 1.5 V vs. Li/Li. Samples 2, 3 and 4 also showed similar discharge curves with capacities of 420, 620 and 570 mAh/g, respectively. When K2NiF6, K2MnF6 and K2PdF6 were used as fluorinating agents, most of the Cx F samples contained H and/or stage 2 and 3 phases, which suggests that relatively larger amounts of HF were cointercalated in these samples, compared with those prepared using KAgF4 (samples 12–14). The a0 values of samples 1–11 in Table 1 were in the range 0.245–0.247 nm, and most of them were 0.245 or 0.246 nm, which also suggests that these samples contain ionic species, HF2 . On the other hand, samples 12, 13 and 14 gave flat discharge potentials at 3.2 V until 400–450 mAh/g at a current density of 10 mA/g and decreased to 1 V as shown in Fig. 4. They had higher fluorine contents without stage 2 and 3 phases and larger a0 values of 0.247–0.248 nm, which suggests that the amounts of HF and HF2  in samples 12–14 were less than those in samples 1–11. The influence of cointercalated HF was more clearly observed on the discharge at high current densities as shown in Fig. 5 [18]. The discharge capacity of Cx F containing a higher amount of HF was largely decreased at a high current density. The discharge capacity of sample 8 was significantly decreased from

Fig. 4. Discharge curves of Cx F samples, obtained at 10 mA/g in 1 mol/dm3 LiClO4-PC at 25°C. 12–14 in the figure: sample numbers in Table 1.

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Fig. 5. Discharge curves of Cx F samples, obtained at 35 mA/g (sample 8) and 40 mA/g (sample 14) in 1 mol/dm3 LiClO4-PC at 25°C. 8 and 14 in the figure: sample numbers in Table 1.

460 mAh/g at a current density of 10 mA/g to 275 mAh/g by the discharge at 35 mA/g, while that of sample 14 was slightly reduced from 515 mAh/g at 10 mA/g to 500 mAh/g at 40 mA/g at the cut-off potential of 1.5 V. The discharge capacities of samples 8–10, obtained at a low current density of 10 mA/g and at a cut-off potential of 1.0 V, reached 93–96% of their theoretical values calculated from the composition. This shows that cointercalation of HF in Cx F causes the decrease in the discharge capacities at both low and high current densities (10 and 35–40 mA/g, respectively). The discharge potential and capacity are affected by several factors such as fluorine content, nature of C–F bond and the amount of cointercalated HF. Cx F with a high fluorine content generally shows a large discharge capacity. The Cx F samples in Table 1 possessed three kinds of C–F bonds, i.e. semi-ionic, covalent and nearly ionic C–F bonds with composition between C2F and C1F. Discharge reaction of graphite fluoride with covalent C–F bond proceeds with the formation of an intermediate discharge product composed

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of C, F, Li and solvent molecules, which is metastable, determining an open-circuit voltage (OCV) of Li/(CF)n cell [1,2,36,37]. The intermediate discharge product finally decomposes to carbon (C), LiF and solvent molecules (S). (CF)n  nLi·S  ne→ nC···F···Li···S → nC  nLiF  nS

(3)

The discharge reaction would be basically similar to that of graphite fluoride; however, the reaction rate of discharge changes depending on whether HF coexists or not: Cx F  Li·S  e → Cx···F···Li···S → x C  LiF  S (fast)

(4)

Cx F  HF  Li·S  e → Cx···F···HF···Li···S → x C  LiHF2  S (slow) (5) CxδHF2δ  Li·S  e → Cx···F···HF···Li···S → x C  LiHF2  S (slow) (6) The discharge capacities at the cut-off potential of 1.5 V vs. Li/Li strongly depended on the amounts of cointercalated HF and coexistence of stage 2 and 3 phases. The C–F bonding in stage 2 and 3 phases is nearly ionic, and the main intercalated fluorine species are mobile HF2δ. All the samples in Table 1 contained small amounts of HF as shown by IR absorption spectra even in case that H was not detected by elemental analysis. The cointercalation of HF in stage 1 phase also suggests the coexistence of a trace of HF2δ. The discharge capacity was larger in stage 1 Cx F samples containing smaller amount of HF without stage 2 and 3 phases as already shown. If HF does not coexist in Cx F, the intermediate discharge product, C x ···F···Li···S would smoothly decompose to carbon, LiF and solvent molecules because of a strong interaction between F and Li due to a large surface charge of a small fluoride ion (Eq. (4)). Therefore the discharge reaction may proceed well without potential decrease for Cx F containing a less amount of cointercalated HF. On the other hand, when HF is contained in Cx F, decomposition of the intermediate discharge product, C x···F···HF···Li···S, may be slow because of a weak interaction between HF2 and Li due to a small surface charge of a large HF2 anion (Eq. (5)). Slow decomposition of the intermediate discharge product would cause the gradual decrease in the discharge potential, leading to the reduction of the discharge capacity at the cut-off potential of 1.5 V. This is more clearly observed at a high current density. The discharge capacities of samples 8 and 10 were 460 and 423 mAh/g at a current density of 10 mA/g, decreasing to 275 and 175 mAh/g at 35 mA/g, respectively. Therefore, complete formation of stage 1 phase with semi-ionic or covalent C–F bond and thorough removal of HF from a prepared sample are needed to obtain a flat discharge potential and high capacity. For this purpose, KAgF4 seems the best fluorinating agent. Chemical diffusion coefficients of Li in the intermediate discharge product were obtained by the impedance measurements. For all samples, the diffusion

Applications of fluorinated carbon materials to primary and secondary lithium batteries

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Table 3 Chemical diffusion coefficients of lithium ion F/C ratio

DLi (1012 cm2/s)

8

0.65

13

10

0.62

11

13

0.80

4.4

14

0.73

5.7

Samplea

a

Sample number is the same as given in Table 1.

coefficients were nearly constant irrespective of the discharge ratios. The data are given in Table 3 [18]. Samples 8 and 10 may have contained larger amounts of HF than samples 13 and 14, because sample 8 had stage 2 and 3 phases and discharge capacities of samples 8 and 10 at 35 mA/g (275 and 175 mAh/g, respectively) were much smaller than those of samples 13 and 14 at 40 mA/g (440 and 500 mAh/g, respectively), although the fluorine contents were somewhat higher in samples 13 and 14. Large HF2 anion would have a weaker interaction with Li cation than small F anion. Therefore HF2and Li would be mobile in the intermediate discharge product. This may be the reason why the larger diffusion coefficients were obtained for samples 8 and 10 than samples 13 and 14. 3. CHARGE/DISCHARGE BEHAVIOR OF SURFACE-FLUORINATED CARBON MATERIALS AS ANODES OF SECONDARY LITHIUM BATTERY Since electrochemical redox reactions occur at the surface of a solid electrode, surface structure is one of the decisive factors for determining the electrode performance. Surface modification is effective for improving the electrochemical characteristics of carbonaceous electrodes for lithium ion secondary battery. Electrode characteristics are governed by crystallinity, surface area, surface pore volume distribution, surface chemical species such as oxygen and so on. Several methods were applied to improve electrochemical behavior of carbon materials. They are surface oxidation [39–41], surface fluorination [22–27], thin metal coating [42], and carbon coating [43–47]. Light oxidation of carbon materials caused increase in their capacities by forming nanochannels at the surface while strong oxidation degraded surface structure, leading to increase in the irreversible capacity [39–41]. Carbon coating is an effective method for increasing the capacity and first coulombic efficiency [43–47]. In the present section, the effect of surface fluorination of natural graphite [22–25] and petroleum coke [26,27] on their electrochemical characteristics is described.

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3.1. Electrochemical behavior of surface-fluorinated natural graphite 3.1.1. Composition and surface structure change of natural graphite by fluorination

Natural graphite samples with average particle sizes of 7, 25 and 40 μm (abbreviated NG7, NG25 and NG40, respectively) were used as starting materials. Surface fluorination was performed between 150 and 500°C by elemental fluorine (purity: 99.4–99.7%) of 3  104 Pa for 2 min in a nickel reactor [22,23,25]. The same natural graphite samples were also subjected to plasma fluorination using CF4 [24,25]. Tables 4 and 5 show fluorine contents obtained by elemental analysis and surface fluorine concentration by XPS, respectively [22–25]. The fluorine contents in the samples fluorinated between 150 and 350°C were less than 1 at%, i.e. 0.2–0.6 at% while those fluorinated between 350 and 500°C had the larger values, i.e. 0.6–4.7 at%. The results coincide well with the fact that only the surface of graphite is fluorinated between 150 and 300°C [1–3]. Surface fluorine concentrations obtained by XPS showed the same trend, being in the range of 3.6–12.0 at% for the samples fluorinated between 150 and 350°C and in the range of 8.1–33.1 at% for those fluorinated between 350 and Table 4 Fluorine contents of surface-fluorinated graphite samples, obtained by elemental analysis Fluorination

Fluorine content (at%)

condition

NG7

NG25

NG40

Fluorinated by F2 150–300°C 350–500°C

0.5–0.6 0.6–2.2

0.3–0.4 0.6–4.7

0.2–0.4 1.7–4.2

Plasma-fluorinated

0

0.3

0.3

0 within detection limit ( 0.2 at%).

Table 5 Surface fluorine concentrations of fluorinated graphite samples, obtained by XPS Fluorination condition

Surface fluorine concentration (at%) NG7

NG25

NG40

Fluorinated by F2 150–300°C 350–500°C

6.0–10.4 11.3–14.6

4.5–6.9 8.1–28.9

3.6–12.0 14.3–33.1

Plasma-fluorinated

6.7–8.8

7.1–11.5

3.3

Applications of fluorinated carbon materials to primary and secondary lithium batteries

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500°C. Fluorine content and surface fluorine concentration were both increased not only with increasing fluorination temperature but also with increasing particle size of graphite, i.e. with decreasing surface area of graphite. On the other hand, fluorine contents were lower in plasma fluorination than the fluorination by elemental fluorine, i.e. ⬃0.3 at%. Plasma-fluorinated samples exhibited rather constant surface fluorine concentrations, less than 10 at% in most cases. Small amounts of surface oxygen were detected by XPS, i.e. 1.0–3.7 at% for natural graphite samples fluorinated between 150 and 500°C, and 1.0–2.8 at% for plasma-fluorinated samples. Table 6 shows the BET surface areas, indicating the significant increase in the surface areas by the fluorination using elemental fluorine at 250 and 350°C and plasma fluorination using CF4 [23–25]. The increments in surface areas were ⬃77% and ⬃55% for the samples fluorinated by elemental fluorine and plasma treatment, respectively. The surface pore volume distribution was also changed by surface fluorination. Fluorination reduced the mesopores with diameters greater than 2–3 nm, and increased those with diameters of 1.5–2 and 2–3 nm. These surface structure changes may have been caused by carbon–carbon bond breaking due to the strong fluorination reactions. Fluorination of natural graphite powder accompanying carbon–carbon bond breaking also induces the increase in the surface structural disorder, which is detected by Raman spectroscopy. Carbon materials exhibit two Raman shifts at 1580 and 1360 cm1. The Raman shift observed at 1580 cm1 is based on the in-plane stretching vibration derived from graphitic structure (E2g2 mode, Gband), and that at 1360 cm1 indicates the A1g vibration mode due to the disordered structure and/or edge of carbon particles (D- band) [33,34]. Surface fluorination enhanced the D- band intensity in any case. The intensity ratio of two Raman shifts, R (ID/IG), shows the degree of surface disordering of carbon Table 6 Surface areas of natural graphite and surface-fluorinated samples, obtained by BET method Surface area (m2/g)

Fluorination condition

NG7

NG25

NG40

Original graphite

4.79

3.71

2.94

Fluorinated by F2 150°C 250°C 350°C

5.61 7.65 8.48

5.16 5.18 6.09

3.50 4.90 4.95

Plasma-fluorinated (60 min)

7.42

4.71

3.69

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Tsuyoshi Nakajima

materials. The R values calculated from the intensity ratios of two Raman shifts are given in Table 7, which indicate that R values were increased by surface fluorination, in particular, with increasing fluorination temperature and particle size of graphite powder [22–25]. Plasma fluorination also increased the R values; however, the values decreased with increasing duration of plasma treatment. This means that the surface-disordered parts are somewhat eliminated by the extended plasma treatment. Fig. 6 shows transmission electron microscopic images of graphite samples fluorinated by elemental fluorine [23,25]. The fluorinated layers exhibit disordered structures in contrast with the unreacted graphene layers. The thickness of fluorinated basal plane is about 3–4 nm, corresponding to 5–7 fluorine-intercalated layers [1–4]. The surface of edge plane is also disordered with a similar thickness. The nature of C–F bonding is evaluated by XPS. Graphite samples fluorinated by elemental fluorine showed C1s and F1s peaks, indicating C–F bond at 288.5 and 687.7 eV on average, respectively. These binding energies indicate that the C–F bonding of surface-fluorinated graphite samples is in an intermediate state between the semi-ionic and covalent bonds [22–25]. Weakly shifted C1s peaks were also observed at 291.0 eV indicating small amounts of covalent C–F bonds. For plasma-fluorinated samples, similar C1s and F1s peaks were observed at 288.1 and 687.6 eV on average, respectively, with weakly shifted C1s peaks at Table 7 R (ID /IG) values calculated from Raman shifts of surface-fluorinated graphite samples Fluorination

R value

condition

NG7

NG25

NG40

Original graphite

0.083

0.082

0.080

Fluorinated by F2 at 150°C 200°C 250°C 300°C 350°C 400°C 450°C 500°C

0.18 – 0.23 – – 0.25 – 0.29

0.15 – 0.19 0.18 – 0.40 0.70 0.56

0.13 0.14 0.16 0.23 0.45 0.74 0.55 0.73

Plasma-fluorinated for 60 min 100 min 140 min 180 min

0.20 0.14 0.13 0.13

0.15 – – –

0.14 – – –

Applications of fluorinated carbon materials to primary and secondary lithium batteries

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Fig. 6. Transmission electron micrographs of surface-fluorinated graphite samples: (a) NG25 fluorinated by F2 at 250°C; (b) NG40 fluorinated by F2 at 300°C.

291.1 eV, indicating covalent C–F bonds. Thus the nature of C–F bonding is the same in graphite samples fluorinated by elemental fluorine and plasma treatment. 3.1.2. Charge/discharge characteristics of surface-fluorinated natural graphite

The profile of charge/discharge curves of surface-fluorinated samples is similar to that for non-fluorinated graphite as shown in Fig. 7, which shows the charge/discharge curves for original NG25 and that fluorinated at 250°C [23]. The charge/discharge characteristics were obtained at a current density of 60 mA/g between 0 and 3 V vs. Li/Li in 1 mol/dm3 LiClO4–ethylene carbonate (EC) / diethyl carbonate (DEC) (1:1 in volume) at 25°C [23,25]. The small plateau at 0.6–0.7 V in the first reduction curve is ascribed to the reduction of EC and subsequent formation of a thin surface film on graphite (solid electrolyte interface, SEI) [48]. The potential plateau indicating the decomposition of EC was no longer observed from the second cycle in any case. The surface-fluorinated NG25 exhibited higher capacity than original NG25 without any change in

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Fig. 7. Charge/discharge curves for (a) original NG25 and (b) NG25 fluorinated at 250 °C. —— : 1st cycle, ······· : 5th cycle, ----- : 10th cycle.

the profile of charge/discharge curves. NG7 and NG40 also showed the similar charge/discharge curves. The discharge capacities of surface-fluorinated NG7 samples were in the range 376–383 mAh/g at 60 mA/g, while original NG7 showed 360–363 mAh/g [22]. The discharge capacities of original and surface-fluorinated NG25 samples are shown in Fig. 8 as a function of cycle number [23]. NG25 had discharge capacities of 350–353 mAh/g smaller than those of NG7. As shown in Fig. 8, NG25 samples fluorinated between 150 and 400°C exhibited higher capacities than original NG25. Among them, the samples fluorinated between 200 and 300°C gave the highest capacities of 387–389 mAh/g, and those fluorinated at 150 and 350°C also provided high capacities of 382–384 mAh/g. The increments of the discharge capacities obtained at 10th cycle were ⬃10%, which is larger than ⬃5% for surface-fluorinated NG7. The results were similar in the case of NG40. The discharge capacities of original NG40 were 330–335 mAh/g smaller than those of NG7 and NG25. The NG40 samples fluorinated between 150 and 300°C showed high capacities of 374–377 mAh/g. The discharge capacities were

Applications of fluorinated carbon materials to primary and secondary lithium batteries

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390

Capacity (mAh/g)

380

370

360

350

340 0

2

4

6 Cycle number

8

10

Fig. 8. Discharge capacities of original NG25 and surface-fluorinated NG25 samples as a function of cycle number. (䉬) NG25, (䊏)150°C, (䉱) 200°C, () 250°C, (䉭) 300°C, (䊉) 350°C, (䊊) 400°C, (ⵧ) 450°C, (䉫) 500°C.

gradually decreased with increase in the fluorination temperature. The increments of the discharge capacities were ⬃13%. Since the theoretical capacity of graphite is 372 mAh/g, corresponding to LiC6, many fluorinated samples had higher capacities than the theoretical value: NG7 fluorinated between 150 and 500°C, NG25 between 150 and 400°C and NG40 fluorinated 150 and 300°C. The first coulombic efficiencies of NG7, NG25 and NG40 were 79.7, 85.6 and 85.1%, respectively. The samples fluorinated between 150 and 300°C showed the same first coulombic efficiencies as those of original graphites. It means that the irreversible capacities induced by surface fluorine are negligible for the samples fluorinated in this temperature range. The surface fluorine was reduced by ⬃40% in the course of electrode preparation [22]. This may be one of the reasons why the first coulombic efficiencies were not decreased for the samples fluorinated between 150 and 300°C. The first coulombic efficiencies were decreased with increasing fluorination temperature from 350 to 500°C, i.e. with increasing fluorine content. Plasma-fluorinated graphite sample had the same potential profile as that of original graphite. Figs. 9 and 10 show the discharge capacities of plasmafluorinated NG7 and NG25 as a function of cycle number [24]. The discharge capacities of plasma-fluorinated samples were highly dependent on the duration of plasma fluorination (Fig. 9). The 30-min-fluorinated NG7 showed 370 mAh/g higher than that of original NG7. The 60-min-fluorinated sample gave the highest

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Capacity (mAh/g)

390

380

370

360

350

3

1

5

7

9

Cycle number

Fig. 9. Discharge capacities of original NG7 and plasma-fluorinated NG7 samples as a function of cycle number. (䉫) NG7, (䉭) 30-min fluorinated, (䊉) 60-min fluorinated, 100-min fluorinated, (䊊)140-min fluorinated, *180-min fluorinated.

Capacity (mAh/g)

400

380

360

340

1

3

5 7 Cycle number

9

Fig. 10. Discharge capacities of original NG25 and plasma-fluorinated NG25 samples as a function of cycle number. (䉬)NG25, (䊉)60-min fluorinated at room temperature, (䉱) 60-min fluorinated at 90°C.

capacities, 382 mAh/g, which is larger than not only that of NG7 but also the theoretical capacity of graphite, 372 mAh/g. When the duration of plasma fluorination was further extended to 100–180 min, the discharge capacities decreased to the values comparable with that of NG7. The observed capacities are closely related to R values of Raman shifts given in Table 7. The plasma-fluorinated NG7 samples having the larger R values exhibited higher capacities. The decrease in the R value with increase in the duration of fluorination probably arose from the elimination of surface-disordered layers by extended plasma treatment. With the increase in the particle size, the capacities of NG25 and NG40 were reduced

Applications of fluorinated carbon materials to primary and secondary lithium batteries

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because the surface area decreases with increasing particle size. The effect of surface fluorination was more distinct as the particle size increases. The discharge capacities increased to 388 and 381 mAh/g for plasma-fluorinated NG25 and NG40, respectively. The increments in the discharge capacities by plasma fluorination were 5, 10 and 15% for NG7, NG25 and NG40, respectively. It was also found that first coulombic efficiencies were the same as those of original graphites, or slightly higher in some cases. Surface fluorination by elemental fluorine and plasma treatment both increased the discharge capacities of graphite samples to the higher value than theoretical capacity of graphite, 372 mAh/g. Surface fluorination increased the surface areas (Table 6) and mesopores with diameters of 1.5–2 and 2–3 nm. These surface structure changes would enhance the reaction kinetics and make possible the preservation of some excess lithium in surface mesopores. 3.2. Electrochemical behavior of surface-fluorinated petroleum coke 3.2.1. Composition and surface structure change of petroleum coke by fluorination

Starting materials were petroleum coke and those heat-treated at 1860, 2300 and 2800°C (abbreviated as PC, PC1860, PC2300 and PC2800C, respectively). The d spacings of (002) diffraction lines were 0.3450, 0.3385, 0.3366 and 0.3361 nm for PC, PC1860, PC2300 and PC2800, respectively. The d values were decreased with increasing heat-treatment (graphitization) temperature, indicating the increase in the crystallinity. However, all the d values were larger than that of natural graphite, 0.3354 nm. This means that the crystallinity of four petroleum coke samples is lower than that of natural graphite. Surface fluorination of petroleum coke was performed at 150, 200 and 300°C by 3  104 Pa elemental fluorine (purity: 99.4–99.7%) for 2 min, using a nickel reactor [26,27]. Table 8 shows composition obtained by elemental analysis and surface composition by XPS [26,27]. Fluorine contents were relatively higher in PC, in particular in PC fluorinated at 300°C. Smaller amounts of fluorine (0.3–0.6 at%) were detected for PC1860, PC2300 and PC2800 fluorinated at 300°C. Fluorine contents in the heat-treated petroleum cokes fluorinated at 150 and 200°C were in the detection limit ( 0.2 at%). Surface fluorine concentration had similar dependence on the heat-treatment temperature to that obtained by elemental analysis, decreasing from 50.2 at% to 5.2 at% with the increase in the graphitization temperature. Hydrogen (2.9 at% in Table 8) was detected in original PC by elemental analysis, which means that PC contained hydroxyl and carboxyl groups. PC and heat-treated petroleum cokes had the larger amounts of surface oxygen, 2.6–8.2 at% than natural graphite samples, 1.5–2 at%. The surface oxygen significantly decreased by heat – treatment from 8.2 to 2.6 at% and fluorination further reduced the surface oxygen to 0.9–2.5 at%. BET surface areas of petroleum cokes and surface-fluorinated samples are summarized in Table 9 [26,27]. The surface area of PC was large, decreasing to

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Table 8 Composition (at%) obtained by elemental analysis (upper column) and surface concentration (at%) by XPS (lower column) for surface-fluorinated petroleum cokes Fluorination temperature

PC

PC1860

C

F

O

C

F

Original 150°C 200°C 300°C

94.8 96.7 96.5 90.4

2.9 a 0.9 1.2 7.3

(2.3) (2.4) (2.3) (2.3)

99.4 99.8 99.8 99.2

– 0.0 0.0 0.6

Original 300°C

91.8 48.8

– 50.2

8.2 1.0

95.7 76.7

– 22.4

Petroleum coke PC2300 O

C

F

(0.6) 100.0 (0.2) 100.0 (0.2) 100.0 (0.2) 99.6 4.3 0.9

96.4 85.8

– 0.0 0.0 0.3 – 11.7

PC2800 O

C

F

O

(0.0) 100.0 – (0.0) (0.0) 99.8 0.0 (0.2) (0.0) 99.8 0.0 (0.2) (0.1) 99.5 0.4 (0.1) 3.6 2.5

97.4 93.2

– 5.2

2.6 1.6

0 for F within detection limit (<0.2 at%). aHydrogen contained in original petroleum coke. No hydrogen was detected in heat-treated petroleum cokes.

Table 9 BET surface areas (m2/g) of fluorinated petroleum cokes Fluorination

Petroleum coke

temperature (°C)

PC

PC1860

PC2300

PC2800

Original 150 200 300

6.35 7.35 8.06 25.2

3.31 3.03 3.31 3.20

2.33 2.21 2.27 2.25

2.43 2.85 2.73 2.81

half or less than half by heat-treatment. Surface area of PC was slightly increased by fluorination at 150 and 200°C, and largely increased at 300°C. The increase in the surface area seems to be proportional to the fluorine content in the sample. A low crystalline carbon with a large surface area is easily fluorinated, forming covalently bonded C–F layers, although the highly disordered parts may be lost by the fluorination as CF4 gas. It is known that the formation of graphite fluoride layers enlarges the surface area [49]. Surface areas of PC1860 and PC2300 were almost the same as or slightly smaller than those before fluorination. Disordered parts were probably eliminated as CF4 gas. On the other hand, fluorination behavior of PC2800 is rather similar to that of natural graphite powder because the surface areas of PC2800 were enlarged by 12–17%. Main reaction would be the formation of C–F covalent bonds accompanying the carbon–carbon bond breaking, which results in the formation of CF2/CF3 groups. Table 10 shows R values (ID/IG) obtained from intensity ratio of two Raman shifts (D and G bands) [26,27]. R value for non-fluorinated petroleum

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Table 10 R (= ID/IG) values obtained from peak intensity ratios of Raman shifts Fluorination

Petroleum coke

condition

PC

PC1860

PC2300 PC2800

Original sample

0.87

0.49

0.21

0.18

Fluorinated at 300°C

0.94

0.78

0.68

0.43

coke decreased from 0.87 to 0.18 with increase in the heat-treatment temperature, i.e. with the increase in the crystallinity of petroleum coke. It means that surface disorder of petroleum coke is significantly decreased by graphitization. It was recently found by transmission electron microscopic observation that heat treatment of a carbon material gives rise to the closure of edge surface by carbon– carbon bond formation. [50]. This coincides with the result obtained by Raman spectroscopy. It is clearly shown that surface fluorination enhanced the peak intensity of D-band at 1360 cm1. R values for fluorinated PC1860, PC2300 and PC2800 were much larger than those for non-fluorinated samples. The result shows that surface disorder of petroleum coke samples was enhanced by light fluorination accompanying carbon–carbon bond breaking and probably simultaneous formation of CF2/CF3 groups in the surface region. 3.2.2. Electrochemical behavior of surface-fluorinated petroleum coke

Cyclic voltammetry study was performed for petroleum coke samples in 1 mol/cm3 LiClO4 -EC/DEC at 25°C [26,27]. A large irreversibility was observed in the reduction and oxidation currents for PC; however, heat-treated petroleum cokes exhibited higher reversibility in reduction and oxidation currents. In the first reduction process, a reduction current peak indicating the reduction of EC and subsequent formation of a surface film (SEI) was observed at 0.55–0.60 V for heat-treated petroleum cokes though it was not clearly observed for PC. The reduction current peak indicating SEI formation was the smallest for PC1860 and increased with increasing heat-treatment temperature from 1860 to 2800°C. The formation of SEI occurs with intercalation of solvated lithium ion into surface region of a carbonaceous electrode and subsequent decomposition of the solvents [51,52]. The results suggest that the SEI formation is easier for a carbon with moderate graphitization degree. Surface fluorination of heat-treated petroleum cokes gives several positive effects to their electrochemical behavior. The oxidation peak potentials were slightly shifted to the lower values and the peak currents were increased by surface fluorination. This means that the reversibility of lithium ion intercalation and deintercalation into and from petroleum coke was improved by surface fluorinaton. No reduction current was observed at around 2.5 V for any of heat-treated petroleum cokes.

54

Tsuyoshi Nakajima

Potential curves for PC gradually decreased and increased with intercalation and deintercalation of lithium ion, respectively. However, the profile of potential curves for heat-treated petroleum cokes was rather similar to that for natural graphite as shown in Fig. 11 [26,27]. The discharge capacities of PC and PC1860 were in the range 246–258 mAh/g and those of PC2300 and PC2800 were in the range 285–306 mAh/g at a current density of 60 mA/g as given in

3.50

Potential (V vs Li/Li+)

3.00 1s t cyc le

2.50

10th cyc le

2.00 1.50 1.00 0.50 0.00

0

100

(a)

200 300 400 Capacity (mAh/g)

500

600

3.50

Potential (V vs Li/Li+)

3.00 1st cycle

2.50

10th cycle 2.00 1.50 1.00 0.50 0.00

(b)

0

100

200

300

400

500

600

Capacity (mAh/g)

Fig. 11. Charge/discharge curves for (a) PC2800 and (b) that fluorinated at 300°C, obtained at a current density of 60 mA/g.

Applications of fluorinated carbon materials to primary and secondary lithium batteries

55

Table 11 Discharge capacities (mAh/g) of petroleum coke samples at 1st and 10th cycles, obtained at a current density of 60 mA/g Fluorination temperature (°C) Original 150 200 300

Petroleum coke PC

PC1860

PC2300

PC2800

258–246 249–243 249–234 249–231

249–246 240–237 240–237 240–240

300–285 297–285 291–282 285–285

306–291 312–291 312–291 306–300

Table 11 [26]. PC1860, PC2300 and PC2800 showed 180–218 mAh/g at a higher current density of 150 mA/g. When crystallinity of a carbon material is low (PC and PC1860), amount of lithium ions accommodated between two graphene layers is relatively small because the carbon takes a turbostratic structure [53]. Heat treatment at higher temperatures than 2000°C improves the crystallinity of a carbon material and increases the discharge capacity. For this reason, PC2300 and PC2800 have higher discharge capacities than PC and PC1860. The potential profile was unchanged by surface fluorination in case of PC2300 and PC2800, except the first lithium intercalation curves below 1 V (Fig. 11). PC2300 and PC2800 also represented similar potential curves at a current density of 150 mA/g. It was found that surface fluorination did not give a positive effect to the discharge capacities of petroleum cokes except one case. PC2800 fluorinated at 300°C exhibited 233–210 mAh/g at 150 mA/g. The discharge capacities of 233–210 mAh/g were 10.9–16.7% higher than those of non-fluorinated sample. First coulombic efficiencies of non-fluorinated petroleum cokes varied depending on the heat-treatment temperature, i.e. the crystallinity of petroleum coke samples [26,27]. PC1860 exhibited the highest first coulombic efficiencies, 90.2 and 89.1%, at 60 and 150 mA/g, respectively, as given in Tables 12 and 13 [26,27]. However, first coulombic efficiencies were decreased with increasing heat-treatment temperature of petroleum coke, i.e. 71.9 and 70.0% for PC2300, and 65.4 and 63.6% for PC2800, at 60 and 150 mA/g, respectively, while natural graphite samples with high crystallinity had high first coulombic efficiencies of 80–85%. The first coulombic efficiencies obtained for petroleum cokes well coincides with the results obtained by cyclic voltammetry. The reduction peak of EC at 0.6 V was the smallest for PC1860, and increased for PC2300 and PC2800. This means that somewhat disordered surface is preferable for accommodating solvated lithium ions in the surface region and subsequent decomposition of the solvents to form SEI. It was recently reported by high resolution transmission electron microscopic study that edge plane of heat-treated carbon is closed by carbon–carbon bond formation [50], which would give some difficulty to

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Tsuyoshi Nakajima

Table 12 First coulombic efficiencies (%) for petroleum coke samples at a current density of 60 mA/g Fluorination

Petroleum coke

temperature (°C)

PC

PC1860

PC2300

PC2800

Original 150 200 300

72.3 68.6 68.6 47.4

90.2 87.9 87.9 83.3

71.9 72.8 72.4 84.1

65.4 61.5 60.5 83.6

Table 13 First coulombic efficiencies (%) for petroleum coke samples at a current density of 150 mA/g Fluorination temperature (°C) Original 150 200 300

Petroleum coke PC1860

PC2300

PC2800

89.1 88.0 88.0 81.8

70.0 75.0 76.5 83.3

63.6 54.7 59.6 79.5

electrochemical insertion of solvated lithium ion into surface region of a carbon material. This may be the main reason why the first coulombic efficiencies of PC2300 and PC2800 were lower than that for PC1860. On the other hand, natural graphite powder is prepared by pulverization of flake sample. Therefore its edge surface may be open though some oxygen atoms are bonded. An interesting result is that first coulombic efficiencies of PC2300 and PC2800 were significantly increased by surface fluorination as given in Tables 12 and 13. The first coulombic efficiencies were increased by 12.2–18.2% when they were fluorinated at 300°C, and those of PC2300 fluorinated at 150°C and 200°C were increased by 0.5–6.5%. This is also shown in Fig. 11, in which the potential at first cycle quickly lowered below 1 V for fluorinated PC2800. Several factors should be considered for the increase in first coulombic efficiencies of PC2300 and PC2800 by fluorination. They are surface structure and influence of surface chemical species such as fluorine and oxygen. The change in the surface areas was small before and after fluorination for PC2300 and PC2800 as given in Table 9. This is an advantage for heat-treated petroleum cokes because electrochemical decomposition of EC does not increase. Another advantage given by fluorination would be the increase in surface disorder of PC2300 and PC2800 by carbon–carbon bond breaking. Edge plane closed by

Applications of fluorinated carbon materials to primary and secondary lithium batteries

57

carbon–carbon bond formation due to heat treatment would be opened by fluorination reaction with carbon–carbon bond breaking and formation of CF2/CF3 groups. The opening of edge plane can make the accommodation of solvated lithium ion at edge surface to facilitate the SEI formation easy. Adsorption of a polar solvent molecule such as EC may have a difficulty on surface-fluorinated petroleum cokes due to the hydrophobic nature of surface CF2 and CF3 groups with covalent bonds. Fluorine atoms bonded to graphite surface is finally removed as LiF by electrochemical reduction, because the amount of fluorine detected by XPS analysis was only 0.2 at% after charge/discharge cycling [22]. LiF thus formed on petroleum coke would facilitate the SEI formation as one of the constituents of SEI. The other effect of surface fluorination would be the decrease in the amounts of surface oxygen as given in Table 8. The decrease in the surface oxygen by fluorination was 1.1–1.0 at% for PC2300 and PC2800, respectively. Since surface oxygen interacting with lithium ion is one of the factors to increase the irreversible capacity, the replacement of surface oxygen to fluorine is found to give a positive effect to the first coulombic efficiency. ACKNOWLEDGEMENTS The present study is partly supported by a grant from the Frontier Research Project “Materials for the 21st century – Materials Development for Environment, Energy and Information” (for 2002–2006 fiscal years) from Ministry of Education, Culture, Sports, Science and Technology. Fluorine gas used in this study was supplied by, the courtesy of Daikin Industries, Ltd. The author gratefully acknowledges them. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13]

N. Watanabe, T. Nakajima, and H. Touhara, Graphite Fluorides, Elsevier, Amsterdam, 1988. T. Nakajima and N. Watanabe, Graphite Fluorides and Carbon–Fluorine Compounds, CRC Press, Boca Raton, 1991. T. Nakajima (Ed.), Fluorine–Carbon and Fluoride–Carbon Materials – Chemistry, Physics and Applications, Marcel Dekker, New York, 1995. T. Nakajima, Advanced Inorganic Fluorides, T. Nakajima, B. Z emva, and A. Tressaud (Eds.), Elsevier, Lausanne, 2000, Chap. 15. T. Nakajima, M. Kawaguchi, and N. Watanabe, Electrochim. Acta, 27 (1982) 1535. R. Hagiwara, M. Lerner, N. Bartlett, and T. Nakajima, J. Electrochem. Soc., 135 (1988) 2393. A. Hamwi, M. Daoud, and J.C. Cousseins, Synth. Metals, 30 (1989) 23. A. Hamwi, J. Phys. Chem. Solids, 57 (1996) 677. P. Hany, R. Yazami, and A. Hamwi, J. Power Sources, 68 (1997) 708. R. Yazami, P. Hany, P. Masset, and A. Hamwi, Mol. Cryst. Liq. Cryst., 310 (1998) 397. T. Mallouk and N. Bartlett, J. Chem. Soc. Chem. Commun., (1983) 103. R. Hagiwara, M. Lerner, and N. Bartlett, J. Chem. Soc. Chem. Commun., (1989) 573. T. Mallouk, B.L. Hawkins, M.P. Conard, K. Zilm, G.E. Maciel, and N. Bartlett, Philos. Trans. R. Soc. Lond. Ser. A, 314 (1985) 179.

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Tsuyoshi Nakajima Y. Sato, T. Kume, R. Hagiwara, and Y. Ito, Carbon, 41 (2003) 351. Y. Sato, S. Shiraishi, Z. Mazej, R. Hagiwara, and Y. Ito, Carbon, 41 (2003) 1971. T. Nakajima, M. Koh, V. Gupta, B. Z emva, and K. Lutar, Electrochim. Acta, 45 (2000) 1655. V. Gupta, T. Nakajima, Y. Ohzawa, and B. Z emva, Mol. Cryst. Liq. Cryst., 386 (2002) 25. T. Nakajima, V. Gupta, Y. Ohzawa, H. Groult, Z. Mazej, and B. Z emva, J. Power Sources, J. Power Sources 137 (2004) 80. K. Kanamura, Advanced Inorganic Fluorides, Synthesis, Characterization and Applications, T. Nakajima, B. Z emva, and A. Tressaud (Eds.), Elsevier, Amsterdam, 2000, Chap. 16. F. Kita, H. Sakata, S. Sinomoto, A. Kawakami, H. Kamizori, T. Sonoda, H. Nagashima, J. Nie, N.V. Pavlenko, and Y.L. Yagupolskii, J. Power Sources, 90 (2000) 27. M. Takashima, S. Yonezawa, and M. Ozawa, Mol. Cryst. Liq. Cryst., 388 (2002) 153. T. Nakajima, M. Koh, R.N. Singh, and M. Shimada, Electrochim. Acta, 44 (1999) 2879. V. Gupta, T. Nakajima, Y. Ohzawa, and H. Iwata, J. Fluorine Chem., 112 (2001) 233. T. Nakajima, V. Gupta, Y. Ohzawa, M. Koh, R.N. Singh, A. Tressaud, and E. Durand, J. Power Sources, 104 (2002) 108. T. Nakajima, V. Gupta, Y. Ohzawa, H. Iwata, A. Tressaud, and E. Durand, J. Fluorine Chem., 114 (2002) 209. T. Nakajima, J. Li, K. Naga, K. Yoneshima, T. Nakai, and Y. Ohzawa, J. Power Sources, 133 (2004) 243. J. Li, K. Naga, Y. Ohzawa, T. Nakajima, A.I. Shames, and A.M. Panich, J. Fluorine Chem., in press. T. Nakajima, M. Mori, V. Gupta, Y. Ohzawa, and H. Iwata, Solid State Sci., 4 (2002) 1385. K. Amine, T. Nakajima, and M. Motoyama, Carbon, 32 (1994) 1067. J.P. Lemmon and M.M. Lerner, Carbon, 31 (1993) 437. V.K. Mahajan, R.B. Badachhape, and J.L. Margrave, Inorg. Nucl. Chem. Lett., 10 (1974) 1103. Y. Kita, N. Watanabe, and Y. Fujii, J. Am. Chem. Soc., 101 (1979) 3832. F. Tuinstra and J.L. Koenig, J. Chem. Phys., 53 (1970) 1126. D.S. Knight and W.B. White, J. Mater. Res., 4 (1989) 385. Y. Sato, R. Hagiwara, and Y. Ito, J. Fluorine Chem., 110 (2001) 31. N. Watanabe, R. Hagiwara, T. Nakajima, H. Touhara, and K. Ueno, Electrochim. Acta, 27 (1982) 1615. H. Touhara, H. Fujimoto, N. Watanabe, and A. Tressaud, Solid State Ionics, 14 (1984) 163. N. Watanabe, T. Nakajima, and R. Hagiwara, J. Power Sources, 20 (1987) 87. J.S. Xue and J.R. Dahn, J. Electrochem. Soc., 142 (1995) 3668. E. Peled, C. Menachem, D. Bar-Tow, and A. Melman, J. Electrochem. Soc., 143 (1996) L4. M. Hara, A. Satoh, N. Tamaki, and T. Ohsaki, Tanso, 165 (1994) 261. R. Takagi, T. Okubo, K. Sekine, and T. Takamura, Denki Kagaku, 65 (1997) 333. M. Yoshino, H. Wang, K. Fukuda, Y. Hara, and Y. Adachi, J. Electrochem. Soc., 147 (2000) 1245. H. Wang, M. Yoshino, T. Abe, and Z. Ogumi, J. Electrochem. Soc., 149 (2002) A499. M. Yoshino, H. Wang, K. Fukuda, T. Umeno, N. Dimov, and Z. Ogumi, J. Electrochem. Soc., 149 (2002) A1598. Y.-S. Han and J.-Y. Lee, Electrochim. Acta, 48 (2003) 1073. Y. Ohzawa, M. Mitani, T. Suzuki, V. Gupta, and T. Nakajima, J. Power Sources, 122 (2003) 153. R. Fong, U. von Sacken, and J.R. Dahn, J. Electrochem. Soc., 137 (1990) 2009. T. Nakajima, A. Mabuchi, R. Hagiwara, N. Watanabe, and F. Nakamura, J. Electrochem. Soc., 135 (1988) 273.

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[50] K. Moriguchi, S. Munetoh, M. Abe, M. Yonemura, and K. Kamei, A. Shintani, A. Omaru, and M. Nagamine, J. Appl. Phys., 88 (2000) 6369. [51] M. Inaba, S.-K. Jeong, T. Abe, and Z. Ogumi, Battery Tech., 13 (2001) 31. [52] M. Inaba, Y. Kawatate, A. Funabiki, S.-K. Jeong, T. Abe, and Z. Ogumi, Electrochim. Acta, 45 (1999) 99. [53] H. Fujimoto, A. Mabuchi, K. Tokumitsu, and T. Kasuh, Carbon, 38 (2000) 871.

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 3

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration Yoshimi Ohzawa Department of Applied Chemistry, Aichi Institute of Technology, Yakusa, Toyota 470-0392, Japan 1. INTRODUCTION Over the last decade, chemical vapor deposition (CVD) has been utilized for the purpose of synthesizing new types of carbon materials required for the preparation of the anode of lithium-ion secondary batteries (LIBs), involving carbonaceous thin films [1,2], carbon soot with a disordered structure [3] and carbon materials containing foreign atoms such as nitrogen and silicon [4–6]. Pyrolytic carbon (pyrocarbon) grown by CVD process shows a capacity of 250–700 mAhg1, depending on the macro-texture, microstructure, crystallinity and existence of foreign atoms. These properties of pyrocarbon can be varied by varying the CVD conditions (e.g. temperature, pressure and hydrocarbon species) [3,7] and the source species for the dopants [4,6]. Recently, pyrocarbon coating by CVD was applied to graphite-based anodes of LIBs to improve anode performance, especially in propylene carbonate (PC)-based solvents [8–12]. In this process, a low-crystalline pyrocarbon shell covered the surface of a high-crystalline graphite core, which successfully prevented the decomposition of the PC solvent. Thus, CVD is effective in synthesizing new carbons or modifying the core carbon; however, it is difficult for conventional CVD to achieve the homogeneous deposition of pyrocarbon with relatively high crystallinity. The source gases in the conventional CVD process are made to flow through a reaction vessel under isothermal and isobaric conditions. In this continuous gas-flow CVD, the source gases are preheated before reaching the surface of substrate. Preheating of source gases leads to the formation of active precursors followed by nucleation, which often results in the formation of high-molecular-weight compounds such as tar or carbon soot in the gas phase. Co-deposition of tar or soot affects the crystallinity of the pyrocarbon film. The low-crystalline carbon

62

Yoshimi Ohzawa

often shows a large irreversible capacity during the first charge – discharge cycle. In addition, thick films are easily formed on the external surface of the porous substrates or particle-packed beds in the continuous gas-flow CVD. Film formation on the external surface prevents the source gas from penetrating into the substrate and creating a uniform coating throughout the substrate. In order to achieve a uniform coating on the conventional CVD, it is necessary to repeat the CVD treatment and remove the films by polishing the external surface or softly grinding the particles several times. Chemical vapor infiltration (CVI) has received attention as a preparation process for fiber- or particle-reinforced composites. In this process, the source gases are flowed through the fibrous or particulate preforms at a high temperature, during which specific material is deposited as a matrix between the fibers or particles [13,14]. In the past few decades, three main methods have been developed: isothermal and isobaric CVI (ICVI) [15,16], forced CVI (FCVI) [17,18] and pressure-pulsed CVI (PCVI) [19–22]. In the PCVI process, the following steps are repeated: evacuation of the reaction vessel, instantaneous introduction of the source gas, and holding to allow deposition. Compared with the conventional CVD process, the PCVI process allows the homogeneous infiltration of matrix through the thickness of the preforms under suitable conditions because of the rapid penetration of the source gas throughout the preform without preheating [23,24]. In addition, the instantaneous introduction of the source gas and the evacuation of the reacting gas are repeated at relatively short intervals in the PCVI process, therefore, nucleation in the gas phase is restrained, and crystal growth is promoted in the evacuation step. These are effective in increasing the crystallinity of pyrocarbon, resulting in high coulombic efficiency during the first cycle [25,26]. This chapter summarizes the recent results of the preparation of new types of carbon-based anodes and surface modifications by coating with pyrocarbon on low-crystalline carbon using PCVI. As mentioned in Chapter 2, surface fluorination effectively modifies the surface structures of graphite particles used as anode material in LIBs [27–31]. Surface fluorination of pyrocarbon film is also described in this chapter. 2. PREPARATION OF NEW CARBON-BASED ANODE AND ITS MACROSCALED STRUCTURE 2.1. Formation of three-dimensionally continuous conduction network

The anode properties of LIB, such as reversible capacity, irreversible capacity, high rate property and cyclability, strongly depend not only on the nanoscaled structure of carbon as anode material but also on the macroscaled structure of electrodes, involving the nature, the structure and the content of organic binders, current collectors and conductive fillers added to assist the construction of

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration

63

conduction network. In general, fine graphite or carbon black powders are utilized as the conductive fillers; however, the use of the fine particles leads to an increase in the active surface area of electrode, resulting in reactions with electrolytes. These reactions cause the irreversible loss and decrease in the electrode capacity during charge – discharge cycling. Instead of the fine carbon particles, the use of fibrous conductive additives was examined. Micron-sized metal fibers with low active area and high conductivity are effective for enhancing electrode capacity, especially at a high rate with very low volumetric content [32]. Using the vapor-grown carbon fibers, the capacity and the cyclic life of electrodes have been improved due to their excellent conductivity and high surface-to-volume ratio [33]. Thus, in order to improve anode performance, it is necessary to optimize not only the nanoscaled structure of carbon but also the macroscaled structure of electrodes. As mentioned in Section 1, the CVI process was developed for matrix filling into fiber preforms to prepare the fiber-reinforced ceramics [13,14]. Recently, the PCVI method has been applied to synthesize porous ceramics, for example, the porous SiC foams prepared by partial densification, with SiC matrix, into the biologically derived porous preforms such as carbonized wood, cotton and paper [34,35]. By a similar process, highly porous electroconductive substrates can be obtained to partially infiltrate electroconductive materials such as TiN and TiC instead of SiC matrix [25,36]. Utilizing these conductive porous substrates as the conductive fillers or current collectors for LIBs, the negative electrodes containing the three-dimensionally continuous current paths are prepared to infiltrate pyrocarbon from the gas phase [25,26,37]. Because the pyrocarbon films are directly deposited on the surface of the current collector in the present process, it is assumed that the contacting resistance between pyrocarbon and current collector is low, even if no organic binder and additional conductive filler are used. 2.2. Highly porous electroconductive substrates for current collector

Highly porous electroconductive substrates are synthesized by partial infiltration, with TiN, into the porous preforms such as carbonized wood, cotton wool and paper, which are prepared by carbonization at 1000°C in Ar for 4 h [36]. Fig. 1 shows the main part of the apparatus for PCVI. The source gas mixture was allowed to flow into a reservoir. It was instantaneously introduced (within 0.1 s) into the reaction vessel up to a pressure of 0.1 MPa, and the pressure was held under the same conditions to allow matrix deposition for the desired time (holding time). Then, the gas was evacuated to below 0.7 kPa within 1.5 s. This cycle of sequential steps was defined as 1 pulse, and it was repeated for the desired number of times. Fig. 2 shows the SEM images of the TiN-coated paper (photographs (a) and (b)) and wood substrates (photographs (c) and (d)) prepared by 10,000 pulses of PCVI at 850°C with a holding time of 1.5 s from the gas system of TiCl4

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Yoshimi Ohzawa

7

8 4 9 1 2 3

6

5

Fig. 1. Main part of apparatus for pressure-pulsed chemical vapor infiltration. 1, source gas; 2, reservoir; 3, electromagnetic valve; 4, pressure gauge; 5, vacuum tank; 6, vacuum pump; 7, furnace; 8, substrates; 9, thermocouple.

Fig. 2. SEM images of electroconductive porous substrates: (a) and (b)TiN-coated paper, (c) and (d) TiN-coated wood. Substrates were prepared by 10,000 pulses of PCVI for TiN.

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration

65

(1%) – N2 (10%) – H2. From low-magnification photograph (a), it can be observed that the fibers in the TiN-coated paper substrate have a relatively random orientation, and that the TiN-coated fibers connect with each other. On the other hand, it can be observed from photograph (c) that the carbonized wood substrate has a honeycomb-shaped cellular structure and pores with the cross-section of a rectangle, penetrating through the substrate. Photographs (b) and (d) show high-magnification images of the cross-section of the substrates (a) and (c), respectively. It can be found that TiN films with a thickness of 0.5–1 μm deposit around the carbonized fibers of about 6 μm in diameter (photograph (c)) or on the cell walls of the carbonized wood (photograph (d)). It appears that TiN films adhere tightly to the carbonized fibers or the cell walls. Table 1 shows the properties of TiN-coated porous substrates as the current collector for the rechargeable batteries, along with those of the foil-type current collector for the LIB and the metal (Ni) foam current collector for the Ni – Cd battery. In the commercial LIB, the volume fraction of the active materials layer per unit volume of the electrode was estimated to be about 75%, based on a simple geometric calculation [38]. The porosity (i.e. free space for filling with active materials) of each substrate is higher than 75%. The apparent resistivity of TiNcoated substrates is in the range 105–106 Ω m, which is low enough for use as current collector. The average pore sizes are below 40 μm, which is lower than the thickness of the active materials layer in the commercial LIB and the pore size of the metal (Ni) foam current collector for the Ni – Cd battery. Large cavities are unsuitable because of the low conductivity of the active materials and the electrolytes for the lithium-ion rechargeable battery. The geometric surface area per unit volume of each TiN-coated porous substrate shows a higher value than that of the conventional current collector, which could lead to the reduction of contacting resistance between active materials and current collectors. TiN is inert Table 1 Specific properties of TiN-coated substrates Substrate

Porosity (%)

Resistivity (Ωm)

Average pore size (μm)

Geometric surface area (m2 m3)

TiN-coated substrates TiN-coated papera TiN-coated wooda

84–88 80–85

9  106 7  106

18 15

8  104 18  104

Foil-type current collector

(75)b

106

–

1  104

Metal-foam current collector

92–96

106

200

3  104

a

Number of pulses in PCVI treatment for TiN is 10,000. b Volume fraction of active material layer per unit electrode volume.

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to the electrochemical reaction with lithium ion under the galvanostatic condition at an apparent current density of 0.3 mAcm2 in a potential range from 0 to 3 V vs. Li/Li. 2.3. Macroscaled structure and anode performance

Utilizing the TiN-coated porous substrates as the conductive fillers and current collectors for LIBs, the new type of negative electrode containing the threedimensionally continuous current paths is prepared to infiltrate pyrocarbon from the gas phase of C3H8 (30%) – H2 at 950°C with a holding time of 1.0 s [25,37]. For the electrode obtained after 40,000 pulses of PCVI, the volume fraction of the current collector (TiN-coated substrate) is 20%; active material (pyrocarbon) occupies 59%, and pores 21%. Mass fraction of pyrocarbon per unit volume of the electrode reaches 0.9 g cm3, which is in a range similar to that of the electrode for the commercial lithium-ion battery [39]. Fig. 3 shows the SEM images of the pyrocarbon-based electrodes obtained from the TiN-coated paper (photographs (a) and (b)) and wood substrates (photographs (c) and (d)). From lowmagnification micrograph (a), it can be observed that the fibers have a relatively random orientation and connect with each other. From image (b), it can be seen that dense films of pyrocarbon with a thickness of ∼3 μm are formed on the TiN films. It appears that the pyrocarbon films adhere tightly to the TiN films. In the

Fig. 3. SEM images of pyrocarbon-based electrodes. (a) and (b) TiN-coated paper, (c) and (d) TiN-coated wood. The number of pulses in PCVI for pyrolytic carbon. (a) and (b) 32,500, (c) and (d) 40,000.

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration

67

case of the TiN-coated wood substrate (photograph (c)), pyrocarbon is deposited onto the pores between the carbonized cell walls. Micrograph (d) shows the image of the cross-section at half thickness of the thin plate sample obtained from the TiN-coated wood substrate. The thickness of the pyrocarbon at half depth from the sample surface is close to that at the position near the surface (photograph (c)). It is considered that the pyrocarbon films are almost uniform in thickness throughout the sample. These SEM observations indicate that threedimensional current paths are formed in the active material layers of the negative electrodes. It is expected that the contacting resistance between active material (pyrocarbon) and current collector (TiN-coated substrate) will be relatively low even if no organic binders and additional conductive fillers are used. Fig. 4 shows the charge curves of pyrocarbon in the sample obtained from TiN-coated paper under different current densities, where galvanostatic charge – discharge cycling is made at 25°C, using a three-electrode cell with metallic lithium as counter and reference electrodes in 1 mol L1 LiClO4 EC/DEC (1:1) solution. The capacity is calculated using only the weight of pyrocarbon in sample, considering the significantly low capacity below 10 mAhg1 of the TiNcoated paper substrate. In the charge curve at current density of 25 mAg1, the plateau is observed at the potential below 0.2 V, after which the potential gradually 3.5

Potential (V vs Li/Li+)

3.0 2.5 2.0 1.5 1.0 0.5 0.0

0

100

200

300

400

500

Capacity (mA h g-1)

Fig. 4. Charge curves of pyrocarbon deposited on TiN-coated paper substrate at several current density: (●) 0.2 mA cm2 (25 mA g1), (▲) 0.8 mA cm2 (100 mA g1), (■) 3.2 mA cm2 (400 mA g1) and (◆) 8 mA cm2 (1000 mA g1). The capacity was calculated using the mass of pyrolytic carbon in the sample. The current density at discharge was kept at 0.2 mA cm2. The number of pulses in PCVI for pyrolytic carbon is 7000.

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increases. The charge capacity at 25 mAg1 is 460 mAhg1 per pyrocarbon weight in the sample. This value is higher than the theoretical capacity of graphite (372 mAhg1), corresponding to the stage 1 lithium-intercalated graphite, LiC6. This behavior is similar to that observed in some non-graphitizing carbon having a disordered structure, and can be related to the nanoscaled structure of pyrocarbon; the details are given in the next section. The capacity at 1000 mAg1 is 80% of that at 25 mAg1, and 92% of the initial charge capacity is maintained until the 60th cycle [36]. It is supposed that the relatively high capacity at high charge rate and the good cyclability results from the formation of three-dimensionally continuous current paths in the active material layers and the tight adhesion between pyrocarbon and TiN, which is effective in reducing internal resistance and restraining the failure of electroconduction during charge – discharge cycling. 3. EFFECT OF SUBSTRATE ON NANOSCALED STRUCTURE OF PYROCARBON AND ITS ELECTROCHEMICAL CHARACTERISTICS As mentioned in the previous section, anode performance is affected by the macroscale structure of the electrode. In addition, the electrochemical properties of carbon strongly depend on nanoscaled structure, crystallinity, existence of foreign atoms and so on. For pyrocarbon grown by the CVD process, these properties can be varied by varying the CVD conditions (e.g. temperature, pressure and hydrocarbon species) [3,7] and the source species for the dopants [4,6]. The type of substrates on which pyrocarbon deposits can also affect the structure of pyrocarbon [25,40] because the nucleation and crystal growth of carbon often proceed on the surface of the substrate in the CVD/CVI process. 3.1. Structure of pyrocarbon deposited on carbon or TiN

In the case of the PCVI process for pyrocarbon, its crystallinity and surface structure are also affected by the kind of substrate used [25,37]. For instance, the pyrocarbon film deposited on a naked carbonaceous substrate has relatively high crystallinity and laminar microstructure, whereas the pyrocarbon on the TiNcoated substrate is disordered. Fig. 5 shows the SEM images of pyrocarbon film (A) deposited directly on the carbonized wood substrate (photographs (a)) and pyrocarbon (B) on the TiN-coated wood substrate (photographs (b)), where the preparation procedures of these substrates and the PCVI conditions for pyrocarbon are the same as those mentioned in section 2. It is found that pyrocarbon has a laminar texture, in which the basal planes of carbon crystallites are oriented parallel to the cell-wall surface of the carbonized wood. The laminar structure appears remarkably in pyrocarbon film (A) deposited directly on the carbonized wood as shown in photograph (a). This laminar structure is effective in reducing irreversible reactions such as the decomposition of the electrolytes, because the basal planes of carbon with low reactivity are mainly exposed to the electrolyte solution.

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration

69

Fig. 5. SEM images of pyrocarbon film deposited directly on the carbonized wood substrate (a) and pyrocarbon on the TiN-coated wood substrate (b). The number of pulses in PCVI for pyrolytic carbon is 40,000.

(200)

: Carbon : TiN

Intensity ( a. u. )

(111)

Sample (B) (002)

(10)

Sample (A)

Carbonized wood

10

20

30 40 Cu Kα 2θ (deg.)

50

60

Fig. 6. X-ray diffraction patterns from the external surface of the carbonized wood/pyrocarbon sample (A), the carbonized wood/TiN/pyrocarbon sample (B), and the original carbonized wood substrate (C). The number of pulses in PCVI for pyrolytic carbon is 40,000.

Fig. 6 shows the X-ray diffraction (XRD) patterns of the carbonized wood/pyrocarbon sample (A), the carbonized wood/TiN/pyrocarbon sample (B), and the original carbonized wood substrate (C). The filling ratios of pyrocarbon in samples (A) and (B) are 67 and 68%, respectively. For the original substrate (C), it can be observed that (002) diffraction peak at 2θ of 22.7° (d002  0.399 nm) is very weak and broad, reflecting the low crystallinity of carbon obtained from the carbonization of cellulose. However, for carbonized wood/pyrocarbon sample (A), a strong (002) peak appears at a higher angle of 25.4° (d002  0.359 nm). These results indicate that the crystallinity of pyrocarbon deposited on the

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Yoshimi Ohzawa

carbonized wood is much higher than that of the substrate carbon. In addition, the intensity of the (10) diffraction peak is relatively strong compared with (002) diffraction peak. The strong (10) peak results from the orientation of the pyrocarbon films perpendicular to the external surface of the plate-like sample as shown in Fig. 5. On the other hand, the intensity of (10) diffraction line for the carbonized wood/TiN/pyrocarbon sample (B) is weaker than that for sample (A) although the strength and peak position of (002) line for sample (B) is close to that for sample (A). As shown in Fig. 5 (SEM image (a)), pyrocarbon in sample (A) has a laminar texture perpendicular to the external surface of the plate-like sample. This orientation results in the strong intensity of the (10) diffraction line of sample (A). On the other hand, it is suggested that the orientation of the pyrocarbon deposited on TiN in sample (B) is disturbed, i.e. the degree of structural disordering in the pyrocarbon of sample (B) is high compared with the pyrocarbon of sample (A). It is also revealed from the Raman spectra that the degree of structural disordering in pyrocarbon deposited on TiN-coated substrate (sample (B)) is high in comparison with pyrocarbon directly deposited on carbonized substrate (sample (A)) [25]. In the case of the continuous gas-flowing CVD, the source gases are preheated before reaching the surface of the substrate, leading to the formation of active precursors followed by nucleation, which often causes the formation of high-molecular-weight compounds such as tar or carbon soot. Co-deposition of tar or soot would affect the crystallinity of pyrocarbon film. In the PCVI process, instantaneous introduction of the source gas and evacuation of the reacting gas are repeated in relatively short intervals. As the source gas rapidly penetrates throughout the preform without preheating, nucleation in gas phase is restrained, and crystal growth is promoted in evacuation step. These may result in the relatively high crystallinity of pyrocarbon in sample (A). However, the degree of structural disordering in pyrocarbon deposited on TiNcoated wood (sample (B)) is high in comparison with pyrocarbon directly deposited on carbonized wood (sample (A)). It is supposed that the formation of C–C bonds between pyrocarbon and substrate carbon for sample (A) is easier than for sample (B), which may affect the crystal growth rate on the surface of the substrate to form laminar structure. In addition, Ti atoms easily react with carbon to form TiC, which may disturb the orientation of crystallites in pyrocarbon deposited on TiN. The surface area data obtained by BET method are given in Table 2 for the original carbonaceous substrates, the carbonized wood or paper/pyrocarbon samples (A), and the carbonized wood or paper/TiN/pyrocarbon samples (B). The surface area of each original substrate significantly decreases owing to the infiltration of pyrocarbon. This result indicates that the dense pyrocarbon film covers most of the surface of porous carbon substrate. The surface area of sample (B-1) or (B-2) is higher than that of sample (A-1) or (A-2), respectively. This result indicates that pyrocarbon deposited on TiN in sample (B) is porous and nanometer-sized as

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration

71

Table 2 BET surface area data of original carbon substrate prepared from paper and wood, pyrocarbons deposited directly on carbon substrates (A-1, A-2) and pyrocarbons deposited on TiN-coated carbon substrates (B-1, B-2) Sample

BET surface area (m2 g1)

Original carbonized paper

170–210

Original carbonized wood

80–120

(A-1) Carbonized paper/pyrolytic carbon

0.81

(A-2) Carbonized wood/pyrolytic carbon

0.58

(B-1) Carbonized paper/TiN/pyrolytic carbon

33

(B-2) Carbonized wood/TiN/pyrolytic carbon

42

Note: Number of pulses in PCVI treatment for pyrocarbon is 1000.

compared with pyrocarbon deposited directly on carbon substrate in sample (A). From pore volume distribution analysis of both samples, it is found that sample (B) has mesopores with a diameter of 1.5 to 7 nm, and the volume fraction of mesopores below 3nm in sample (B) is extremely large compared with that in sample (A) [25]. It is expected that these mesopores below 3 nm may accommodate some amount of excess lithium in addition to lithium insertion into the interlayer of graphite-like stacking. 3.2. Charge – discharge behavior of pyrocarbon deposited on carbon or TiN

Fig. 7 shows the first charge – discharge curves of the pyrocarbon anodes deposited on carbonized paper (a) and TiN-coated paper (b). In the case of the carbonized paper/pyrocarbon sample (A), it can be seen that the potential gradually decreases and increases with lithium intercalation and de-intercalation, respectively. The charge – discharge profiles are similar to those usually observed for low-crystalline carbon materials having a laminar microstructure (i.e. soft carbon). From the SEM image and XRD pattern of sample (A), the structure of pyrocarbon deposited directly on carbon substrates is considered to be similar to that of soft carbons. It is supposed that a sloping profile over a range of potentials reflects the crystallinity of pyrocarbon in sample (A), although sample (A) contains 50 wt% of the carbonized fibers having a disordered structure. As shown in Table 3, the charge capacities (lithium de-intercalation process) per total weight of sample are 298 and 290 mAhg1 for the carbonized paper/pyrocarbon (A-1) and the carbonized wood/pyrocarbon (A-2), respectively. The substrate carbon shows a capacity of 285–320 mAhg1; therefore, the capacity of pyrocarbon in sample (A) seems to be in a similar range, assuming additivity in capacity.

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Yoshimi Ohzawa 3.5 (a)

Potential (V vs Li/Li+)

3.0 2.5 2.0 1.5 1.0 0.5 0.0

0

100

200

300

400

3.5 (b) Potential (V vs Li/Li+)

3.0 2.5 2.0 1.5 1.0 0.5 0.0 0

100

200

300

400

500

600

700

Capacity (mA h g-1)

Fig. 7. First charge – discharge curves of carbonized wood/pyrocarbon sample (a) and carbonized wood/TiN/pyrocarbon sample (b) obtained after 7000 pulses of PCVI treatment. The current density was 0.2 mA cm2. The capacity of samples (a) and (b) were calculated using the total mass of the sample and the mass of pyrolytic carbon in the sample, respectively.

One noticeable feature of the pyrocarbon sample (A) deposited on the carbonized substrate is its high first coulombic efficiency, for example, 85% for the carbonized paper/pyrocarbon and 87% for the carbonized wood/pyrocarbon, which is higher than those of the original carbonaceous substrates (63–68%). In sample (A), the dense pyrocarbon film with higher crystallinity covers most of the surface of porous carbon substrates with a disordered structure. In addition, the pyrocarbon film has a laminar texture oriented parallel to the surface of the substrate. It is suggested that the reactive sites (e.g. the crystal edges and the functional groups) at the surface of the carbonized substrate are covered, and the basal planes with low reactivity are exposed mainly to the electrolyte solution. These

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration

73

Table 3 Capacity and first coulombic efficiency data of original carbon substrate prepared from paper and wood, pyrocarbons deposited directly on carbon substrates (A-1, A-2) and pyrocarbons deposited on TiN-coated carbon substrates (B-1, B-2) Capacitya (mAh g1)

Capacityb (mAh g1)

Coulombic efficiency (%)

Original carbonized paper

–

320

68

Original carbonized wood

–

302

63

(A-1) Carbonized paper/pyrolytic carbon

–

298

85

(A-2) Carbonized wood/pyrolytic carbon

–

290

87

(B-1) Carbonized paper/TiN/pyrolytic carbon

460

120

73

(B-2) Carbonized wood/TiN/pyrolytic carbon

442

115

68

Sample

a

Note: Number of pulses in PCVI treatment for pyrocarbon, 5000; current density, 0.2 mA cm2. Capacity per mass of pyrocarbon. b Capacity per total mass of sample.

surface structures reduce irreversible reactions such as the decomposition of electrolytes, resulting in high coulombic efficiency at the first cycle. On the other hand, the first charge – discharge curves of the carbonized paper/TiN/pyrocarbon sample (B) are shown in Fig. 7(b), where the capacity is calculated using only the weight of pyrocarbon in sample (B), considering the significantly low capacity below 10 mAh g1 of the TiN-coated substrate. On the discharge reaction (i.e. lithium intercalation), the potential gradually decreases up to 0.1 V, below which the long plateau appears. In the case of the charging process, the plateau is also observed at a potential below 0.2 V. This behavior is similar to that observed in non-graphitizing carbon with a disordered structure. As shown in Fig. 7(a), no low-voltage plateau is observed on the discharge or charge curve of pyrocarbon deposited directly on the carbonized paper (sample (A)). It was also reported that the electrochemical properties of pyrocarbon films by CVD were similar to those of the soft carbons [1,4]. From the results of XRD and Raman spectroscopy, it is suggested that the degree of structural disordering in pyrocarbon deposited on the TiN-coated substrate (sample (B)) is high in comparison with pyrocarbon deposited directly on carbonized substrate (sample (A)). The plateau of low voltage is attributed to the disordering structure of pyrocarbon in sample (B). The charge capacity is 442–460 mAhg1 per pyrocarbon weight in sample (B). This value is larger than the theoretical capacity of graphite (372 mAhg1), corresponding to the stage 1 lithium-intercalated graphite, LiC6. It is supposed that excess lithium is stored in the nanosized pores in the pyrocarbon of sample (B).

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Thus, the pyrocarbon deposited on the TiN-coated substrate in sample (B) showed a larger capacity than graphite, but the capacity per total electrode mass became lower because of the significantly low capacity of the TiN-coated substrate used as the conductive filler or the current corrector. On the other hand, the charge capacity was 290–298 mAhg1 per total mass of carbonized wood or paper/pyrocarbon sample (A). The carbonized substrates in sample (A) could act not only as the conductive filler or current collector but also as the host material for lithium intercalation/de-intercalation, whereas thick TiN film prevents the intercalation of lithium into the carbon substrate for sample (B). In addition, the first coulombic efficiency of sample (A) is 85–87%, which is higher than 68–73% for sample (B). As mentioned above, the degree of structural disordering of pyrocarbon in sample (B) is high in comparison with pyrocarbon in sample (A), and BET surface area of sample (B) is also higher than that of sample (A). These structural features of pyrocarbon in sample (B) would cause the increase of irreversible reactions such as the decomposition of electrolytes, resulting in low coulombic efficiency at the first cycle. In order to achieve a high capacity per total mass of electrode and coulombic efficiency for sample (B), it is necessary to control the thickness of the TiN film and to modify the surface of pyrocarbon with a disordered structure. 3.3. Structure and electrochemical property of pyrocarbon deposited on Ni

It is well known that some metals such as Ni and Fe promote the graphitization of carbonaceous materials and organic precursors, which is often termed “catalyzed graphitization.” In the case of the CVD process, high crystalline pyrocarbon can be deposited on an Ni substrate at a low temperature below 1000°C [40]. In general, the capacity of graphitized carbon approaches the theoretical value of 372 mAhg1 with an increase in its crystallinity; however, high-temperature treatment above ca. 2800°C is needed in order to form high-crystalline graphite. Catalyzed graphitization is useful for obtaining graphite at low temperatures. Using PCVI technique, pyrocarbon was deposited at a low temperature below 1100°C from the source gases of C6H6 (6%) – H2 into the porous Ni substrate (Ni foam) as catalyst. Pyrocarbon could be deposited at 650°C or higher. Above 900°C, carbon film with laminar structure covered the Ni substrate, whereas thin plate-like particles were formed in the pores of Ni foam below 1000°C. Fig. 8 shows the Raman spectra of pyrocarbon as deposited on Ni at several temperatures. Two Raman shifts are observed at 1580 and 1360 cm1 indicating graphitic structure (E2g2 mode, G band) and disordered structure (A1g mode, D band), respectively. The strong D-band peak is observed for pyrocarbon deposited at 700°C. The peak intensity of the D band is weak, and the peak shape of G band sharpens by increasing the temperature to 850°C. From these results and XRD, it is considered that the crystallinity of pyrocarbon is increased with an increase in the temperature up to 900°C. However, the peak intensity of

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration

G-band

75

D-band

Intensity/a.u.

950° C

850° C

700° C

2000

1800

1600 1400 Raman shift / cm-1

1200

1000

Fig. 8. Raman spectra of pyrocarbon as-deposited on Ni at several temperatures. The number of pulses in PCVI for pyrolytic carbon is 1000.

D band is again increased, and both G and D bands become broad for the samples obtained at 950°C and higher temperatures. In this temperature region, the formation rate of active precursors and nucleation increase rapidly in the gas phase even if no Ni catalyst is used. Low-crystalline carbon would be codeposited at temperatures above 900°C. Fig. 9 shows the XRD patterns of pyrocarbon powder deposited at 850°C, compared with those of high-crystalline natural graphite powder. Pyrocarbon powder is obtained by the removal of Ni using concentrated hydrochloric acid. The lattice parameters are measured from (002) and (110) peaks, and summarized in Table 4. The d002 of pyrocarbon was 0.3360 nm, which is slightly higher than that of natural graphite, but is much lower than that of the low-crystalline carbon formed at ∼1000°C (i.e.  0.340nm). Furthermore, (112) peak was clearly observed on the pyrocarbon deposited at 850°C, which is an evidence for the formation of graphite stacking. It was found by XRD and Raman spectroscopy that high-crystalline pyrocarbon particles having mainly a graphite phase were deposited at 800–900°C. The particle size of pyrocarbon deposited at 850°C was around 1 μm. Crystallite size is also low, compared with that of natural graphite. As a consequence, pyrocarbon powder has a higher surface area than natural graphite powder. Fig. 10 shows the first charge – discharge curves of the pyrocarbon deposited at 850 and 1100°C. For the sample obtained at 850°C, the charge – discharge behavior was similar to that of high-crystalline natural graphite. On the other hand, a “sloping region” in charge – discharge curves was observed for the sample

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Intensity (a.u.)

C(002)

Si Pyrocarbon (x5)

Natural Graphite

24

26

28

Si

30

C(006)

Si

Intensity (a.u.)

C(110)

C(112) Pyrocarbon

Natural Graphite (x3)

75

80 85 Cu Kα 2θ (deg.)

90

Fig. 9. XRD patterns of pyrocarbon powder deposited at 850°C, compared with those of highcrystalline natural graphite powder. Pyrocarbon powder is obtained by removal of Ni using concentrated Hydrochloric acid.

Table 4 Structural data of pyrocarbon deposited in Ni foam at 850°C and natural graphite powder Sample

Lattice constants(nm)

Crystallite size(nm)

Particle size(μm)

BET surface area (m2 g1)

d002

a0

Lc002

La

Pyrocarbon

0.3360

0.2461

48

73

1

10.2

Natural graphite

0.3354

0.2461

∼200

∼500

7

4.8

obtained at 1100°C, resulting from the co-deposition of low crystalline carbon. The sample obtained at 800–900°C showed the highest reversible capacity of 352 mAhg1 at a current density of 30 mA g1. One noticeable feature of the

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration

77

3.0 1100˚C

850°C

+

Potential (V vs. Li/Li )

2.5 2.0 1.5 1.0 0.5 0.0 0

100

200

Capacity (mA hg

300

400

-1)

Fig. 10. First charge – discharge curves of the pyrocarbon deposited at 850 and 1100°C. Current density  30 mA cm2. 110 100

Capacity ratio (%)

90 80 70 60 50 40 30

0

500

1000

1500

2000

Current density (mA g-1)

Fig. 11. Dependence of current density on the capacity for pyrocarbon deposited at 850°C (●) and natural graphite (◆). The current density in discharge is kept at 30 mA g1.

present pyrocarbon is its high-reaction rate. Fig. 11 shows the dependence of current density on the capacity for pyrocarbon deposited at 850°C and natural graphite. It was found that 92% of the capacity at 30mAg1 was maintained even at 1500 mAg1. This high-reaction rate property can be attributed to the small crystalline size and moderately high surface area of pyrocarbon, which effectively

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reduces the diffusion length of lithium ion in carbon and increases the reaction rate at the surface. 4. SURFACE FLUORINATION OF PYROCARBON As mentioned in the previous section, high coulombic efficiency at first cycle was achieved for pyrocarbon deposited directly on the carbonaceous substrate, a result owing to the structural features of pyrocarbon such as high crystallinity, laminar structure and low surface area [25,26]. However, the capacity was slightly decreased by coating with pyrocarbon film as shown in Table 3. This may be because the excessive reduction of surface area resulted in a decrease in the reaction rate for lithium insertion/extraction at the surface. In addition, the lithium ion must diffuse across the graphene layers in pyrocarbon film, which has a laminar structure oriented parallel to the surface of the carbon substrate. Surface fluorination is an effective method for improving electrode kinetics (see Chapter 2). For instance, light fluorination of natural graphite powder by elemental fluorine increases the surface area and mesopores with diameters of 2 and 3 nm, with the result that the high-reaction rate at the surface leads to an increase in the capacity without a decrease in first coulombic efficiency [27–31]. Surface fluorination of pyrocarbon film was also used to improve the electrode characteristics for the carbonized paper/pyrocarbon sample [41], where fluorination treatment was performed between 200 and 500°C by NF3 gas of 3  104 Pa for 10 min in a nickel reactor. Pyrocarbon is infiltrated from the gas phase of C3H8 (30%) – H2 at 950°C by 5000 pulses of PCVI treatment with a holding time of 1.0 s. Table 5 shows the surface fluorine concentrations measured by energy-dispersive X-ray spectroscopy for the samples fluorinated at several temperatures. Fluorine contents were significantly low for the samples fluorinated at 200 and 300°C, whereas increase in the temperature above 400°C rapidly Table 5 Elemental analysis of carbonized paper/pyrocarbon samples fluorinated at several temperatures by energy-dispersive X-ray spectroscopy Fluorination temperature(°C)

Concentration (at%) C

O

F

200

98.6

1.2

0.2

300

98.1

1.7

0.2

400

97.3

1.5

1.2

500

91.8

1.8

6.4

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration

79

increased the fluorine content. When graphite with high crystallinity was fluorinated by F2, several different products were obtained [42]. Fluorine – graphite intercalation compounds with planar graphene layers (sp2 character) are formed below ca. 100°C. On the other hand, graphite fluorides, (CF)n and (C2F)n, with puckered graphene layers (sp3 character) are obtained between ca. 350 and 600°C. In the temperature range of ca. 100–350°C, only the surface of graphite is fluorinated. Surface fluorination of graphite also proceeds at a high temperature, 350°C, with the accelerated fluorination rate. Although the crystallinity of pyrocarbon is much lower than that of graphite, the results for pyrocarbon given in Table 5 coincide well with the trend observed for high-crystalline graphite. The present pyrocarbon film has a laminar texture oriented parallel to the surface of the substrate; therefore, NF3 mainly attack the low reactive basal plane in carbon film. These surface structures may reduce the fluorination rate by NF3. Table 6 shows the structural properties of the original pyrocarbon deposited on the carbonized paper substrate and the samples fluorinated at several temperatures. The d002 obtained by XRD was 0.355 nm for the original pyrocarbon. No remarkable change in d002 was observed for the samples after fluorination treatment. This result indicates that only the surface of pyrocarbon film was fluorinated under the present condition. Raman spectroscopy effectively reveals the surface disorder of carbon materials [43,44]. Two Raman shifts are usually observed at 1580 and 1360 cm1 indicating graphitic structure (E2g2 mode, G band) and disordered structure (A1g mode, D band), respectively. The intensity ratio of Raman shifts (R  ID/IG) is often used to determine the degree of surface disorder. As shown in Table 6, the R value of the original pyrocarbon was slightly decreased by fluorination at 200 or 300°C, and increased above 400°C. These results indicate that the degree of surface disorder on the sample fluorinated at 200–300°C is lower than that of the original pyrocarbon, and that surface disorder Table 6 Structural properties of carbonized paper/pyrocarbon samples fluorinated at several temperatures Fluorination temperature (°C)

d002 by XRD (nm)

R (IDIG) value by Raman Spectroscopy

Surface area by BET (m2 g1)

Original

0.355

0.98

0.58

200

0.355

0.95

0.48

300

0.354

0.89

0.49

400

0.356

1.28

0.80

500

0.355

1.58

3.7

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Yoshimi Ohzawa

is significantly increased by the fluorination at high temperature (400°C or higher). A similar tendency can be observed in the BET surface area of the original pyrocarbon and surface-fluorinated samples. The surface area of each sample fluorinated at 200 or 300°C was lower than that of original pyrocarbon, whereas surface area was increased by fluorination above 400°C. It is supposed that the highly disordered parts of the pyrocarbon surface were eliminated to form CF4 gas during fluorination at 200 and 300°C, which would cause a decrease in surface disorder and surface area in the samples fluorinated at 200 and 300°C. The fluorination level was slightly high (above 400°C). The main reaction would be the formation of C–F covalent bonds accompanying the carbon – carbon bond breaking, which results in the formation of CF2/CF3 groups and lattice defects. These reactions could enlarge the surface disorder and the surface area in the sample fluorinated above 400°C. Fig. 12 shows the pore volume distributions of the original pyrocarbon deposited on the carbonized paper substrate and the samples fluorinated at several temperatures. Pore distributions fluorinated at 200 and 300°C are nearly close to that of the original pyrocarbon. It can be observed that only the mesopores with diameters between 1.5 and 2.0 nm are increased by fluorination at 400°C. Fluorination at 500°C increased not only the small pore but also the relatively large pore with a diameter of 1.5–10 nm. Fig. 13 shows the first charge – discharge curves of the original pyrocarbon deposited on the carbonized paper substrate and the samples fluorinated at 300–500°C. For the samples fluorinated at 300 and 400°C, charge – discharge profiles are similar to that for original pyrocarbon. On the other hand, a small

Pore Volume, dV/dD (X10-10m3g-1)

45 40 35 Fluorination temp. (°C) 30

500

25

400 300

20

200 15

Original

10 5 0 1

10 Pore diameter (nm)

100

Fig. 12. Pore volume distributions of the original pyrocarbon deposited on the carbonized paper substrate and the samples fluorinated at several temperatures.

Synthesis and electrochemical properties of new carbon anodes prepared by chemical vapor infiltration 3.5

Original

3.0

Potential (V, vs. Li/Li+)

Potential (V, vs. Li/Li+)

3.5

2.5 2.0 1.5 1.0 0.5 0.0

300 °C

3.0 2.5 2.0 1.5 1.0 0.5 0.0

0

100

200

300

0

400

-1

100

200

300

400

Capacity (mA h g-1)

Capacity (mA h g ) 3.5

4.0

400 °C

3.0

Potential (V, vs. Li/Li+)

Potential (V, vs. Li/Li+)

81

2.5 2.0 1.5 1.0 0.5 0.0

500 °C

3.5 3.0 2.5 2.0 1.5 1.0 0.5 0.0

0

100

200

300

400

0

Capacity (mA hg-1)

100

200

300

400

Capacity ( mA h g-1 )

Fig. 13. First charge – discharge curves of the original pyrocarbon deposited on the carbonized paper substrate and the samples fluorinated at 300–500°C. Current density  0.2 mA cm2.

Table 7 Capacity and coulombic efficiency data at first cycle of original pyrocarbon and surfacefluorinated samples Capacity (mAh g1)

Coulombic efficiency (%)

Original

298

85

200

300

85

300

298

83

400

320

82

500

310

72

Fluorinated temperature (°C)

plateau appeared at about 1.7 V in the first reduction curve of the sample fluorinated at 500°C. It is also observed that the discharge capacity between 0.5 and 1.0 V is larger than those of the original pyrocarbon and the samples fluorinated at 400°C or lower. Table 7 shows the capacity and coulombic efficiency data at

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Yoshimi Ohzawa

first cycle of the original pyrocarbon and surface-fluorinated samples. The capacities at the initial stage of charge – discharge cycling are also shown in Fig.14. No change is observed between the original pyrocarbon and the samples fluorinated at 200 and 300°C. The effect of surface fluorination on capacity is remarkable for the sample fluorinated at 400°C, i.e. the capacity was increased by 7% compared with the original pyrocarbon. Surface fluorination at 400°C enlarged the surface area of pyrocarbon as shown in Table 6. The enlargement of surface area effectively increases the reaction rate for pyrocarbon electrode. Surface fluorination at 400°C or higher also changed the pore volume distribution of the pyrocarbon surface as shown in Fig. 12. Fluorination at 400°C increased mainly the mesopores with the diameters between 1.5 and 2.0 nm, in which excess lithium would be stored. The first coulombic efficiencies are also summarized in Table 7. The decrease in coulombic efficiency was extremely small for the samples fluorinated below 400°C. This means that the changes in the surface area and the fluorine contents on the surface are too small to affect the coulombic efficiency. On the other hand, the first coulombic efficiency was significantly reduced by fluorination at 500°C. Both the fluorine content and the surface area of the sample fluorinated at 500°C were much higher than those of the samples fluorinated below 500°C. The excessive enlargement of surface area and the large amount of fluorine results in the increase of irreversible reactions on the carbon surface. Thus, in order to improve the electrochemical characteristics of pyrocarbon-based anode, the suitable temperature is considered to be around 400°C for the surface fluorination 340 Fluorination temp. (°C)

330 Capacity (mA h g-1)

400 320 500 200 300

310 300 290

Original 280 270

0

1

2

3 4 Cycle number

5

6

Fig. 14. Charge capacities (Li deintercalation) of original pyrocarbon and samples fluorinated at several temperatures as a function of cycle number at the initial stage of charge – discharge tests. Current density  0.2 mA cm2.

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of pyrocarbon by NF3. Under the present condition, light fluorination of pyrocarbon film moderately increases the surface area and mesopores with diameters of 1.5–3 nm, so that the high-reaction rate at the surface leads to an increase in the capacity without a significant decrease in the first coulombic efficiency. 5. SURFACE MODIFICATION BY COATING WITH PYROCARBON ON LOW-CRYSTALLINE CARBON Electrochemical properties of carbonaceous anodes strongly depend not only on the bulk structure but also on the surface properties such as crystallinity, surface area and chemical species, because the electrochemical reactions occur on the surface of the electrode. Surface modification is one of the effective methods for improving the electrochemical characteristics of carbonaceous anodes for LIB. Several methods of surface modification were applied to carbon materials, involving surface oxidation [45–47], surface fluorination [27–31], thin metal coating [48] and carbon coating [8–12,49]. Light oxidation of carbon materials increases their capacities due to the formation of nanochannels at the surface. However, strong oxidation degraded the surface structure, leading to an increase in the irreversible capacity [45,46]. As shown in Chapter 2 and in the previous section in this chapter, surface fluorination is an effective method to improve the electrochemical characteristics of carbon anodes. Pyrocarbon coating by CVD was recently applied to graphite-based anodes of lithium-ion secondary battery to improve the anode performance, especially in propylene carbonate (PC) based solvent [8–12]. In this process, a low-crystalline pyrocarbon shell covered the surface of a high-crystalline graphite core; therefore, the decomposition of the PC solvent was successfully prevented. In order to achieve high coulombic efficiency in the first cycle without reducing the reversible capacity, the coating of thin pyrocarbon film with uniform thickness is desired because the capacity of pyrocarbon (soft carbon) is generally lower than that of high-crystalline graphite. As mentioned in Section 1, it is necessary to repeat the CVD treatment and the soft grind of particles several times in order to achieve a uniform coating on the conventional CVD. Rotation of a reactor was recently attempted to obtain a uniformly coated carbon layer [12]. The PCVI process allows homogeneous infiltration of matrix through the thickness of the porous substrate under suitable conditions because of the rapid penetration of the source gas throughout the preform without preheating [23,24]. In addition, instantaneous introduction of the source gas and evacuation of the reacting gas are repeated at relatively short intervals in the PCVI process; therefore, nucleation in the gas phase is restrained, and crystal growth is promoted in the evacuation step. These are effective in increasing the crystallinity of pyrocarbon, resulting in high coulombic efficiency at first cycle [25]. In case low crystalline carbon such as non-graphitizing carbon (hard carbon) is used as the core

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material, the pyrocarbon shell has higher crystallinity compared with the core carbon. This structure would lead to a reduction in the irreversible capacity observed in the naked hard carbon. Pyrocarbon was coated on powdery hard-carbon beads with an average diameter of 3 μm from CH4 (50%) – H2 at 1100°C by PCVI technique [50]. Fig. 15 shows the SEM images of the original hard-carbon powder (a) and the sample coated with 7 wt% pyrocarbon (b). From the wide range of the SEM observation, it appears that the thin pyrocarbon films are coated on most of the surfaces of the spherical carbon particles. It can be observed that the pyrocarbon film has pebblelike projections sized below 0.1 μm. Therefore, the roughness of the particle surface appears to increase at the submicron scale by coating with the pebble-like pyrocarbon film. Table 8 shows the structural properties of the original hard-carbon powder and the sample coated with pyrocarbon. The d002 of the pyrocarbon-coated sample was 0.348 nm, which is much lower than that of the original hard carbon. The R value calculated from Raman shifts of the original pyrocarbon was slightly decreased by coating with pyrocarbon. These results by XRD and Raman spectroscopy demonstrate that the crystallinity of the pyrocarbon shell is higher than that of core carbon. The BET surface area was decreased from 25 m2g1 of the original particles to 8.5 m2g1 after coating with 7 wt% pyrocarbon. From pore

Fig. 15. SEM images of the original hard-carbon powder (a) and the sample coated with 7 wt% pyrocarbon (b). The number of pulses in PCVI for pyrolytic carbon is 500.

Table 8 Structural properties of original carbon beads and pyrocarbon-coated sample Sample

d002 by XRDa R (IDIG) value by Raman Surface area by BETb (nm) spectroscopya (m2 g1)

Original carbon beads

0.373

1.42

25

Pyrocarbon-coated carbon beads

0.348

1.22

8.5

a Measured for the sample with 47% pyrocarbon after 5000 pulses in PCVI treatment. bMeasured for the sample with 7% pyrocarbon after 500 pulses in PCVI treatment.

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volume distribution analysis, it was found that the pores with diameters of 1.5–5 nm were extremely decreased by pyrocarbon coating; however, the volume of larger pores above 10 nm was rather increased. Fig. 16 shows the charge–discharge curves at first and 10th cycles of the original carbon powder (a), the samples coated with 7 wt% pyrocarbon (b) and

Potential (V vs Li/Li+)

3.5 3.0

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2.5 2.0 1.5 1.0 0.5 0.0

0

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Potential (V vs Li/Li+)

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(b)

10th

2.5 2.0 1.5 1.0 0.5 0.0

0

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600

400

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Potential (V vs Li/Li+)

3.5 1st

3.0

(c) 10th

2.5 2.0 1.5 1.0 0.5 0.0 0

200

Capacity (mA h g-1)

Fig. 16. Charge – discharge curves at the first and 10th cycles of the original carbon powder (a), the samples coated with 7 wt% pyrocarbon (b), and 47 wt% pyrocarbon (c). The number of pulses in PCVI for pyrolytic carbon: (b) 500, (c) 5000. Discharging, constant current of 60 mA g1 followed at constant potential of 3 mV vs. Li/Li for 24 h; charging, constant current of 60 mA g1.

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47 wt% pyrocarbon (c). The charge–discharge profiles of the original carbon powder (a) and the samples coated with 7 wt% pyrocarbon (b) are similar to that observed in typical non-graphitizing carbon with a disordered structure. High irreversible capacity of 200 mAh g1 was observed in the original carbon beads, reflecting the low crystalline disordered structure and high surface area. Irreversible capacity was reduced to around 100 mAh g1 by coating with 7 wt% pyrocarbon. As mentioned above, the coated pyrocarbon film has higher crystallinity and lower surface area than the core carbon. In addition, it is supposed that the surface functional groups containing oxygen atoms will be decreased while pyrocarbon is coated under the reductive condition with C3H8–H2 gas system at high temperatures. These structural features of pyrocarbon might cause the decrease in the rates of irreversible reactions such as decomposition of the electrolytes and the trapping of lithium ions. As shown in Fig. 16(c), irreversible capacity was decreased with increasing the mass fraction of pyrocarbon, however, the charge capacity (Li de-intercalation) was also decreased. Thin pyrocarbon film with uniform thickness is desired in order to achieve high-coulombic efficiency at the first cycle without reducing the reversible capacity. For the original carbon beads, the decrease in the charge capacity is observed after the 10th charge – discharge cycling as shown in Fig. 16(a). The cycleability can be improved by coating with pyrocarbon (Fig. 16(b) and (c)). From the SEM images, it was observed that the roughness of the particle surface appeared to increase at the submicron scale by coating with the pebble-like pyrocarbon film, which would effectively increase the adhesion among the particles by the organic binders. ACKNOWLEDGMENTS The author gratefully acknowledges Prof. T. Nakajima for many valuable discussions and suggestions. Special thanks are due to Prof. B. Zemva for preparation of the surface-fluorinated pyrocarbon samples and for fruitful discussion. Quallion LLC and Mitsui Miring Co. Ltd are acknowledged for their kind supply of carbon beads. REFERENCES [1] [2] [3] [4] [5] [6] [7]

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[41] Y. Ohzawa, V. Gupta, B. Zemva, and T. Nakajima, Abstract of the Third French-Japanese Seminar on Fluorine in Inorganic Chemistry and Electrochemistry (3rd FJSF-2003), Paris, France, 2003, pp. 25–27. [42] T. Nakajima, Fluorine-Carbon and Fluoride-Carbon Materials – Chemistry, Physics and Applications, Marcel Dekker, New York, USA., 1995, pp.1–31, Chap. 1. [43] F. Tunistra and J.L. Koenig, J. Chem. Phys., 53 (1970) 1126. [44] D.S. Night and W.B. White, J. Mater. Res., 4 (1989) 385. [45] J.S. Xue and J.R. Dahn, J. Electrochem. Soc., 142 (1995) 3668. [46] E. Peled, C. Menachem, D. Bar-Tow, and A. Melman, J. Electrochem. Soc., 143 (1996) L 4. [47] M. Hara, A. Satoh, N. Tamaki, and T. Ohsaki, Tanso, 165 (1994) 261. [48] R. Takagi, T. Okubo, K. Sekine, and T. Takamura, Denki Kagaku, 65 (1997) 333. [49] Y.Sato, Y. Kikuchi, T. Nakano, G. Okuno, K. Kobayakawa, T. Kawai, and A. Yokoyama, J. Power Sources, 81–82 (1999) 182. [50] Y. Ohzawa, Y. Yamanaka, K. Naga, and J. Li. R.Chandrasekaran, T. Nakajima, Abstract of International Meeting on Lithium Batteries (IMLB 12), Nara, Japan, 2004, Abs. No. 60.

Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 4

Electrochemical properties of fluorinated carbon nanotubes Hidekazu Touhara Department of Chemistry, Faculty of Textile Science and Technology, Shinshu University, 3-15-1 Tokida, Ueda 386-8567, Japan 1. INTRODUCTION Fluorination is one of the most effective chemical methods of modifying and controlling the physicochemical properties of carbon materials over a wide range. It provides an opportunity to prepare a very broad array of fluorine–carbon materials with new functionalities [1,2]. Of the many forms of carbon materials, the fluorination of carbon nanotubes (CNTs) is of great interest from the point of view of fluorine doping or intercalation and sidewall chemical functionality. It is already known that both single-wall carbon nanotubes (SWNTs) and multi-wall carbon nanotubes (MWNTs) are amphoteric in their chemical properties, and form donor and acceptor compounds such as graphite intercalation compounds (GICs) [3–8]. These properties, together with their one-dimensional-type host lattice structure with a central hollow core, suggest energy- related applications such as lithium cells, electric double-layer capacitor and fuel cells. In addition to these interesting features, the fluorination of CNTs is expected to bring about further functionalities for energy storage and conversion. The fluorination chemistry of CNTs is extensive and diverse, and a wide array of fluorinated CNTs have so far been prepared. In this chapter, the electrochemical properties of fluorinated carbon nanotubes are presented with an emphasis on electrochemical lithium storage by fluorinated MWNTs and on discharge performances of primary lithium cells with fluorinated CNT cathodes. Before dealing with the topic of lithium cells with fluorinated CNTs cathodes, we will briefly review the fluorination of CNTs.

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2. PRIMARY LITHIUM CELLS WITH FLUORINATED MULTI-WALL CARBON NANOTUBES A primary battery based on fluoride and lithium is, theoretically, the optimum redox system for a high-energy density power source. Such a battery has been realized commercially by the use of graphite fluoride (CF)n as a cathode material for lithium battery with an electrolyte – aprotic solvent system. In this context, the potential use of fluorinated MWNTs as high-density energy conversion materials has made it important and interesting to investigate the performances of lithium cells with fluorinated MWNTs cathodes, particularly deeply fluorinated MWNTs. The electrochemical properties of fluorinated MWNTs (F-MWNTs) as a cathode material for lithium cells were investigated on the Li/1M LiClO4-PC/FMWNTs cell (PC  propylene carbonate) [9]. The pristine MWNTs used were prepared by the thermal decomposition of acetylene over silica-supported cobalt catalysts. The optimized experiments allowed long MWNTs up to 10 μm in length and 20 to 40 nm in outer diameter. Fluorination was carried out at temperatures of 279 and 773 K, using elemental fluorine. The reaction was carried out at room temperature with an F2, HF, and IF5 mixture. Fluorination at 773 K for 4 h yielded white compounds for which XRD patterns were quite similar to those of (CF)n obtained at 873 K fluorination of natural graphite, and TEM (002) lattice fringe images showed the destruction of the tube structure. The values of open-circuit voltage and current densities are not clear in the galvanostatic discharge experiments; however, it should be noted that the (HT)CFx cathode (fluorinated MWNTs at a high temperature of 753 K) shows a flat and stable discharge potential at ca. 2.4 V with a large capacity of 620 Ah/kg. It is well known that discharge performances of conventional cathode materials are strongly dependent on the structural properties of pristine carbon materials, fluorine content, C–F bonding nature, and fluorination temperature. Nevertheless, a cathode behavior similar to that, for example, of fluorinated activated carbon fibers [10], is observed for an F-MWNTs cathode. This suggests that thick F-MWNTs with 30 nm or so in diameter show a discharge performance substantially similar to that of conventional C–F materials. 3. ELECTROCHEMICAL LITHIUM STORAGE BY FLUORINATED MULTI-WALL CARBON NANOTUBES Electrochemical energy storage has been investigated on the template-synthesized MWNT -membrane [11], and on the MWNTs produced by the catalytic decomposition of acetylene [12], using lithium test cells with nonaqueous electrolyte solutions. Early work by Martin et al. has revealed that the highly ordered MWNT -membrane can be a good candidate for the anode material of a lithiumion secondary battery [11].

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We studied in detail the selective fluorination of the inner surfaces of template-synthesized MWNTs and the properties of resultant materials [2,8,13,14], and the improvement in electrochemical lithium storage characteristics by the fluorination of MWNTs was also investigated. The pristine MWNTs were ca. 75 μm in thickness with straight channels of ca. 30 nm in diameter and a central hollow core of ca. 27 nm in diameter. Fluorination was performed by the direct reaction of MWNTs embedded in the Al2O3 template with elemental fluorine, and only the inner surfaces of the tubes were fluorinated. The reaction was run for 5 days in the temperature range of 323 to 473 K with 1 atm of fluorine gas. The TEM and SEM observations revealed that the form and the tubular morphology of MWNTs were unchanged upon fluorination, but the inner-surface (002) lattice fringe image, parallel to the tube axis, suggests a slight modification of the tubes arising from the existence of buckled fluorinated carbon layers. The C–F bonding nature of fluorinated MWNTs was investigated by C1s and F1s XPS spectra [8]. The C1s peaks observed at 288.2–289.4 eV were ascribed to sp3-hybridized carbon atoms with covalent C–F bonds, which are similar to those in the covalent graphite fluorides compounds, (CF)n and (C2F)n. With an increase in temperature, the C1s and F1s peaks assigned to fluorine functional groups shift to higher binding energy side and their intensities become stronger. The surface compositions (F/C) determined by the peak area ratio of C1s and F1s XPS spectra were 0.52, 0.86 and 1.42 at fluorination temperatures of 323, 373 and 473 K, respectively. The peaks assigned to fluorine functional groups almost disappeared by 30 s argon-ion sputtering etching, and the C1s peak became narrower and more symmetric. These depth profiles indicated that fluorination was limited to the outermost layers of surfaces. The close similarities in Raman spectra of pristine SWNTs and F-SWNTs also indicated that the internal carbon atoms below the outermost layers retain their sp2 hybridization [8]. We investigated electrochemical Li insertion in template-synthesized MWNTs, and inner-surface fluorinated MWNTs (F-MWNTs) at 5°C by cyclic voltammetry (CV) and charge – discharge experiments on the Li/1M LiClO4(EC  DEC)/WE ((EC  DEC) represents a mixture of ethylene carbonate (EC) and diethyl carbonate (DEC) (1:1, v/v), WE (working electrode)  pristine MWNTs, and F-MWNTs) cells, and the effects of inner-surface fluorination on the electrochemical properties [13,14]. In the electrochemical measurements, the Al2O3 template was removed by treatment with an aqueous solution of 48% HF. In these test cells, Li is inserted into MWNTs and F-MWNTs during the cathodic (discharge) process and extracted during the anodic (charging) process. Inner-surface fluorination results in significant changes in the electrochemical properties of MWNTs. The OCV (open circuit voltage) value 2.98 V of pristine MWNTs increases up to 3.64 V by fluorination. This high OCV value is comparable to that of (C2F)n prepared by the fluorination of activated carbon fibers [10]. Fig. 1 shows the CVs of pristine MWNTs and F-MWNTs fluorinated

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Fig. 1. Cyclic voltammograms for (A) pristine multi-wall carbon nanotubes and (B) innersurface fluorinated multi-wall carbon nanotubes at 323 K at the sweep rate of 0.1 mV/s (reproduced with permission from Touhara et al. [14]).

at 323 K [14]. It can be seen from Fig. 1(A) that the pristine MWNTs show three major peaks. Peaks A (at 1.4 V) and B (at 1.26 V) can be attributed to the reduction of surface-oxygenated functional groups and the electrolyte LiClO4, respectively. Peak C (at 0.27 V) is due to the formation of solid electrolyte interphase (SEI) by the decomposition of solvent EC to gaseous ethylene and lithium insertion into available defect sites of MWNTs. The discharge scan is completed with a sharp current declination from 0.25 V due to Li ion intercalation into graphene layers parallel to the tube axis of the MWNTs. On the reverse sweep (first charge) a current plateau starting at 0.4 V and extending to around 1.5 V is noticed. This plateau is attributed to different extraction processes of Li from various insertion sites of the MWNTs, occurring at different oxidation potentials. In the case of FMWNTs, the first discharge sweep starts with the reduction of C–F bonds (peak B at 3.0 V). Current peak A, due to the reduction of oxygenated groups and electrolyte, becomes negligibly small in Fig. 1(B). The replacement of oxygenated groups by elemental fluorine leads to that observation. It is noted that the peak due to the reduction of EC (peak C of pristine in Fig. 1(A) is shifted to more positive potentials for F-MWNTs. Inner-surface fluorination removes the amorphous carbon on the inner surfaces of the CNM [8] and produces a relatively defect-free microtexture, which accounts for the decrease in voltage hysteresis. The galvanostatic discharge – charge curves of pristine MWNTs and F-MWNTs for the 1st to 5th cycle are shown in Fig. 2. In Fig. 2(A), a large irreversible capacity in the reduction processes was observed for the first cycle. This is attributed to the large BET surface area of 28 m2/g, which is about 5–20 times that of powder graphite, and to the presence of oxygenated functional groups on the MWNTs. The reversibility of the insertion and extraction process is found to be enhanced greatly in the second and consecutive cycles. For example, it is

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Fig. 2. Galvanostatic discharge – charge curves for (A) pristine multi-wall carbon nanotubes and (B) inner-surface fluorinated multi-wall carbon nanotubes at 323 K under the current density of 50 mA/g (reproduced with permission from Touhara et al. [14]).

observed from galvanostatic charge – discharge experiments that the 1st cycle coulomb efficiency of 22% is enhanced to about 82% with a reversible capacity of 460 mAh/g at 6th cycle. The effect of the fluorination of MWNTs on their discharge – charge properties is shown in Fig. 2(B). The irreversible capacity during the 1st cycle is found to be common for both pristine and fluorinated MWNTs. It is worth remarking that there is a drastic increase in columbic efficiency from ca. 23% to ca. 90% with a reversible capacity of ca. 400 mAh/g after the second cycle. Surface lithiation and formation of SEI is believed to be the reason for the irreversible charge loss in the first discharge sweep, which is reduced in the second and consecutive sweeps solely due to the aging of the SEI. 4. PRIMARY LITHIUM CELLS WITH FLUORINATED SINGLE-WALL CARBON NANOTUBES 4.1. Fluorination of single-walled carbon nanotubes, structure and properties

Fluorination of SWNTs has currently attracted considerable interest in nanoscience and nanotechnology, and in potential applications [16–23]. Mickelson et al. [15], were the first to report the fluorination of purified and end-closed SWNTs which is of great interest for a wide variety of sidewall chemical functionalizations. The SWNTs used were produced by the dual-pulsed laser vaporization of Co/Ni-doped graphite rods. Fluorination was carried out on “bucky paper,” which is a free-standing film (10 μm thick), using elemental fluorine diluted with helium, in the temperature range 423–873 K for 5 h. The stoichiometries, by gravimetry, of the products fluorinated at 423, 523, 598, and 623 K, were CF0.114, CF0.521, CF0.495, and CF0.565, respectively. There is a limiting stoichiometry of C2F

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for which the outside-wall-fluorinated SWNTs can still maintain their tube-like structure. This seems to be supported by the fluorine content of the products determined by gravimetric analysis. The values of F/C ratio were close to 0.5 for the products whose tube-like structure remained intact after lower temperature fluorination at 523 and 623 K, while the values were fairly higher than 0.5 for the materials obtained by fluorination at higher temperatures which led to the breaking of C–C bonds and, hence, destruction of the tube. Upon fluorination, the electronic properties of the SWNTs changed drastically; while the pristine SWNTs were good conductors (10–15 Ω two-probe resistance across the length of the ca. 10 mm × 3 mm × 30 μm bucky paper samples), the tubes fluorinated at temperatures of 523 K and above were insulators (two-probe resistance 20 MΩ). TEM studies have shown that at fluorination temperatures as high as 598 K, the majority of the fluorination products maintained a tube-like structure. Infrared spectroscopy confirmed the presence of covalently bonded fluorine (peaks in 1220–1250 cm1) in the sample fluorinated at a temperature of 523 K and higher. No C–F stretching frequency was observed for the sample fluorinated at 423 K. Infrared spectroscopy, together with the product stoichiometries and resistance measurements, indicated that reaction temperatures higher than 150°C were necessary to covalently add significant amounts of fluorine to the tube wall. Another interesting feature is the defluorination of once fluorinated SWNTs by anhydrous hydrazine via the following reaction: CFn  14 nN2H4 → C  nHF  12 nN2 We have investigated the fluorination reactions of end-closed SWNTs (C-SWNTs) and open-end SWNTs (O-SWNTs) [23]. High-purity end-closed SWNTs (C-SWNTs) were prepared by the laser-ablation method using a metalcarbon composite rod as target [24]. The diameters of the tubes were in the range of 1.4 to 1.5 nm. The open-end of the tubes confirmed that O-SWNTs adsorb almost twice the quantity of Ar gas compared with C-SWNTs. Open-end SWNTs (O-SWNTs) were obtained by heat treatment of C-SWNTs at 698 K in air. Fluorination was carried out using 1 atm elemental fluorine at a temperature ranging from room temperature to 523 K. The XRD profiles of pristine and endclosed F-SWNTs prepared by 473 K fluorination are given in Fig. 3 [23,28]. Reflections can be indexed in terms of a 2-D triangular lattice. Upon fluorination, the lattice constant a  1.74 nm of the pristine tube increases up to ca. 2 nm. The most important observation is that the fluorination mechanism between C-SWNTs and O-SWNTs is quite different. Fluorine atoms can access carbon atoms not only from the outside of the tube but also from the inside of the tube, they are concentrated in the restricted part, and the nonreacted part of SWNTs remains up to F/C  0.5. On the other hand, in the case of C-SWNT samples,

(21)

Gr (002)

95

F-SWNTs

0

10

(21)

(20)

(11)

(10)

Intensity/arb.units

(11)

(10)

Electrochemical properties of fluorinated carbon nanotubes

20 30 2θ (CuKα)/deg.

Pristine

40

Fig. 3. XRD patterns of pristine and fluorinated SWNTs (F-tube) CF0.48 obtained by 473 K fluorination of end-closed SWNTs prepared by laser-ablation method (reproduced with permission from Touhara and Yamada [27]).

since fluorine atoms cannot access the inside wall of the tube, a wider area is attacked by the same amount of fluorine atoms compared with O-SWNT samples. The selective fluorination on the outside of C-SWNT tubes results in a larger lattice constant than fluorinated O-SWNTs with the same F/C value, and also results in the disappearance of RBM (radial breathing mode) in the Raman spectrum of fluorinated C-SWNTs with F/C  0.48. In the case of fluorinated O-SWNTs with F/C  0.52, RBM is clearly observed [23]. We have also investigated the fluorination reaction of HiPco (high pressure carbon monoxide) – SWNTs [25] using 1 atm elemental fluorine at a temperature ranging from room temperature to 573 K. The behavior of HiPco – SWNTs to fluorination was almost similar to that of C-MWNTs; however, both pristine HiPco – SWNTs and F-SWNTs show no diffraction line in XRD patterns, which is the most important difference between HiPco – SWNTs and SWNTs prepared by the laser-ablation method. Fig. 4 shows the XPS C1s spectra of pristine HiPc – SWNTs and F – SWNTs obtained by fluorination at 473 K. The peak assigned to the C–F bond is observed at ca. 288 eV, indicating semi-ionic – covalent of the bond nature in C–F. Dispersion property of fluorinated tubes in ethanol and water strongly improved because of the formation of semi-ionic C–F bonds on the sidewall of tubes. The F/C values determined by the XPS spectra increase with an increase in fluorination temperature, and the composition of fluorinated

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285.0 e V 288.0 e V

Intensity / arb. units

291.0 eV

F-SWNTs

284. 5 eV

Pristine

300

290 Binding Energy/eV

280

Fig. 4. C1s XPS spectra of pristine and fluorinated SWNTs (CF0.48) at 473 K (reproduced with permission from Touhara and Yamada [27]).

HiPco – SWNTs at 301 and 473 K were CF0.28 and CF0.48, respectively. The value 0.48 of F-SWNTs fluorinated at 473 K is close to the limiting composition C2F of the saturated fluorotube. TEM and SEM images are given in Figs. 5 and 6. TEM and SEM observations indicate that the bundle structure and tubular morphology of pristine SWNTs were preserved up to 473 K fluorination. Further fluorination at 573 K led to the breaking of C–C bonds, and hence, the partial destruction of tubes was observed. 4.2. Electrochemical properties of lithium cells with fluorinated single-wall carbon nanotubes (fluorotubes)

Fluorinated SWNTs are reminiscent of graphite – fluorine compounds and fluorinated fullerenes whose C–F bonds undergo electrochemical reduction, and are cathode materials of lithium–carbon fluoride electrochemical power sources [24]. In this section, fluorinated SWNTS are conveniently abbreviated as fluorotubes [17]. Peng et al. [17] studied in detail the discharging performance of twoelectrode lithium cells. Li/1M LiBF4–(DME  PC)/fluorotubes (DME,1,2dimethoxyetahne; PC,propylene carbonate), and compared it with that of a Li/carbon monofluoride (CFx) cell. Fluorotubes (C2F) as cathodes were prepared by fluorination laser oven SWNTs (L-SWNTs) and HiPco-SWNTs (H-SWNTs)

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Fig. 5. TEM images of pristine and fluorinated SWNTs (CF0.48) at 473 K (reproduced with permission from Touhara and Yamada [27]).

Fig. 6. SEM images of pristine and fluorinated SWNTs at 473 K (CF0.48) (reproduced with permission from Touhara and Yamada [27]).

with F2/HF at 523 K for 12 h. The discharge experiments were carried out under a 2 kΩ resistance load. The discharge curves for Li/fluorotubes cells show two distinct slopes with the inflection point occurring when ⬃65–70% of the fluorine was removed. This suggests that C2F fluorotubes are first converted into C4F while fluorine is removed. When compared with the Li/ CFx cell, the cell potential, which was obtained across 200 kΩ resistors, for Li/fluorotubes (fluorinated

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L-SWNTs) cell was ⬃0.7 V higher than that of Li/ CFx cell. During discharge, the potential of the cell drops more rapidly for the Li/ fluorotube cell than for the Li/CFx cell. These results indicate the following two interesting features: First, that the large change in discharge potential is due to a C–F bond energy dependence on the extent of sidewall fluorination; C–F bonds are the strongest when the nanotube is slightly fluorinated. Secondly, that the C–F bonds of the fluorotube are weaker than those of the carbon monofluoride. It is interesting to note that thermodynamic calculations using the heat of formation ΔHf0, of fluorotubes from (10,10) arm chair SWNTs and carbon monofluoride indicated a 0.4 V higher potential for the Li/fluorotube cell than for the Li/carbon monofluoride cell. We studied the discharging performance of fluorotubes on a Li/1MLiClO4–(EC  DEC)/ F-SWNTs cell under a current density of 100 μA/cm2. A typical discharge curve of fluorotubes CF0.48 obtained by the fluorination of purified HiPco tubes at 473 K is shown in Fig. 7 together with that of graphite fluoride (CF)n as a comparison material. The OCV value 3.9 V vs. Li reference of fluorotubes is ca. 0.7 V higher than that of (CF)n, which vividly reflects the lower bond energy of the semi-ionic C–F bond in the fluorotubes. The discharge potential of the fluorotubes decreases with increasing cathode utilization. It was also observed that the OCV value of the fluorotubes decreases with increasing cathode utilization. In comparison, in (CF)n, as is well known, the discharge potential of

4

OCV=3.9V

Potential vs. Li ref. /V

F-SWNTs

3

OCV=3.2 V (CF)n

2

1 0

50 Cathode utilization /%

100

Fig. 7. Discharge curves of Li/1M-LiClO4–(ECDEC)/ F-SWNTs (CF0.48) and (CF)n cells with a current density of 0.1 mA/cm2 (reproduced with permission from Touhara and Yamada [27]).

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the electrode is flat until cathode utilization reaches 80%, and the OCV values are also unchanged during discharge. The original nanotube was recovered by the discharge of fluorotube cathode without any additive and was characterized by SEM and TEM observations and by XPS and Raman spectoroscopy. Figs. 6 and 8 show SEM and TEM images, respectively, of discharged fluorotubes. These images indicate that tubular morphology and bundle structures are preserved after the electrochemical reduction of C–F bonds. The STEM image of discharged fluorochubs is given in Fig. 9, where one can clearly see the LiF crystal formed during the discharge. During the discharge

Intensity / arb. units

Fig. 8. STEM and TEM images of 100% discharged F-SWNTs (CF0.48) (by courtesy of Dr. Yanagiuchi, TDK Corporation).

100% discharge

CF0.05

80% discharge

CF0.10

70% discharge

CF0.15

F-SWNT

CF0.48

Pristine

300

290 Binding energy/eV

280

Fig. 9. C1s spectra of pristine, F-SWNTs (CF0.48), 70, 80, and 100% discharged products.

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of flurotube CF0.48 , XPS C1s and RBM of Raman spectra changed as shown in Figs. 9 and 10, respectively. The relative intensity of the peak assigned to C–F bond at ca. 288 eV decreases with increasing discharge percentage, and almost disappeared at 100% discharge. The fluorotube CF0.48 shows no RBM mode, but it is completely recovered to that of pristine SWNTs, as is shown in Fig. 10. The discharge reaction of (CF)n cathode is known to proceed as follows [27]. (CF)n  xLi  zSol. → (CF)nx  x[C-(F-Li)] ⋅ zSol. → (CF)nx  xC  xLiF  zSol.,

(1)

where x[C-(F-Li)] ⋅ zSol. is the ternary intermediate phase formed during the discharge of (CF)n cathode and z is the salvation number. On the other hand, the OCV values of fluorotubes decrease with increasing cathode utilization. These results clearly indicate that the discharge reaction of fluorotubes is quite different from that of (CF)n, and on the basis of relevant experimental evidence, the discharge reaction in the Li/fluorotubes cell has been deduced as. CFx  yLi → CFxy  yLiF

(2)

The discharge of fluorotube CF0.48 electrode proceeds homogeneously, forming discharged product CF0.48-0.48x and LiF, where the fluorine concentration decreases with the discharge ratio x (0  x 1.0).

Intensity / arb. units

100% discharge

80% discharge

F-SWNTs(CF0.48)

pristine

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200 Raman Shift / cm-1

300

Fig. 10. Raman RBM spectra of pristine, F-SWNTs (CF0.48), 80, and 100% discharged products.

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As a concluding remark, it should be pointed out that the electrochemical study of flurotubes is considerably useful for a deeper understanding of nanostructure and properties of carbon nanotubes. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14]

[15] [16] [17] [18] [19] [20] [21] [22] [23] [24]

F. Okino and H. Touhara, Comprehensive Supramolecular Chemistry, G. Alberti and T. Bein (Eds.), Pergamon, Oxford, 1996, Vol. 7, Chapter 2. pp. 25–76. H. Touhara and F. Okino, Carbon, 38 (2000) 242. O. Zhou, R.M. Fleming, D.W. Murphy, C.H. Chen, R.C. Haddon, A.P. Ramirez, and S.H. Glarum, Science, 263 (1994) 1744. R.S. Lee, H.J. Kim, J.E. Fisher, A. Thess, and R.E. Smalley, Nature, 388 (1997) 255. A.M. Rao, P.C. Eklund, S. Bandow, A. Thess, and R.E. Smalley, Nature, 388 (1997) 257. V.A. Nalimova, D.E. Sklovsky, G.N. Bondarenko, H. Alvergnat-Gaucher, S. Bonnamy, and F. Beguin, Synth. Met., 88 (1997) 89. T. Nakajima, S. Kasamatsu, and Y. Matsuno, Europ. J. Solid State Inorg. Chem., 33 (1996) 831. Y. Hattori, Y. Watanabe, S. Kawasaki, F. Okino, B.K. Pradham, T. Kyotani, A. Tomita, and H. Touhara, Carbon, 37 (1999) 1033. A. Hamwi, P. Gendraud, H. Gaucher, S. Bonnamy, and F. Beguin, Mol. Cryst. Liq. Cryst., 310 (1998) 185. H. Touhara, K. Kadono, N. Watanabe, and J.-J. Braconnier, J. Electrochem. Soc., 134 (1987) 1071. G. Che, B.B. Lakshmi, E.R. Fisher, and C.R. Martin, Nature, 393 (1998) 346. E. Frackowiak, S. Gautier, H. Gaucher, S. Bonnamy, and F. Beguin, Carbon, 37 (1999) 61. H. Touhara, Carbon Alloy, E. Yasuda, M. Inagaki, K. Kaneko, M. Endou, A. Oya, and Y. Tanabe (Eds.), Elesevier, Amsterdam, 2003, p. 485, Chapter 30. H. Touhara, J. Inahara, T. Mizuno, Y. Yokoyama, S. Okano, K. Yanagiuchi, I. Mukopadhyay, S. Kawasaki, F. Okino, H. SHirai, W.H. Xu, T. Kyotani, and A. Tomita, J. Fluorine Chem., 114 (2002) 181. E.T. Mickelson, C.B. Huffman, A.G. Rinzler, R.E. Smalley, and J.L. Margrave, Chem. Phys.Lett., 296 (1998) 188. E.T. Mickelson, I.W. chiang, J.L. Zimmerman, P.J. Boul, J. Lozano, J. Liu, R.E. Smalley, R.H. Hauge, and J.L. Margrave, J. Phys. Chem. B, 103 (1998) 4318. H. Peng, Z. Gu, J. Yang, J.L. Zimmerman, P.A. Willis, M.J. Bronikowski, R.E. Smalley, R.H. Hauge, and J.L. Margrave, Nano Lett., 1 (2001) 625. Z. Gu, H. Peng, R.H. Hauge, R.E. Smalley, and J.L. Margrave, Nano Lett., 2 (2002)1009. P.R. Marcoux, J. Schreiber, P. Batail, S. Lefrant, and J. Renouard, Phys. Chem. Chem. Phys., 4 (2002) 2278. R. Bandyopadhyaya, E.N.-Roth, O. Regev, and R.Y.-Rozen, Nano Lett., 2 (2002) 25. Y.S. Lee, T.H. Cho, B.K. Kee, J.S. Rho, K.H. An, and Y.H. Kee, J. Fluorine Chem., 120 (2003) 99. P.E. Pehrsson, W. Zhao, J.W. Baldwin, C. Song, J. Liu, S. Kooi, and B. Zheng, J. Phys. Chem. B, 107 (2003) 5690. S. Kawasaki, K. Komatsu, F. Okino, H. Touhara, and K. Kataura, Phys. Chem. Chem. Phys., 6 (2004) 1769. H. Kataura, Y. Maniwa, M. Abe, A. Fujiwara, T. Kodama, K. Kikuchi, H. Imahori, Y. Misaki, S. Suzuki, and Y. Achiba, Appl. Phys. A, 74 (2002) 349.

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[25] H.Touhara and F. Okino, Advanced Inorganic Fluorides, T. Nakajima, A. Tressaud, and B. Cemva, (Eds.), Elsevier, Amsterdam, 2000, Chapter 17. [26] H. Touhara, The Japan-Korea Joint Seminar on Fluorine Chemistry Abstracts, JSPS-155 Fluorine Committee, (2003) 17. [27] H. Touhara and I. Yamada (Eds.), Introduction to Fluorine Chemistry, JSPS-155 Fluorine Committee, Sankyo Pub. Co., Tokyo, 2004, pp. 362–370, Chapter 8.11. [28] H. Touhara, H, Fujimoto, N. Watanabe, and A. Tressaud, Solid State Ionics, 14 (1984) 163.

Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Fluorine-doped tin oxide electrodes for lithium batteries C.W. Kwona, H. Kimb, T. Toupancec, B. Jousseaumec, and G. Campetd a

Samsung Corning Precision Glass, 544 Myungam-ri, Tangjung-myun, Asan-si, Chungnam 336-840, South Korea

b

Samsung Advanced Institute of Technology (SAIT), San 14-1, Nongseo-ri, Giheung-eup, Yongin-si, Gyeonggi-do 449-712, South Korea c

Laboratoire de Chimie Organique et Organométallique, Université Bordeaux I, 351 Cours de la libération, 33405 Talence Cedex, France

d

Institut de Chimie de la Matière Condensée de Bordeaux (ICMCB-CNRS), 87 Avenue du Dr. A. Schweitzer, 33608 Pessac Cedex, France 1. INTRODUCTION The rapid development of modern electronic technology has created a popular demand for portable power sources. Lithium batteries are considered to be the best choice, as they provide high output power and have moderate lifetimes. But their capacity is limited by the electrode materials they contain. In this respect, much effort has been made to improve the performance of electrode materials. Traditional electrochemical working principles are based on the redox potential difference of the electrodes in the course of intercalation/deintercalation reactions [1,2]. They are generally well-crystalline host compounds either with a layered structure such as graphite, LiCoO2 and LiNiO2, or with a tunnel structure such as LiMn2O4. For a long time, only well-crystalline materials have been considered as good electrodes for lithium-ion cells, because their crystal regularity insures the easy diffusion of Li ion in the iono-covalent lattices from the viewpoint of “solid state ionics”. However, nanocrystalline materials are also in the process of being re-evaluated as “nanoscience” advances. Along with the progress of nanoscience, it was found that the electrochemical behavior of nanoparticles is quite different from that of microparticles because the former have a high density of surface defects [3–6]. O’Regan and Gratzel [7] proposed dye-sensitized

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solar cells based on nanostructured TiO2 particles surface-tailored with photoreceptor dyes. Kim and Manthiram [8–10] reported that amorphous manganese oxyiodide Li1.5Na0.5MnO2.85I0.12 showed excellent reversibility involving both the Mn3/Mn4 and Mn2/Mn3 couples, unlike most other manganese oxide systems, which generally show reversibility only for the Mn3/Mn4 couple. Related to this, a few years ago some of us showed, using various examples (SnO2, LiNiO2, TiO2), that the control of crystallite size is a key factor in determining the specific capacity and the cycling efficiency of electrodes [4,11]. Recently, the studies on alloy-based anode materials have also shown the effects of crystallite size on the dimensional stability and capacity retention of the electrode [12]. The disadvantage of the alloy-based anode materials studied is their drastic volume variation during Li insertion/extraction cycles, which leads to fast capacity fading. To overcome this disadvantage, nanocrystalline materials have been intensively investigated. On the basis of these studies, it was concluded that their crack and pulverization can be alleviated in the case of superfine alloy particles due to relatively small absolute volume change [13–15]. Nanocrystalline materials were shown to exhibit an enhanced electrochemical activity, compared with their microcrystalline homologues, only when the first significant electrochemical step is an insertion of Li ions (corresponding to a charge of the Li battery) [4,16]. This insertion begins with the electrochemical grafting of Li ions, promoted by structural defects at or near the nanocrystallite surface. This concept will be explained in detail at the end of this chapter. In this chapter, we start with a brief history of tin oxide as an anode electrode for lithium batteries. A discussion of the doping approach for improving tin oxide, especially fluorine doping, will follow. Finally, we will highlight the surface effects of electrode materials, based on the nanocrystalline materials and on our “electrochemical grafting model”. 2. TIN OXIDE ELECTRODES Tin is one of the oldest metals known to mankind. Its history as pure metal can be traced back to about 1750 B.C., and it was used in bronze (as an alloy with copper) before 2500 B.C. The most common form of tin ore is the oxide cassiterite (SnO2). The study of Li–Sn system in Li-based cells began with the pioneering work of Foster et al. and Wen and Huggins [17,18]. The binary lithium – tin system shows various phases, including Li2Sn5, LiSn, Li7Sn3, Li5Sn2, Li13Sn5, Li7Sn2, and Li22Sn5. Lithium alloyed with tin was electroactive, but not reversible. Therefore, this subject remained forgotten for a long time. However, the current progressive improvement of Li-based electrolytes, driven by the need for enhanced electrochemical performance, now makes it possible to revisit this old subject. Efforts to upgrade anodes led to the discovery that some metal oxides could act as a lithium insertion – extraction anode after

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decomposition/lithiation in the first step [19]. Previously, graphite was commonly used as an anode material for lithium-ion batteries and its capacity approached theoretical value (372 mAh/g) through the development of highly graphitized carbonaceous materials and suitable electrolyte composition [20]. For this reason, various types of alternative anode materials with higher energy density such as metal oxides [21–24], nitride [25–27], phosphide [28–32], intermetallic compounds [33–44], and multiphase/amorphous alloys [45–48] are now being studied. Of these materials, tin-based oxide materials were first suggested as alternatives to graphite and drew considerable attention from a large number of research groups [21,49–64]. These materials showed higher reversible capacities ( 600 mAh/g) than graphite and enhanced capacity retention characteristics relative to lithium binary alloys such as Li–Si and Li–Sn alloys, which showed poor cyclability due to a large volume change during cycling (Fig. 1). Plausible lithiation/delithiation mechanism of tin-based oxide materials had been suggested by Courtney and Dahn [49,50]. Their studies on crystalline tin oxide using in situ XRD and electrochemical characterization techniques

Fig. 1. Voltage profiles for the first 2.5 cycles of the studied materials: (a) Sn, (b) SnO, (c) SnO2, (d) SiSnO3, and (e) Li2SnO3 (from Ref. [49], Fig. 2).

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have shown that the reaction of crystalline tin oxide (SnO/SnO2) with lithium follows a two-step process. The reaction schemes of crystalline tin oxide are as follows: SnOx  2xLi → Sn  xLi2O (amorphous), Sn  yLi  Li2O (amorphous) → LiySn  Li2O (amorphous)

(y  4.4).

In the first step, tin oxide reacts with lithium to form an amorphous Li2O and the Sn metal. These suggested structural changes represent the decomposition of tin oxide and the reduction of Sn(II) or Sn(IV) into metallic Sn, which was also confirmed by in situ XRD [49,50], Raman spectroscopy [51], solid-state NMR [65], and high-resolution TEM analysis [52,66,77] (Fig. 2). Subsequently, the further reaction of lithium with the newly formed Sn embedded in the Li2O matrix leads to the formation of Li–Sn alloys. The reaction responsible for the large reversible capacity in tin-based oxide materials essentially includes an alloying/dealloying reaction between lithium and metallic Sn formed at the first lithiation. The improved reversibility of tin-based oxide materials can be attributed to the finely dispersed Sn metal in the amorphous Li2O matrix, which hinders the Sn atoms from aggregating and growing into large Sn grains. It is believed that the electrodes, in the form of nanocrystalline active materials, can relieve the stresses efficiently due to the small volume change, and remains relatively free from cracking and crumbling as the absolute changes in grain dimensions are still small.

Fig. 2. High-resolution TEM image corresponding to the enlargement of the tin grain; arrow C points the amorphous part surrounding the tin crystal. EDX analysis performed on region C exhibits tin and oxygen; note that the thickness of the amorphous part C surrounding all tin crystallites is always in the range 5–10 nm (from Ref. [77], Fig. 8).

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Courtney and Dahn [49,50] suggested that capacity fading of tin-based oxide materials could be attributed to the aggregation of Sn clusters upon repeated cycles. The aggregation of Sn clusters on cycling, which was identified by TEM analysis for thin-film and powder electrodes, causes a large volume change in the active material, and the degradation of the electrode. This indicates that the Li2O matrix is not completely effective in avoiding Sn atoms from aggregating. In order to avoid Sn aggregation and improve capacity retention, there were attempts to add one or more elements to tin-based oxide materials [67,68]. Amorphous tin composite oxides containing a glass-forming element (B, P, Al, etc.), reported by Idota et al. [21] are a typical example of these attempts. These amorphous solids contain Sn(II) as an electrochemically active center for alloying/dealloying reactions with lithium and oxygen bonded to other glass-forming element networks that limit the size of Sn clusters and also allow fast lithium-ion diffusion and delocalization of the Sn(II) active center. For example, Fuji’s composition, i.e. Sn1.0B0.5P0.5Al0.4M0.103.7 (M  alkaline metals such as K) showed that 90% of the initial reversible capacity was retained after 100 cycles. On the other hand, amorphous tin composite oxide still boses some problems. These materials exhibit a large irreversible capacity and a lower volumetric capacity than crystalline tin oxides. 3. DOPING FOR TIN OXIDES As previously mentioned, tin-based oxide materials have some drawbacks as an anode material in lithium-ion batteries. It is difficult to avoid the large irreversible capacity loss in the first cycle due to the reaction of lithium with oxygen. Also, it is difficult for the tin-based oxide as an anode material to meet both good cyclability and high-energy density requirements. For instance, the addition of glass-forming elements into tin oxide lowered its energy density, although its capacity retention was greatly enhanced. The introduction of heteroatoms into tin oxide might be a plausible solution to these problems, that is, irreversible capacity loss and poor cyclability. It is well known that a dopant can change the characteristics of materials with respect to grain size, surface charge density, and electrical/optical properties. For this reason, several doped SnO2 compounds have been investigated as an anode material for lithium-ion batteries. A number of doped SnO2, MxSn1xO2 (where M  Mo, Si, In, B, and Al) compounds have been synthesized by hydrothermal and other soft chemistry approaches and examined as anode materials in lithium-ion batteries. While using molybdenum as a dopant, beneficial effects on the electrochemical reversibility and cycle performance were reported [69–71]. SnO2 doped with Mo was prepared under hydrothermal conditions and tested in lithium half-cells. Xray diffraction and IR analysis revealed the formation of single-phase products with a rutile-like structure that is maintained upon calcination at 800°C (Fig. 3).

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Fig. 3. XRD patterns for SnO2 and MoySnxO2 samples (from Ref. [70], Fig. 1).

Mo6 ions (0.62 Å), which are in high valence state and smaller relative to Sn4 ions (0.71 Å), change the growth patterns of crystals and are randomly distributed at octahedral positions, thus promoting the formation of cation vacancies. The addition of Mo increases the reversibility of the lithium insertion/de-insertion process as reflected in the simplified differential specific capacity plots obtained, which result in a single, rather symmetric peak in the anodic and cathodic waves (Fig. 4). Furthermore, increasing the Mo content improves retention capacity at the expense of reversible capacity because the transition element acts as an inactive component in the electrochemical process (Fig. 5). The following reasons account for the improved performance of these mixed oxides: (i) a diluent effect of Mo atoms, which facilitates the dispersion of the Sn

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Fig. 4. Differential capacity plots for the LixSnO2 and LiMoySnxO2 cells (from Ref. [70], Fig. 5).

atoms formed during the reduction process; (ii) hindered formation of large clusters and decreased interfacial strains; (iii) restriction of the number of alloying/ dealloying processes through a decreased reversible capacity during the first few cycles; and (iv) an increased chemical diffusion coefficient for lithium, which results partially from the structural disorder caused by the replacement of tin with molybdenum. However, it is difficult to control the Mo/Sn ratio, which has significant effects on electrochemical properties. In order to overcome this drawback, Sn1xMoxO2 mixed oxides of low crystallinity have been synthesized by mechanochemical methods and investigated as electrode materials for lithium batteries. X-ray diffraction, IR spectroscopy, and XPS results suggest that the Modoped samples are solid solutions with a cassiterite-type structure and Mo in the tetravalent oxidation state. Significant amounts of amorphous silica over the range 12–23% (by wt), as determined by energy-dispersive analysis (EDX) and originating from the agate jar and balls of milling apparatus used for the experiments,

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Fig. 5. (a) Changes in specific capacity and (b) charge recovery of (䊉) LiSnO2, (O) LiMo0.02Sn0.97O2, (䊐) LixMo0.14Sn0.78O2, (䉫) LiMo0.17Sn0.74O2, and (䉭) LiMo0.26Sn0.61O2 cells on cycling (from Ref. [70], Fig. 6).

were also detected. The addition of Mo has two favorable effects, namely: (i) it increases the discharge capacity and (ii) it improves capacity retention in cells cycled between 1.0 and 0.0 V (Fig. 6). The formation of a Li–Mo–O oxide-conductive matrix during the electrochemical insertion of lithium may account for this enhanced performance. By this reasoning, the silica present in the samples is assumed to play a minor role on account of the stability of the Si–O bond.

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Fig. 6. Changes in (a) specific capacity, (b) extracted Li/Sn ratio, and (c) charge recovery of Li/SnO2 (䊐 or —), Li/Sn0.96Mo0.04O2 (O or - - -), and Li/Sn0.86Mo0.14O2 (䊉 or …) cells on cycling (from Ref. [71], Fig. 6).

By doping SnO2 with the element that decreases the oxidation state of Sn and the amount of electrochemically active oxygen, one can reduce the irreversible capacity loss at the first cycle. Si-doped SnO2 has been tested for this purpose [72]. Because the Si–O bond is too strong to be broken up by Li in SiO2, Sn1xSixO2 was prepared by the ultrasonic spray pyrolysis method. XRD results revealed that Sn-doped SnO2 consisted of a solid solution of Sn1xSnxO2 and the

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amorphous SiO2 highly dispersed in the former. The major phase obtained in Sndoped SnO2 is rutile-type SnO2, showing a sheet of the lattice of SnO2 upon Si addition, suggesting that traces of silicon dissolved in the SnO2 rutile structure to form a solid solution. The electrochemical results showed low irreversible capacity loss and high reversible capacity ( 900 mA h/g), which is much higher than the theoretical value (783 mA h/g) of SnO2 (Fig. 7). Deviation from theoretical reversible capacity is probably due to the formation of Li6Sn, which is known to exist only at 400°C. Huang et al. stated that the enhanced performance is possibly due to the stabilization of Li6Sn by the presence of Si in either the lattice or as a secondary phase oxide. Antimony-doped SnO2 thin films were also studied as anode materials [73]. Sb-doped SnO2 was prepared by a sol–gel technique and the films obtained were homogenous in composition and morphology and showed a remarkable decrease in grain size, (a few nanometers) and resistivity. Sb-doped SnO2 (5%) showed enhanced capacity retention characteristics relative to undoped SnO2 (Fig. 8), which was due to the decrease in grain size and the increase in the electrical conductivity of active materials by doping with Sb. Effects of doping with In and B were also investigated [74]. In- and B-doped SnO2 nanoparticle synthesized by a hydrothermal method exhibited their best cycle performance between 1.0 and 0.0 V. Undoped SnO2 and particularly low Bdoped samples exhibited the best electrochemical characteristics, as reflected in an increased specific capacity and improved cycling properties. These cells retained 75% of their initial capacity after 30 cycles. By contrast, increased B content or

Fig. 7. The discharge profiles at a current rate of 0.02 mA; after three cycles at a current rate of 0.1 mA (from Ref. [72], Fig. 5).

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Fig. 8. Capacity (mAh/g) versus cycle number of cells with antimony–tin oxide thin films as anodes (from Ref. [73], Fig. 3).

Fig. 9. Delivered anode capacity from galvanostatic cycling of (䉫) Li/SnO2, (䊉) Li/[SnO2/B (2%)], (O) Li/[SnO2/B (10%)], (䊏) Li/[SnO2/In (1.3%)], (䊐) Li/[SnO2/In (5%)], (x) Li/In2O3, (䉬) Li/B2O3, and (·) Li/(graphite–acetylene black–PTFE) cells (from Ref. [74], Fig. 6).

doping with In caused a significant drop in the capacity; in fact, cells hardly retained 30% of their initial capacity after only a few cycles (Fig. 9). The adverse effects of the doping elements are ascribed to the poor electrochemical performance of bulk BO and InO rather than to other factors such as particle size, shape, or crystallinity. The metal-citrate method was applied to the preparation of microcrystalline tin dioxide and Al-doped materials [75]. The thermal decomposition of the

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citrate precursors leads to the formation of rutile-type SnO2. For Al-doped samples, 27Al-MAS-NMR signals at ca. 0 ppm and 119Sn-Mössbauer signals at 0 mm/s isomer shift and 0.3 mm/s quadruple splitting evidence that Al3 ions occupy octahedral sites isomorphic to Sn4. XAS results of the Al-doped samples showed that the presence of Al atoms in the structure does not modify the relative positions of the different anti-bonding states in SnO2. XRD line-broadening analysis evidenced a large microstrain content in the ex-citrate products, which are released by successive thermal treatments at 450°C. Al-containing exbiscitrate oxide (10%) exhibited the best electrochemical performance as an anode material in lithium-ion batteries (Fig. 10).

Fig. 10. Cycling behavior of lithium cells using the excitrate powdered solids as active cathode material: (a) cell voltage vs. capacity for sample Al21 and (b) capacity vs. cycle number for selected samples (from Ref. [75], Fig. 8).

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4. FLUORINE-DOPED TIN OXIDES Besides the above-mentioned studies, there are several significant works, such as the use of tin oxide electrodes for electrochromic device [76], cycling properties of nanocrystalline tin oxides depending on synthetic conditions [77,78] and doping with Cl, Sb, Mo, or F [79–82]. Among these works, fluorine doping is known to be the most efficient for enhancing the electrochemical properties of tin oxide, achieving the highest electrical conductivity up to ⬃5  103/Ω cm at an optimum atomic F/Sn ratio of ⬃3%; for the sake of clarity, let us recall that for each fluorine ion, an electron is introduced into the conduction band according to xF  xe  SnO2 → (Sn4  xe)FxO2-x  xO2 [83]. The enhancement of conductivity with F-doping may increase the reversible electrochemical capacity of this kind of material. Inspired by this idea, our group recently attempted to obtain nanocrystalline F-doped SnO2 via a single molecular precursor by sol–gel process. Methods to stabilize fluoroalcoxytin complexes, which are ideal precursors for preparing highly conductive F-doped SnO2 nanocrystalline powders [84,85] have been reported. The precursors are types of mixed-valence fluorotin alkoxides. The nanocrystalline F-doped SnO2 powder was prepared by the method reported in Ref. [90]. The simplified procedure is illustrated by the following scheme [91]: R O SnF 2

F

+ F

II

IV

Sn

Sn

R O

Hydrolysis

OR

Thermal treatment

Nanocrystalline Conductive F-doped SnO2

Sn(OR)4 O R R = -C(CH3 )3

The thermolysis of xerosol was carried out first by drying the powder at 50°C, and then calcining it at 550°C in air for 15 min [91]. Commercial SnO2 powder was also compared as a reference. Fig. 11 shows the X-ray diffraction patterns of the dried xerosol precursor, the nanocrystalline F-doped SnO2 and undoped SnO2 powders. For the dried xerosol precursor, no distinguishable peak appears, indicating an amorphous state or the absence of long-range order. After heat treatment at 550°C in air, as expected, XRD peaks appear to be broader compared with the well-crystalline Aldrich sample. All the XRD reflections are properly indexed with tetragonal

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Fig. 11. The X-ray diffraction patterns of the dried xerosol precursor, the nanocrystalline F-doped and well-crystalline undoped SnO2 powders (from Ref. [91], Fig. 1).

symmetry and the cell parameters obtained from the least-squares fitting analysis are listed in Table 1. It is evident that the cell parameters become smaller for the F-doped nanocrystalline sample. The size of crystallites was estimated to be 7 nm using Scherrer’s relation t  0.9λ / (Β cos θ ), where λ is the X-ray wavelength, θ the Bragg angle and B the angular full-width halfmaximum of the chosen (hkl) reflection in radians. For the doped sample, the F/Sn ratio of ⬃0.14, based on the result of elemental analysis, leads to a diminished resistivity (ρs⬃0.7 Ω cm). On the contrary, for the undoped SnO2, the resistivity is higher than 105 Ω cm. At this point, it should be noted that the above-reported F/Sn ratio of 0.14 is about five times higher than the maximum value that can be incorporated into the lattice [83]. It might be that some amorphous (as not detected by XRD) secondary phases exist at the grain boundary such as Sn4F4 [83]. The oxidation state of tin in the nanocrystalline F-doped SnO2 was determined by Mössbauer spectroscopy at room temperature. The 119Sn-Mössbauer spectrum of the F-doped nanocrystalline SnO2 and its Lorentzian fits are presented in Fig. 12 [91]. These spectra show the existence of the Sn4 state only, although this material was prepared by heating a mixed valence precursor [85]. This finding is consistent with the observed low resistivity (0.7 Ω cm), which implies that the F-doped SnO2 may generate mobile electrons in the conduction band, instead of trapped electrons leading to Sn2 formation (Sn4  2e → Sn2).

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Table 1 Elemental analysis, electrical resistivity, crystallite size and cell parameter data of the Aldrich SnO2, and the nanocrystalline F-doped SnO2 Undoped SnO2

F-doped SnO2

Elemental analysis (mole ratio) F/Sn C/Sn

– –

0.14

0.04

Electrical resistivity ρs (Ω cm)

 105

0.7

Crystallite size l (nm)

 200

⬃7

Cell parameters a (Å) c (Å)

4.73(8) 3.18(7)

4.69(3) 3.17(6)

Fig. 12. The 119Sn-Mössbauer spectrum of the F-doped nanocrystalline SnO2 and its Lorentzian fits (from Ref. [91], Fig. 2).

Fig. 13 shows the first discharge curves for Li/SnO2 cells. A distinctive plateau at about 1.0 V can be observed for the undoped SnO2 sample, which is due to the initial formation of Sn. In contrast, the plateau is less obvious for the doped nanocrystalline sample. For the voltage region above 1.5 V, the F-doped SnO2 shows a capacity of ⬃30 mAh/g, which is about six times larger than that of the microcrystalline undoped sample ( 5 mAh/g for the first cycle). However, let us recall that, in the latter case, a gradual textural change, namely a reduction in the size of the crystallites was observed upon Li insertion, because the insidecrystallite structure is not adapted for lithium intercalation; such a reduction in the crystallite size leads to an increase in the total number of Sn atoms lying near the grain surface and are acceptable of sustaining a reversible electrochemical

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Fig. 13. The discharge curves of Li/SnO2 and Li/F-doped nanocrystalline SnO2 cells (from Ref. [91], Fig. 3).

reduction as occuring in nanosized materials [11]. A similar phenomenon was also observed in the case of nanocrystalline SnO2 prepared by the hydrolysis of SnCl4 and heating at various temperatures [51]. As reported previously by our group, in general, a higher insertion voltage can be observed at the same insertion level for an n-type semiconducting electrode with smaller particles; moreover, the potential decreases more smoothly as lithium insertion progresses in these types of electrodes composed of nanocrystalline materials such as LixSnO2, LixTiO2 and LixWO3 [45,77–90]. It suggests that a regular and smooth change of Fermi energy in the electrode without any significant structural change, owing to the existence of many sub-band gap states between the valence and conduction band, may be due to surface defects and surface dangling bonds [90]. It is worth noting here that the first discharge capacity (⬃1800 mAh/g) of the doped sample is larger than the theoretical one (⬃1490 mAh/g), probably due to the influence of carbon black (25 wt%). The first two cycling curves of the cells are shown in Fig. 14. The undoped sample shows the reversible capacity of ⬃600 mAh/g, which corresponds well to the literature values [49,51,69,77,78]. For the F-doped nanocrystalline SnO2, the reversible capacity is close to the expected maximum one (⬃800 mAh/g) [91]. Thus, it seems that highly conductive nanoparticles may provide a way of forming the Li–Sn alloy more effectively, as they reach a ratio of 22:5 (Li22Sn5) [92]. On the whole, it can be rationalized that a nanocrystalline and conductive matrix facilitates the diffusion of Li and the dispersion of Sn atoms in the electrode and that the formation of an alloy around the Sn atoms occurs more

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Fig. 14. The first two cycle curves of Li/SnO2 and Li/F-doped nanocrystalline SnO2 cells (from Ref. [91], Fig. 4).

intactly. Even though further systematic investigations should be conducted on the long-term cycling behavior and on the microstructural change induced by dopage of F, the F-doped nanocrystalline SnO2 reported here is expected to be a promising anode material for Li-ion batteries. 5. ELECTROCHEMICAL GRAFTING MODEL As mentioned earlier, the surface effects of nanocrystalline electrodes exhibit a unique electrochemical behavior as distinct from crystalline ones. It is worth mentioning a few key features of the nanocrystalline electrodes: (i) They work better when the insertion of Li is the first electrochemical step. (ii) At the first Li insertion step, a higher insertion voltage can be observed at the same insertion level compared with well-crystalline homologues. (iii) After the first Li insertion step, the subsequent cycles are reversible. These trends agree well with our previous reports on n-type semiconducting electrodes composed of nanoparticles such as LixSnO2, LixTiO2, and LixWO3 [86–90]. Based on the above phenomenological facts, we here suggest the “electrochemical grafting” model. Fig. 15 depicts a schematic band model for this electrochemical grafting process. In this model, the nanocrystalline materials have many subband-gap states between the conduction and valence bands owing

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Fig. 15. A schematic band model for the electrochemical grafting process (from Ref. [93]).

to their surface defects and surface dangling bonds. For example, the metal subband represents deep subband-gap energy states arising from cation defects adjacent to an anion vacancy and slightly lower than the conduction band of metal. The first electrochemical Li insertion process fills the subband-gap states, and consequently gives a more smooth discharge curve compared with the well-crystalline homologue. It results in a regular and smooth change of Fermi energy in the electrode without undergoing any significant structural change. Therefore, the nanocrystalline materials can endure better for the structural phase transition. Once the subband-gap states are filled, they are cured, and the subsequent electrochemical processes are reversible. In this regard, we named this step electrochemical grafting. In this way, nanocrystalline materials can be more efficient than their well-crystalline homologes. We believe that this model can deliver useful insights for developing more effective electrodes. 6. CONCLUSION The electrochemistry of various modified tin oxide-based materials was reviewed on the basis of their different behavior when compared with the pristine wellcrystalline stannic oxide. The electrochemistry of nanocrystalline materials differs from that of traditional well-crystalline ones due to their significant surface effects, and the electrochemical properties of doped-SnO2 show improved performance mainly due to increased conductivity. Especially, it is emphasized that the nanocrystalline F-doping approach can be profitably used to overcome conductivity problems in lithium insertion electrodes by facilitating alloy formation through a judicious molecular precursor route. The above-presented

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Synthesis of fluorinated cathodes and fluoride electrolytes for lithium-ion batteries Susumu Yonezawa and Masayuki Takashima Department of Materials Science and Engineering, Faculty of Engineering, University of Fukui, 3-9-1 Bunkyo, Fukui-shi 910-8507, Japan 1. GENERAL INTRODUCTION Fluorine chemistry plays an important role in the development of materials for lithium-ion batteries. In this chapter, the incorporation of fluorine into lithiumcontaining transition metal oxides for obtaining a new cathode-active material, and the new preparation method of the electrolyte salt LiPF6, are described. 2. CATHODE-ACTIVE -MATERIALS 2.1. Introduction

The oxides, sulfides and oxide chlorides of transition metals were investigated for as their use cathode-active materials of lithium secondary batteries during the 1980s. Since then, lithium-containing transition metal oxides have been investigated as cathode-active materials whereas carbon materials have been used as anode-active materials. The “lithium-ion battery” appeared in the market in 1991. LiCoO2 was the first cathode-active material used for a practical cell. LiNiO2, having the same crystal structure as LiCoO2, and spinel LiMn2O4, which has cost merit, have also been studied as cathode-active materials [1,2]. Currently, “lithium-ion batteries” are widely used as electric sources of small electronic apparatus such as mobile phones and notebook computers. There are some problems, however, in preparing lithium-ion batteries large in size and with high-power density that can be used as an electric source in electric vehicles and the battery for load leveling. The electric conductivities of LiCoO2, LiNiO2, LiMn2O4 and their derivatives are so low that they have to be used together with acetylene black an electroconductive materials. The crystal lattices of these compounds are destroyed during the charge/discharge cycles at which the insertion/deinsertion of the lithium ion takes place, and oxidative degradation of the electrolyte solution occurs during

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the charge process. These phenomena greatly affect the lifetime of the battery. The dissolution of Mn2 shortens the lifetime of batteries that have a LiMn2O4 cathode. The charge/discharge reaction occurs at the three-phase interface consisting of an active material, a carbon material and an electrolyte solution, that is, the reaction site is the surface of the active material. Therefore, the surface structure of the active material greatly influences the charge/discharge properties. As shown in Fig. 1, it is expected that several improvements in electrochemical properties are achieved when the surface of the cathode-active material is treated by fluorine: (1) Since the fluoride ion has a greater interaction with the Li ion than the oxide ion, it is expected that the M–F bond at the surface of the cathode-active material promotes the Li ion transfer at the interface of the active material and the electrolyte

Fig. 1. Illustration of the influence of surface fluorination of metal oxides on their electrochemical properties. (a) Li ion transfer across the interface of the active material and the electrolyte. (b) Protecting the surface from HF attacking. (c) Elimination of H2O and –OH group on the active material surface.

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solution in which the Li ion is solvated, as illustrated in Fig. 1(a). (2) LiPF6 is used as an electrolyte salt. LiPF6 is, however, easily decomposed to HF due to hydrolysis with a trace of water in the cell, as shown in Fig. 1(b). HF will attack the oxideactive material to reproduce water. This recycling process continues until all active sites of the cathode-active material are attacked. A fluorinated surface can prevent this kind of degradation because the M–F bond will no longer be attacked by HF. (3) There exist adsorbed water molecules and hydroxyl groups at the surface of the oxides. Fluorination by elemental fluorine can remove these impurities at the surface of the active materials, as shown in Fig. 1(c). Several approaches in preparing new cathode-active materials by fluorination treatment will be described in the following sections. 2.2. Introduction of fluorine into the bulk of oxides

The cathode-active materials for lithium secondary batteries, such as LiCoO2, LiNiO2 and LiMn2O4, are usually prepared by heating the mixture of transition metal oxide powder (Co3O4, NiO, etc.) and the lithium source such as LiNO3 and LiOH. Several trials to prepare LiMxOyFz, in which some oxide ions are substituted by fluoride ions have been performed by adding LiF to the starting mixture e.g. LiNiO2 by Kubo et al. [3] and LiCoO2 by Yonezawa et al. [4]. It was reported in both papers that cycleability was improved by adding LiF to the starting mixture. On the other hand, another paper reported that the addition of alkali metal fluorides such as LiF to the cathode mixture helped to improve battery performance [5]. It is not clear whether a fluoride ion is substituted or not for an oxide ion in the crystal lattice of cathode material. In the case of LiCoO2, it was reported that the addition of LiF influenced the crystal orientation and the morphology of the sample particles [4]. Regarding LiNiO2, the substitution of an oxide ion with a fluoride ion in the lattice was shown by the Rietveld analysis of XRD results. However, it is very difficult to distinguish a fluoride ion from an oxide ion by using XRD data. Therefore, there are still some doubts as to whether the fluoride ion exists in the lattice of LiCoO2 and LiNiO2 or not. A preparation method that involved adding lanthanide trifluorides or fluoro-organic compounds to the starting mixture was proposed for insertion of the fluorine into the bulk of the compounds [6,7]. In addition to the oxides, Li2MPO4F (M  Fe, Co, etc.) was studied as a fluorine-containing cathode-active material [8]. Li2MPO4F was prepared by heating LiMPO4 with LiF at around 800°C. When M  Co, the dissolution of cobalt ion into the electrolyte solution was prevented. Li2CoPO4F was stable even at 60°C. Li2MPO4F may be one of the materials expected to be used as a cathodeactive material for lithium-ion batteries. On the other hand, LiMn2O4 sometimes contains an oxygen defect in its crystal lattice. Usually, lithium manganese oxide with oxygen defects, LixMn2O4-δ , is prepared by changing the Li content in the starting mixture. Normally, an excess amount of a Li compound is added. It was reported that a fluoride ion can be inserted into the crystal structure by a reaction

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between LixMn2O4-δ and F2, ClF3 or NF3 [9]. The schematic illustration of the tightly sealed fluorination line used in this study is shown in Fig. 2. All the parts, except the reaction vessel are made of SUS316L. The reaction vessel, made of Ni, was used at a reaction temperature higher than 200°C. The introduction of an excess amount of fluorine into the crystal lattice of the cathode-active material reduced the discharge capacity (mAh g1) and the average discharge potential, possibly due to the formation of metal fluoride with a high-potential barrier against electron conduction. Surface fluorination is therefore one of the solutions to prevent the formation of metal fluoride. There are several methods of introducing fluorine to the surface region of active material. In order to obtain pure and

Fig. 2. Schematic illustration of fluorination line: 1, pirani gauge; 2, Bourdon tube gauge; 3,4, activated alumina and soda lime-filled column; 5, oil rotary pump; 6, reaction vessel; 7, electric furnace; 8,11, fluorination reagent gas cylinder; 9,10, buffer; 12, argon gas cylinder.

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dry samples, a solid–gas reaction was carried out in the present study: NF3, ClF3, F2 or other gaseous fluorinating reagents were used. Among these, NF3 is so stable that a glass can be used as a reactor at a temperature lower than 200°C. However, the reaction between NF3 and oxides proceeds even at room temperature. The surfaces of the oxides act as some kind of catalyst to activate NF3. 2.3. Surface fluorination

The electrochemical reaction takes place at the interface between the electrolyte solution and the electrode. The introduction of fluorine to the surface of the cathode-active material provides the following advantages: (1) the solvation/desolvation of the lithium ion proceeds smoothly; (2) the electric conductivity at the surface increases; (3) the oxidative degradation of the electrolyte solution is lowered; (4) the crystal structure at the sample surface is stable during the charge/discharge cycles; and (5) the surface is protected from HF attack. HF is generated from the hydrolysis of LiPF6, which is widely used as an electrolyte salt and is very sensitive to water. The surface fluorination of LiCoO2, LiNiO2, LiNixCo1-xO2 and LiMn2O4 and their electrochemical properties were reported [10]. Fig. 3 shows the cyclic voltammograms (CV) of LiCoO2 whose surface is treated with NF3 (NF3/LiCoO2) at room temperature (rt), 100 and 200°C. Surface fluorination of the sample at a temperature below 100°C resulted in sharper peaks in the CV and larger peak currents, compared with those for untreated samples. On the other hand, surface fluorination of the sample above 200°C resulted in broader peaks CV and reduced the peak currents. XPS spectra of NF3/LiCoO2 (LiCoO2 treated with NF3 gas) and ClF3/LiCoO2 (LiCoO2 treated with ClF3 gas) are shown in Fig. 4. From the results of XPS measurements, it was found that fluorination proceeded to the inner part of the sample particle at a temperature  200°C, and the surface was covered with a continuous film that consists mainly of LiF with high resistivity. These results are reflected in the shape of the CV. In case of the surface fluorination of LiCoO2 below 100°C, it was suggested that the sample surface was partly covered by fluorine and that this type of partly fluorinated surface contributed to the improvement of the electrochemical performances of the cathode. Evidently, the characteristics of the surface-fluorinated sample strongly depend on the moisture in atmosphere. It is therefore important that the fluorination line is sealed and completely cleaned. It is necessary to keep a fluorinated sample as dry as possible. The investigation described above, suggests that the surface fluorination of the active material improved battery performance by promoting Li ion transfer. In order to improve the electrochemical process, not only the ion transfer but also the electron transfer by a carbon material should be maintained at a sufficient level. There is another approach to obtain new cathode-active materials, the surfaces of which are modified by nanothickness carbon coating and fluorination. It was reported in a previous paper that the modification of LiMOx surface by coating with nanothickness carbon and fluorine improves its performance as a cathode-active

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Fig. 3. Cyclic voltammograms (0.02 mV s1, 25°C, first cycle) of LiCoO2 (a) and NF3/LiCoO2 (treated at room temperature (b), 100°C (c) and 200°C (d)) in 1.0 mol dm3 LiClO4/ PC  DME (50:50, v/v).

material [11]. The discharge capacities of LiMn2O4 with a fluorine–carbon nanocomposite surface were increased by 5–10% at the first charge/discharge cycle. The reversibility of the electrochemical process was thus improved by coating with nanothickness carbon. It is expected that coating the active material with a carbon material contributes to the inhibition of the oxidative decomposition of the electrolyte solution [12,13] and the dissolution of Mn2 from the active material [14]. Oxidative degradation of the electrolyte solution decreases the discharge capacity and lowers the charge/discharge efficiency. Surface fluorination by NF3 (F–LiMn2O4) was carried out before (nC–F–LiMn2O4) or after carbon coatings (F–nC–LiMn2O4), where n corresponds

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Fig. 4. XPS spectra of F1s level in LiCoO2 treated with NF3 and ClF3 at 1.3 kPa before (a) and after argon sputtering for 30 min (b). ((1), untreated; (2), NF3, room temperature; (3), NF3, 100°C; (4), NF3, 200°C; (5), NF3, 300°C; (6), ClF3, room temperature; (7), ClF3, 100°C.).

to the thickness of the coated carbon layer [15]. Fig. 5 shows the discharge capacities with cycle number for F–LiMn2O4, 30C–F–LiMn2O4 and F–30C–LiMn2O4. The maximum values of the discharge capacities were 104, 102 and 100 mAh g1 for F–LiMn2O4, F–30C–LiMn2O4 and 30C–F–LiMn2O4, respectively. The decrease in discharge capacities against maximum value at the 50th cycle amounted to 19, 3 and 8% for F–LiMn2O4, 30C–F–LiMn2O4 and F–30C– LiMn2O4, respectively. There is a synergistic effect of the nanothickness carbon coating and surface fluorination on the charge/discharge capacity and the cycle ability of LiMn2O4. The difference in behavior between 30C–F–LiMn2O4 and F–30C–LiMn2O4 reveals that the arrangement of the nanothickness carbon film and the fluorine or fluoride ion has an influence on the electrochemical properties. Fig. 6 shows the charge/discharge curves of LiMn2O4, F–LiMn2O4 and 30C–F–LiMn2O4. The charge was continued at 4.3 V (vs. Li/Li) after the electrode potential reached 4.3 V until charging time reached 18 h. The charging times until the potential reached 4.3 V were 13.1, 13.5, and 13.9 h for 30C–F–LiMn2O4, F-LiMn2O4 and LiMn2O4, respectively. The current decay for 30C–F–LiMn2O4 after the potential reached 4.3 V was the fastest among them, and the average

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Fig. 5. Change in discharge capacities of F–LiMn2O4 (䊉), 30C–F–LiMn2O4 (䉱), and F–30C–LiMn2O4 (䊏) along cycle number. Charge and discharge was carried out at a constant current corresponding to the rate of 0.3 C. Cut off potential was set at 4.5 V.

current during the potentiostatic charge at 4.3 V was the smallest (1.7 mA g1) among them. The average current for F–LiMn2O4 (2.1 mA g1) was slightly smaller than that for LiMn2O4 (2.2 mA g1). The order of the average currents, 30C–F–LiMn2O4 F– LiMn2O4 LiMn2O4, corresponds to that of the decrease in discharge capacity against the maximum value through 50 cycles. It seems that the oxidative degradation of the electrolyte solution occurs above 4.2 V in this case. The quick decay of the current and the low average current during the potentiostatic charge at 4.3 V reveals that the electrolyte solution was protected from the oxidative degradation. Surface fluorination is effective in improving the charge/discharge cycleability. However, fluorine at the surface of LiMn2O4 may be gradually removed through the charge/discharge cycles when it is in direct contact with the electrolyte solution under polarization. Because fluorine in 30C–F–LiMn2O4 is covered with nanothickness carbon, it can be preserved at the surface of LiMn2O4, where the electrochemical reaction takes place through charge/discharge cycles. Charge/discharge capacity and cycleability of LiMn2O4 as a cathode-active material of lithium secondary batteries will be improved by optimizing the arrangement of the nanothickness carbon film and the surface fluorine of LiMn2O4. 2.4. Fluorides

It has been reported that some fluorides such as LiCaCoF6 have a higher potential to be used as cathode-active materials, having a discharge potential of about 5.8 V (vs. Li/Li) [16]. This was proven by molecular orbital calculation

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Fig. 6. Charge/discharge curves in first cycle (A) and the magnified one in the region from 11 to 23 h (B). LiMn2O4 ((a), —), F–LiMn2O4 ((b), —), and 30C–F–LiMn2O4 ((c), —).

with the first principle. Since there is no electrolyte that can be used at around 5.8 V now, it is expected that some new materials will be developed to test this type of fluoride in the future. 3. NEW PREPARATION METHOD OF THE ELECTROLYTE SALT LiPF6 3.1. Introduction

LiMFx (M  B, P, As) has been studied intensively as an electrolyte salt for lithium-ion batteries. Of these, the most popular electrolyte salt used in the lithium-ion battery is lithium hexafluorophosphate (LiPF6), which has good solubility in various solvents such as PC (propylene carbonate) [17]. A lot of efforts

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have been made to obtain highly pure LiPF6 because even a trace of water can cause the battery performance to deteriorate. Generally, liquid anhydrous hydrogen fluoride (L-AHF) is used as a medium for the preparation reaction of LiPF6 between LiF and PF5 [18,19]. Acetonitrile can also be used as a medium for the preparation reaction. In both cases, LiPF6 must be purified by recrystallization in a dry organic media to remove H2O and HF remaining in the products. Lithium oxyfluorophosphate (LiPOxFy) is also produced as a by-product, which is partially dissolved in an HF solution [20]. The removal of a trace of water in AHF has been attempted by using F2 gas [21]. The process of reducing the amount of HF remaining in the product consists of forming an adduct, Li(CH3CN)4PF6, with highly dried acetonitrile. It is, however, very difficult to remove it completely. 3.2. Preparation of LiPF6 by direct method

It has been reported that LiPF6 can be prepared by reacting the mixture of LiF and P (red phosphorus) with elemental fluorine (F2 direct method) [22]. Fluorine gas (purity, 99.4–99.7%; supplied by Daikin Industries, Ltd), LiF and red phosphorus (purity, 99.9%) were used as starting reagents. LiPF6 has been successfully prepared by the reaction between the mixture of an equi-molar ratio of LiF and P and F2 gas at 300°C under an F2 pressure of 0.4 MPa within 5 min. The amounts of LiF, P, and F2 were 1.3  103, 1.3  103 and 3.9 103 mol, respectively, in the present study. The volume of the reactor, made of nickel, was 1.57  105 m3. It was found that the stepwise introduction of F2 into the reactor was effective for obtaining LiPF6 at a high yield. F2 of p/n (n, steps) was introduced for each step, where p corresponds to its stoichiometric amount of F2 required to complete the reaction. Fig. 7 shows the results of the XRD pattern of the sample prepared at 300°C at various steps, n. The peaks corresponding to LiF appeared at 39° and 45° in Fig. 7(a). This means that the reaction between LiF and P in F2 gas is not complete under this condition. An increase in the steps, n, diminished these peaks, and finally, no peak corresponding to LiF was detected for n  4 (Fig. 7(d)). The stepwise introduction of fluorine gas may be much more important for carrying out the reaction efficiently and homogeneously than the other factors such as temperature, F2 pressure and reaction time, in this case. There may be a certain equibrium giving rise to the observed phenomenon during the reaction. It is known that LiPF6 dissociates into LiF and PF5 at a temperature higher than 220oC. However, LiPF6 was efficiently prepared at 300°C in the present study. F2 may protect LiPF6 and PF5 from dissociation and hydrolysis, or play an important role in promoting the reaction LiF  PF5  LiPF6. Essentially, the product obtained by this method was the same as that prepared by the liquid AHF (L-AHF) method. But the moisture sensitivity of LiPF6 obtained in the present study was different from that of the conventional sample. LiPF6 prepared by the L-AHF method was more easily decomposed than that prepared by the F2 direct

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Fig. 7. XRD patterns of the products obtained at various steps, n. n is the number of steps in which (3.9  103)/n mol F2 was introduced into the reactor (1.57  105 m3) for each step. The reaction temperature and time were 573 K and 5 min, respectively. ((a) n  1; (b) n  2; (c) n  3; (d) n  4; (e) the product prepared by L-AHF) 䊊 and 䊉: the peaks due to LiPF6 and LiF, respectively.

method. A trace of HF adsorbed at the surface of the product may be responsible for this phenomenon. LiPF6 prepared by the F2 direct method has the advantage of stability against hydrolysis, compared with that prepared by the L-AHF method. Considering the results of 31P- and 19F-NMR measurements and the conductivity measurement of the PC solution of the products [22], LiPF6 prepared by the F2 direct method is pure, and can be used as an electrolyte salt for lithiumion batteries as produced. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8]

T. Ohzuku and A. Ueda, J. Electrochem. Soc., 141 (1994) 2972. J.M. Tarascon and D. Guyomard, Electrochim. Acta, 38 (1993) 1221. K. Kubo, M. Fujiwara, S.Yamada, S. Arai, and M. Kanda, J. Power Sources, 68 (1997) 553. S. Yonezawa, T. Okayama, H. Tsuda, and M. Takashima, J. Fluorine Chem., 87 (1998) 141. Published patent application, Japan, H07–220758. Published patent application, Japan, 2000–353523. Published patent application, Japan, 2000–353524. S. Okada, S. Sawa, M. Egashira, J.-i. Yamaki, M. Tabuchi, H. Kageyama, T. Konishi, and A. Yoshino, J. Power Sources, 97–98 (2001) 430. [9] T. Tanida, S. Yonezawa, and M. Takashima, Abstr. 43rd Battery Symp., Japan, vol. 1A18, 2002. [10] S. Yonezawa, M. Ohe, T. Tanida, M. Takashima, Abstr. of 13th European Symp. on Fluorine Chemistry, vol. B25, 2001. [11] M. Takashima, S. Yonezawa, and M. Ozawa, Mol. Crys. Liq. Crys., 388 (2002) [567]153.

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[12] E. Endo, T. Yasuda, A. Kita, K. Yamaura, and K. Sekai, J. Electrochem. Soc., 147 (2000) 1291. [13] T. Eriksson, A.M. Andersson, A.G. Bishop, C. Gejke, T. Gustafsson, and J.O. Thomas, J. Electrochem. Soc., 149 (2002) A69. [14] D.H. Jang and S.M. Oh, J. Electrochem. Soc., 144 (1997) 3342. [15] S. Yonezawa, M. Ozawa, and M. Takashima, TANSO, 205 (2002) 260. [16] Y. Koyama, I. Tanaka, and H. Adachi, J. Electrochem. Soc., 147 (2000) 3633. [17] N. Katayama, T. Kawamura, Y. Baba, and J. Yamaki, J. Power Sources, 109 (2002) 321. [18] W.N. Smith and J.E. Pa, U.S. Patent, 3, 607, 020, 1971. [19] D.J. Salmon and D.W. Barnette, U.S. Patent, 5, 378, 445, 1993. [20] R.A. Wiesboeck, U.S. Patent, 3, 654, 330, 1972. [21] D. Na, B. Woo, S. Park, and J. Lee, U.S. Patent, 6, 387, 340, 2002. [22] J.H. Kim, S. Yonezawa, and M. Takashima, Chem. Lett., 33, 2004, 884–885.

Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 7

Physicochemical properties of fluorine-containing electrolytes for lithium batteries D. Lemordant, F. Blanchard, G. Bosser, M. Caillon-Caravanier, B. Carré, A. Chagnes, B. Montigny, and R. Naejus Laboratoire Chimie-physique des Interfaces et des Milieux Electrolytiques, EA 2098, Faculté des Sciences, Parc de Grandmont, Université F. Rabelais, 37200 Tours, France 1. INTRODUCTION: FLUORINE AND LITHIUM BATTERIES 1.1. Lithium as metal anode

Lithium is the lightest metal of all elements of the periodic classification: its density is only 0.534 g/cm3 at 293 K. Li has the lowest standard oxidation potential: E o ⬇3 V vs. the standard hydrogen electrode (SHE). The standard Gibbs energy for the reaction Li → Li  e is 䉭Go   FEo ⬇ 290 kJ/mol, which corresponds to a specific energy density of 41 kJ/g. For this reason, lithium is by far the most energy-dense battery material. As lithium is a very reactive metal, it is not stable towards most of the solvent. Protic solvents like water cannot be used; only a restricted number of dipolar aprotic solvents are convenient for use with lithium. These solvents are mostly alkylcarbonates and lactone derivatives (cyclic or acyclic), but some aliphatic esters or ethers may also be employed. Primary (non-rechargeable) batteries make use of metallic lithium as anode and metal oxides like MnO2 and V2O5 as cathode. Commercial secondary lithium batteries (rechargeable) make use of carbon (graphite, coke, etc.) as anode material instead of metallic lithium. The main reason for this is that lithium cannot be deposited without dendrites formation. The growth of dendrites during the subsequent charge – discharge cycles will inevitably lead to short circuit or “dead”

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lithium deposition, i.e. lithium that is not in contact with current collectors. Moreover, the reduction of the solvent on the lithium surface will lead to the formation of a passive film that is partially destroyed during the discharge cycle and will have to be reformed during the following charge that corresponds to lithium deposition. This phenomenon leads to electrolyte and anode material consumption drastically reducing the life cycle of the battery. Using graphite as anode material, an intercalation compound will be formed in the host structure: LixC6 (0  x  1). The intercalation proceeds by successive stages at low potentials (from 0.25 to 0.08 V vs. Li/Li). When fully charged, the stoichiometry of the intercalation compound is LiC6 (x  1), which has a molecular weight of 79 g/mol. The presence of the host matrix will reduce the faradic capacity of the electrode from 3860 mAh/g for metallic lithium to 360 mAh/g for LiC6. Nevertheless, as the intercalation potential is not well above that of the reduction of lithium ions, the loss in energy density will be limited. 1.2. Li-ion cells

In Li-ion cells, the lithium anode is replaced by an alternative source of Li. To date, the most energy-dense material is lithiated carbon (Li(C)), as the potential of the Li/Li(C) is only slightly less negative than Li. Different types of carbon may be employed, such as pitch carbon and hard and soft carbon (more or less graphitized carbon). A typical Li-ion cell, as displayed on the scheme reported in Fig. 1, associates a carbon anode, a lithiated cobalt oxide cathode (LiCoO2) and an electrolyte containing a lithium salt. The reactions occurring at the electrodes during the charge are given by xLi xe  C → Lix(C)

electrons

x Li+ + e- + Li1-xCoO2 → LiCoO2

LixC→ x Li+ + e- + (C)

Li+

Fig. 1. A typical Li-ion cell (charge state) composed of a carbon anode, a partially lithiated cobalt oxide cathode (Li1-xCoO2) and an electrolyte. Reactions are quoted for the discharge process.

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and LiCoO2 → xLi  xe  Li1xCoO2

(x  0–0.9)

As the potential of the LiCoO2/Li1xCoO2 system is near 1.2 V vs. ESH, the cell potentials at the discharge and charge states are, respectively, 3.6 and 4.2 V. If practically, only 0.6 Li per LiCoO2 are removed from the oxide matrix; the corresponding faradic capacity of this cathode is of the order of 160 mAh/g and the energy density of the order of 1.2 kJ/g, which is far less than that of Li or even LiC6 (3.55 kJ/g). 1.3. Electrolytes

Li batteries require electrolytes that contain Li ions such as Li-ionconducting liquids or Li-ion-conducting polymers. Liquid electrolytes are easily obtained by dissolving lithium salts such as LiPF6 in an aprotic solvent such as ethylenecarbonate (EC), propylenecarbonate (PC), dimethylcarbonate (DMC), diethylcarbonate (DEC) or mixtures of them, but many other solvents and salts may be used for this purpose. The choice of the salt – solvent combination will determine some important properties such as the conductivity and the thermal and electrochemical stability of Li cells. An alternative to liquid electrolyte is the use of solid Li-ion electrolytes (SLIE). The main advantage of SLIE is flexibility, processability of plastics structure and reduced flammability. In the absence of free liquid, cheap, low-weight plastic containers may be used instead of heavier and more expensive stainless steel containers. In addition, the membrane separator, used to avoid short circuiting between electrodes in liquid electrolytes, can be suppressed, and sealing requires only low-temperature heating. Different types of SLIE may be used, depending on the aim and intended application: (i) Li-ion-conducting glasses and ceramics [1], (ii) Li-doped plastic crystals [2], (iii) Li salt in solid polymer electrolytes [3] such as poly(ethyleneoxide) (PEO) and (iv) plasticised electrolytes or polymer-gel electrolytes (PGE). Solid electrolytes (i)–(iii) exhibit conductivities in the range 105–104 mS cm1 at ambient temperature. These conductivities are too low for applications such as video cameras, laptop computers or cellular phones. Only PGE electrolytes can compete with liquid electrolytes with conductivities near or above 102 mS cm1 at room temperature. PGEs are obtained by mixing a polymer and a Li-ion-conducting liquid electrolyte. Many polymer matrices may incorporate a liquid electrolyte, but at present, the most commonly used polymers are, poly(vinylidene fluoride) (PVDF), poly(ethylene oxide) (PEO), poly(acrylonitrile) (PAN), poly(methyl methacrylate) (PMMA), poly(vinylidene carbonate) (PVdC), poly(vinyl chloride) (PVC), poly(vinyl sulfone) (PVS), poly(ethylene glycol acrylate)(PEGA), poly( p-phenylene terephthalamide) (PPTA) and poly(vinyl pyrrolidone) (PVP).

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1.4. Lithium salts

Many lithium salts may be dissolved in aprotic solvent to make up electrolytes: LiClO4, LiPF6, LiBF4, LiAsF6, LiCF3SO3 (LiTf), LiN(CF3SO2)2 (LiTFSI), LiC(CF3SO2)2, LiN[(C2F5SO2)2] (LiBETI), etc. With the notable exception of LiClO4, all of these salts have fluorinated anions; the anions are those of superacids like HPF6, HBF4, etc. The salts are generally highly soluble in dipolar aprotic solvent owing to the delocalised charge of the large anions and the electron-withdrawing properties of fluorine atoms. But LiF, for example, is not very soluble and is often found as a solid deposit on the surface of the electrodes. Fluorine atoms, which are introduced by means of the salt anion in the electrolyte, play various roles in Li-ion batteries: (i) they are components of the passivation film formed on Al current collectors when the cathode is raised at high potentials; and (ii) they are found in the form of solid deposits such as LiF, fluorophosphates LiPOxFy [4] or fluoroborates LiBOxFy [5] in the solid electrolyte interface (SEI) formed on the carbon electrode at the first cycle. It has also been shown that dissolved fluorinated salts decrease the surface tension of electrolytes and enhance the wettability of composite electrodes and porous separators [6]. 1.5. Lithium hexafluorophosphate

Of the different lithium salts that can be used as a component of the electrolyte, only LiPF6 is really employed at the industrial scale as seen in Table 1 where is reported the nature of electrode materials and the composition for commercial batteries. The success of LiPF6 is mainly due to a combination of wellbalanced properties such as ion mobility, ion-pair dissociation, solubility, thermal stability, chemical inertness, surface chemistry (solid electrolyte interface (SEI)) and collector passivation. The properties of LiPF6 are compared with those of other lithium salts and are classified from best to worst in Table 2. A major drawback of LiPF6 solutions is their poor stability at elevated temperatures. The decomposition reaction may be written as LiPF6 y PF5  LiF PF5  H2O → 2 HF  POF3 which may be summed up as LiPF6  H2O → LiF  POF3  2 HF Additives such as organo – silicon compounds [7] may decrease the detrimental effect of LiPF6 decomposition products. In this chapter, the use of fluorine in lithium-ion batteries will be examined with a special focus on fluorinated lithium salts such as LiPF6 and fluorinated

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Table 1 Secondary Li or Li-ion batteries [16] Electrolyte (salt/ solventa)

Negative/positive

Company or Institute

Li/MoS2

LiAsF6/PC co-solvent

Moli Energy (Canada)

Li-Al/TiS2

LiPF6/MeDOLDMEadditive

Hitachi Maxell (Japan)

Li alloy/C

LiClO4/PC

Matsushita Battery (Japan)

Li-Al/polyaniline

LiClO4/PC

Bridgestone-Seiko (Japan)

Li-C/LiCoO2

LiPF6/PC DEC

Sony Energytec (Japan)

Li-C/LiCoO2

LiBF4/PC  EC  BL

A&T Battery (Japan)

Li-C/LiCoO2

LiPF6/EC  DEC  co-solvent

Matsushita Battery (Japan)

Li-C/LiCoO2

LiPF6/EC  co-solvent

Sanyo (Japan)

Li-C/Li1xMn2O4

LiPF6/EC  DMC

Bellcore (USA)

Li-C/LiNiO2

(LiPF6 or LiN(CF3S02)2/ EC  co-solvent)

Rayovac (USA)

Lu LixMnO2

(organic electrolyte)

Tadiran (Israei)

Li/TiS2

Li-Li3PO4-P2S2

Everready (USA)

Li/V6013

LiX/PEO-based polymer

Valence Technology (USA)

a

MeDOL, 4-methyl-I,3-dioxolane; DEC, diethyl carbonate; EC, ethylene carbonate; DMC, dimethyl carbonate.

Table 2 Classification of Li salts [42] From Best → to Worst

Property Ion mobility

LiBF4

LiClO4

LiPF6

LiAsF6

LiTf a

LiTFSI

Ion pair dissociation

LiTFSI

LiAsF6

LiPF6

LiClO4

LiBF4

LiTf

Solubility

LiTFSI

LiPF6

LiAsF6

LiBF4

LiTf

Thermal stability

LiTFSI

LiTf

LiAsF6

LiBF4

LiPF6

Chem. inertness

LiTf

LiTFSI

LiAsF6

LiBF4

LiPF6

SEI formation

LiPF6

LiAsF6

LiTFSI

LiBF4

Al corrosion

LiAsF6

LiPF6

LiBF4

LiClO4

LiTf  lithium triflate.

a

LiTf

LiTFSI

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D. Lemordant et al.

polymers such as PVDF. Here we will identify the properties of the electrolytes based mainly on examples taken from our own work, but references to other works will be given when necessary. 2. LiPF6 AS FLUORINATED LITHIUM SALT 2.1. Preparation of LiPF6

There are two shortcomings in the synthesis of LiPF6: (i) the small size of the Li ion and (ii) the thermal instability of the salt. The first prevents the crystallisation of the salt with large anions such as PF6 and the second prevents the removal of the solvent used for the synthesis. Historically, Lange and Muller [8] realised the first synthesis of alkali metal derivatives of HPF6 with the exception of LiPF6. The presence of an excess of LiOH, of water and of some by-products, contributes to impede the crystallisation of the salt. Syed Mohamed et al. [9] relate that it is difficult to re-crystallise LiPF6 obtained by the reaction of pyridinium hexafluorophosphate (C5H5NHPF6) with LiOH in water. In non-aqueous media, LiPF6 may be obtained by the reaction of BrF3 and LiF in the presence of an excess of P2O5 or by reaction of PF5 with LiF in anhydrous HF. The latter reaction leads to an excess of LiF and LiHF2 as by-products. The synthesis, known as the Wiesboek process [10], gives battery-grade LiPF6. The acetonitrile complex Li(CH3CN)4PF6 is first obtained as a solid product by the reaction of PF5 and LiF at high pressure and low temperature (40 to 80°C) in the presence of acetonitrile. The complex is then decomposed under vacuum at low temperature. In this chapter, the production of battery-grade LiPF6 will be examined, taking a commercial HPF6 aqueous solution as starting material (HPF6; 65%, by wt). HPF6 itself, synthesised from P2O5 and HF or PF5/HF and SO2, is obtained as an oily product soluble in water. Strong aqueous solutions fume in air and gradually decompose. HPF6 forms hydrates, like H3OPF6 HF·4H2O, solvates and insoluble salts with many organic bases. 2.2. Characterisation of HPF6 aqueous solutions

The conductimetric titration curve of an aqueous HPF6 solution by lithium hydroxide (LiOH) is reported in Fig. 2. Two endpoints, denoted as V1 and V2, are clearly visible on the titration curve. The first endpoint, which represents 64% of the total acidity, is relative to the neutralisation of the strong acid HPF6 (the conductivity decreases as the hydronium ion is replaced by Li). Between V1 and V2, all weak acids (by-products of the synthesis or decomposition products) such as HPO2F2, H3PO3F, H3PO4 and HF [11] are neutralised. The main advantage of the conductimetric titration is that it makes it possible to visualise the first endpoint that corresponds to the strong HPF6 alone. If the neutralisation is stopped at this

Physicochemical properties of fluorine-containing electrolytes for lithium batteries

143

Fig. 2. Conductimetric titration of a commercial HPF6 aqueous solution (0.196 g in 20 mL H2O) by LiOH (0.098 mol/L). The first end point (V1) corresponds to the neutralization of the strong acid HPF6. The second end point is relative to acidic by-products.

point, the formation of Li salts other than LiPF6 is avoided. This also explains why an excess of Li is obtained when the titration is followed by an acid – base indicator. 2.3. Synthesis of LiPF6 from HPF6 solution and LiOH

The neutralisation of a sample of a commercial HPF6 aqueous solution by LiOH has been realised in the presence of an excess of acetonitrile. As LiOH is poorly soluble in acetonitrile, the reaction is achieved only after 72 h under reflux. The solvate Li(CH3CN)4 PF6 precipitates from the solution and is isolated by filtration. It is then decomposed under vacuum with a global 66% yield in LiPF6. 2.4. Synthesis of LiPF6 from C5H5NHPF6 and LiOH

As it has been shown that the pyridinium cation in C5H5NHPF6 can be easily exchanged [9] with Na, NH4 or K using the corresponding base, a novel method for the preparation of battery-grade LiPF6 has been proposed. 2.4.1. First step

The commercial HPF6 aqueous solution is first neutralised by the exact amount of pyridine at 0°C, which is determined using a conductimetric titration as described in Section 2.2. The pyridinium salt of HPF6 precipitates in water and is isolated by filtration. C5H5NHPF6 is then re-crystallised in water and in absolute ethanol. The yield is 60–70% and the purity of this product is confirmed by the elemental analysis reported in Table 3.

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Table 3 Elemental analysis of C5H5NHPF6 obtained by neutralization of HPF6 acid with pyridine C (%)

N (%)

F (%)a

P (%)

Calculated

26.68

6.22

50.64

13.76

Obtained

26.67

6.11

45.90

13.50

a

The amount of fluorine in lithium salt cannot be determined as precisely as the other elements.

2.4.2. Second step

The second step is the exchange reaction of the pyridinium by the lithium ion. This can be realised either (i) directly by the action of LiOH or (ii) indirectly using the solvate Li(C5H5N)PF6 as an intermediate product. (i) The reaction between LiOH and C5H5NHPF6, dissolved in a stoichiometric amount in dry methanol, is fast. The neutralised solution is concentrated using a rotary evaporator. In order to remove the last traces of water, an adequate volume of benzene is added and the ternary azeotrope methanol/benzene/water is removed by distillation. The pyridine is then evaporated under vacuum at moderate temperature. (ii) As the last step in the preceding method is delicate, the use of lithium ethylate or lithium methylate (LiOR), as indicated below, seems preferable. LiOR is formed in situ by the reaction of metallic lithium with the corresponding alcohol. C5H5NHPF6, added stoichiometrically to LiOR in alcoholic solution, reacts immediately. After the evaporation of the solvent at 30°C under partial vacuum, a white powder identified as the pyridine solvate of LiPF6 (Li(Py)PF6), remains. The decomposition of Li(Py)PF6 is achieved in a vacuum oven (P 1 Pa and t  50°C) and leads to crystallised LiPF6 with a yield of ca. 96%. The removal of pyridine is easily controlled by the disappearance of the IR absorption bands of this compound. The purity of LiPF6 obtained by this method exceeds 99%. 3. TRANSPORT PROPERTIES OF SOLUTIONS OF LiPF6 AND OTHER FLUORINATED LITHIUM SALTS The basic requirements for electrolytes used in Li-ion batteries are: (i) a good solubility of Li-salts, i.e. 1 mol/L (1 M); (ii) a conductivity of 10 mS/cm at room temperature; (iii) a wide electrochemical window (0–4.2 V vs. Li/Li); (iv) the formation of a stable solid electrolyte interface (SEI) at the carbon electrode; and (v) the formation of a passive film on the cathodic current collectors (Al). The transport and electrochemical properties of Li battery electrolytes are examined in the next section. The transport and electrochemical properties that

Physicochemical properties of fluorine-containing electrolytes for lithium batteries

145

are of concern to this study are viscosity, conductivity, electrochemical window and formation of an SEI at the negative electrode. 3.1. Viscosity 3.1.1. Influence of the electrolyte concentration

Viscosity studies provide useful insights into the mobility of ions in liquid or gel electrolytes. The geometry of molecules and ions is an important factor to take into account when the viscosity of a solution is considered. When a salt is dissolved in a dipolar aprotic solvent or a mixture of dipolar aprotic solvents, the viscosity of the solution (η) increases with the salt concentration C. The increase in viscosity is mainly due to ion – solvent (ion – dipole) and coulombic ion – ion interactions. The variation of the relative viscosity (ηr  η/ηo) can be expressed as a polynomial development in the power of C [12]:

ηr  1  B C  D C 2  …

(1)

At concentrations in the range 0.1 M C 2 M (where M1 mol/L), it is not necessary to take into account polynomial terms over C 2. Eq. (1) is in accordance with the equation proposed by Einstein [13] for the variations of the relative viscosity of solutions containing spherical unsolvated particles:

ηr  1  2.5 Φ  o(Φ2)

(2)

where Φ is the volume fraction of the particles in the fluid. As Φ can be expressed as a function of C and the hydrodynamic molar volume Vi of the particles, Eq. (2) may be rewritten as

ηr  1  0.0025 Vi C  o(C 2)

(3)

where Vi is expressed in mol/cm3. If it is supposed that Eq. (3) holds at the molecular scale for large molecules in organic solvents, the identification of Eqs. (1) and (3) leads to Bcalc ≡ 0.0025Vi. Using this equation, it is possible to deduce a value of the hydrodynamic radius of the anion of the salt if the radius of the solvated cation is known. As an example, the B and D values obtained by curve fitting, for solutions of LiBF4, LiPF6, LiAsF6 and LiTFSI in γ-butyrolactone as solvent are displayed in Table 4. In contrast to B, for which the variations with the temperature are negligible, the D coefficient decreases as an inverse function of the temperature (thermal agitation). As B is linked to the hydrodynamic volume of ions, it is not too surprising to find it almost independent of the temperature. From the mean value of B (Bmean), the volume of the solvated salt in the range of temperature investigated has been deduced using Eq. (3) and the relation Bmean  Bcalc ≡ 0.0025Vi.

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Table 4 B (M1) and D (M2) coefficients from Eq. (1). The radius of the anion of the salt (R in nm) is deduced from Bcalc(≡ 0.0025Vi ) and compared with literature values (refer to text for details) LiBF4 t (°C)

LiPF6

LiAsF6

LiTFSI

B

D

B

D

B

D

B

D

25

0.30

0.75

0.40

1.10

0.40

1.15

0.50

1.20

35

0.30

0.70

0.40

0.95

0.42

1.00

0.47

1.10

45

0.30

0.65

0.38

0.85

0.42

0.85

0.46

1.05

55

0.25

0.55

0.42

0.68

0.41

0.75

0.50

0.90

Bmean

0.288

0.400

0.413

0.483

0.013

0.008

0.005

0.013

a R(anion)cal

0.14

0.28

0.29

0.33

R(anion)lit

0.227

0.254

0.259

0.326

Taking R(Li)  0.346 nm.

a

Taking r (Li, solvated)  0.346 nm [14], the volume and the radius of the anions may be calculated. The results are also reported in Table 4, considering anions as non-solvated spheres. The agreement between the experimental and calculated values of the anions radii is good with the exception of BF4, for which the radius value is too low. This could be attributed to the fact that the ions in LiBF4 associate more easily in ion pairs than in the other salts. In conclusion, viscosity experiments support the validity of Eqs.(1) and (2) and show that volume effects are predominant when salts are composed of large anions with delocalised charge and small strongly solvated cations. 3.1.2. Influence of the temperature

Electrolyte solutions follow an Arrhenius (Eq. (4)) or a VTF (Eq.(5)) law:

η  A exp(Ea,η /RT)

(4)

η  A exp (B/R(T  T o))

(5)

With γ-butyrolactone as solvent, LiBF4, LiPF6, LiAsF6 and LiTFSI follow an Arrhenius law at all concentrations in the range 10–50°C [15]. The energy of activation for the viscous flow of the electrolyte solutions has been determined

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147

by plotting ln(η) vs. 1/T for all electrolyte solutions. The variation in activation energy between the electrolyte solution and the pure solvent (Ea,η  10.7 kJ/mol),

δEa,η  Ea,η(Csalt ⬆ 0)  Ea,η(Csalt  0) is reported in Table 5 as a function of the concentration in salt. As expected, Ea,η increases with the salt concentration for all electrolytes, but, with the exception of LiBF4, the differences in δEa,η values between salts are in the range of the experimental error. This means that there is no influence of the nature of the salt. 3.2. Conductivity 3.2.1. Influence of the solvent

The conductivity at 25°C of EC-based binary solvents (mixtures are 1:1, v/v) containing lithium triflate (LiTf), LiPF6 and LiTFSI at the same concentration C  1 mol/L are reported in Table 6 [16]. LiPF6 and LiTFSI give by far the most conductive solutions. As LiTf is generally less dissociated than the other two salts in most dipolar aprotic solvents, its conductivity is lower. This explains why this salt, thermally stable and non-hygroscopic, cannot be used in powerful Li-ion batteries. One may also remark that the less viscous the co-solvent, the higher is the conductivity. In concentrated electrolyte solutions, the viscosity of the solvent medium controls the conductivity when the salt dissociation is high. 3.2.2. Influence of the salt concentration 3.2.2.1. Variation of the molar conductivity with the salt concentration The molar con-

ductivity (Λ) is defined as Λ  κ /C, where κ is the conductivity and C the concentration of the salt. When Λ is plotted against C1/3, a straight line is obtained as

Table 5 Variations in activation energy for the viscosity δEa,η  Ea,η(Csalt ⬆ 0)  Ea,η(Csalt  0) as a function of the salt concentration (refer to text for details) C (mol L1)

0.1

0.2

0.5

1.0

1.5

LiBF4

2.4

2.6

3.7

4.8

6.5

LiPF6

0.4

1.3

3.3

7.0

8.7

LiAsF6

1.3

1.7

2.0

4.9

6.6

LiTFSI

0.6

1.0

1.8

3.8

6.8

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Table 6 Conductivity (in mS/cm) of EC-based binary solvent mixtures (1:1, by vol.) containing LiTf, LiPF6 and LiTFSI. The co-solvents are dimethoxyethane (DME), methylpropionate (MP), dimethylcarbonate (DMC) and diethylcarbonate (DEC) Co-solvent

Co-solvent viscosity (mPa s)a

LiPF6

LiTFSI

LiTf

DME

0.455

16.6

13.3

8.3

MP

0.49b

13.3

10.8

3.7

DMC

0.58

11.2

9.2

3.1

DEC

0.748

7.8

6.5

2.1

a

From Ref. [42]. Estimated from homologues esters.

b

50 45

Λ /mS cm-1 M-1

40 35 30

LiAsF6

25 20 15 10

LiBF4

5 0 0.3

0.4

0.5

0.6

0.7 0.8 C1/3 / M1/3

0.9

1.0

1.1

1.2

Fig. 3. Verification of the cube root law (Eq. (6)) for the molar conductivity of LiBF4 and LiAsF6 in BL at 25°C.

shown in Fig. 3 for LiBF4 and LiAsF6 in butyrolactone at 25°C. This means that the cube root law [17] Λ  ΛoS C1/3

(6)

which can be deduced from the pseudo-lattice model [18,19], is verified by these solutions. In Eq. (6), Λo is the ordinate value, which is obtained by extrapolation of Λ variations at C  0. Generally, this value is slightly different from that

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149

obtained using the extrapolation law in C1/2 issued from the Debye and Hückel theory. The slope S is dependent on Λo, some characteristics of the solvent such as the relative permittivity and the viscosity, and of the temperature: S  9.61Λo/εr  2.88 106/η

(7)

In Table 7, Λo, Λo and S (experimental Sexp and calculated Scalc) for LiBF4, LiPF6, LiAsF6, LiTFSI in γ-BL at 25°C are reported. For the salts under study  Sexp ⬇Scalc , but as expected Λo ⬆Λo, as these values are obtained from two different extrapolation laws. At infinite dilution, Λo values fall in the following order: LiPF6  LiAsF6  LiTFSI  LiBF4 This is, with the exception of LiBF4, the increasing order of the size of the anions. As it has been reported that the dissociation coefficient of LiBF4(1M) at 25°C in BL is 0.34 [20], organic solutions of this salt involve free ions and ion pairs, leading to an underestimated value for the ordinate in the Λ  f (C1/3) plot. 3.2.2.2. Existence of a maximum in the conductivity curves κ  f(C) When the con-

centration of salt is raised in an electrolyte solution, the number of charge carriers increases and, as seen previously, the viscosity increases. The consequence is that, inevitably, ion mobility decreases. The competition between the increase in number of charge carriers and the decrease of their mobility leads to a maximum in the conductivity–concentration relationship. This maximum is often observed around 1 M in these organic electrolytes. For evident reasons, the determination of this maximum in conductivity is important for battery applications. Table 7 Λo, Λo and Sexp values from Eq. (6) and Scalc from Eq. (7) for LiBF4, LiPF6, LiAsF6, LiTFSI in γ-butyrolactone at 25°C Salts

LiBF4

LiPF6

LiAsF6

LiTFSI

Λο

42.64a

40.73a

39.91a

33.77a

Scalc

26.94

26.40

29.49

24.76

ο

Λ

30.29

48.32

46.77

35.74

Sexp

22.92

36.55

36.00

26.47

a

From Ref. [43].

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D. Lemordant et al.

Using Eq. (6) and the definition of the molar conductivity, the specific conductivity κ, may be expressed as

κ  Λo C  S C4/3

(8)

The maximum conductivity value (Cmax) is easily deduced from Eq. (8) by the derivation of κ relative to the concentration: Cmax  (3Λo/4S)3

(9)

Cmax is then related to the molar conductivity at infinite dilution Λo and to the slope S. As an example, the variation of κ with the concentration is given in Fig. 4 for LiBF4 and LiAsF6 at 25°C in BL. As seen in Table 8, a good agreement is obtained between the calculated (by Eq. (9)) and the experimental values of Cmax in BL. Cmax can be predicted for a different solvent than BL, using only Λo as experimental entry (approximated by Λo if necessary) as all other parameters in S may be calculated. Nevertheless, as the pseudo-lattice model disregards the formation of nonconducting contact ion pairs, the quasi-lattice theory cannot be readily extended to all organic electrolytes, especially those having low permittivities. 3.2.3. Influence of the temperature

The influence of temperature on the molar conductivity (or the specific conductivity) of electrolyte solutions is well described by either an Arrhenius equation (Eq. (10)), at temperatures well over the vitreous transition (if it occurs),

12

K / mS cm-1

10

LiAsF6

8 LiBF4

6

4

2 0

0.40

0.80

1.2

1.60

C /mol L-1

Fig. 4. Variation of the conductivity of electrolytes solutions with salt concentration: LiPF6 and LiBF4 in γ-butyrolactone at 25°C.

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151

Table 8 Salt concentration (in M) at the maximum of conductivity (Cmax) at 25°C in BL. Eq. (9) is used for calculated values Salt

Cmax,exp

Cmax,calc

LiBF4

1.02

0.90

LiPF6

0.98

0.90

LiAsF6

1.01

0.90

LiTFSI

0.97

0.90

or by the VTF equation (Eq. (11)) : Λ  Λo exp(Ea,Λ/RT)

(10)

Λ  Λo exp(BΛ /R(T  T o)

(11)

From the pseudo-lattice theory, a relation between the activation energy for the conductivity (Ea,Λ) and the salt concentration may be inferred [21]. This relation has been successfully applied to the BL–LiClO4 system [15]: o Ea,Λ  Ea,Λ + kC 4/3

(12)

where k is a coefficient taking into account ion–ion and ion–dipole interactions. Ea,Λ vs. C 4/3 has been plotted in Fig. 5 for LiBF4, LiAsF6, LiPF6 and LiTFSI in γBL. Experimental slopes (k) and ordinate (E oa,Λ) values obtained by linear regression for the different salts are reported in Table 9. The ordinate of the correlation o Ea,Λ , which represents the infinitesimal dilution activation energy for the conductivity, reaches a value that is close to the activation energy for the viscosity of pure γ-BL (10.7 kJ/mol). At infinitesimal dilution, the activation energy for the conductivity represents the interactions of 1 mol of “discharged” ions (their distance is considered too large for even long-range interactions to occur). The smallest k value is obtained for LiBF4. 3.2.4. Ion association

In 1926, Bjerrum [22] introduced the notion of ion pairs formed by the following chemical equilibrium: A  B → (A,B)

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14 14 13

13 Ea,Λ/kJ.mol-1

Ea,Λ/kJ.mol-1

12 11 LiBF4 LiAsF6

10

12 11 LiPF6 LiTFSI

10

9

9

8

8

7

7 0

0.2

0.4

0.6

0.8

1

1.2

0

C4/3 / M4/3

0.2

0.4

0.6

0.8

1

1.2

C4/3 / M4/3

Fig. 5. Activation energy for conductivity vs. C4/3 for LiBF4, LiPF6, LiAsF6 and LiTFSI in BL.

Table 9 Slopes of activation energy for conductivity vs. C 4/3 and activation energy at infinitesio mal dilution (Ea,Λ ) Salt

LiBF4

LiPF6

LiAsF6

LiTFSI

kexp (M4/3)

1870

2120

2300

2690

o Ea,Λ (kJ mol1)

10.8

10.7

11.0

10.9

This equilibrium is characterised by the thermodynamic constant Ka, Ka  (1α)/(α2 C γ 2 )

(13)

In Eq. (13), γ is the mean activity coefficient and α the dissociation coefficient. Usually, α is determinated by conductimetry, but at high ionic strength (I  0.1M) the determination of both α and γ is difficult as the classical laws cannot be applied. For these reasons, α is best spectroscopically determined using IR or Raman spectroscopy. At high ionic strength, ion pairing may occur and associated species like tight ion pairs (contact ion pair: Li…X) or solvent separated ion pairs (Li…S…X) have been evidenced by Raman spectroscopy. An example is the dissociation of LiAsF6 in PC or DMC, where the presence of contact ion pairs

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153

Li…AsF6, free AsF6, solvated ion pairs [Li(DMC)4]…AsF6 and polymeric ion pairs [Li…AsF6]x, has been shown in the concentration range 0.1–4.5 M [23]. Even in high permittivity solvents, ion pairs may be formed. This is in particular the case of lithium perchlorate in PC, where ion solvent separated pairs and contact ion pairs are observed [24]. Ion association has been investigated in BL for LiPF6. No association is expected for the TFSI anion, which has the longest delocalised structure. The Raman spectra of solid LiPF6 (as powder) and LiPF6 in BL at 0.5 and 1.5 M are reported in Fig. 6. The peaks present a Gaussian shape area, which is a function of the salt concentration (A  0.0285C with regression coefficient r  0.91). There is no evidence of a new band, which could be attributed to ion pairing. As a result, if ion pairs are formed at high concentration in salts, these would be solvent-separated or solvent – solvent-separated ion pairs. 3.3. Electrochemical properties of electrolytes containing fluorinated lithium salts 3.3.1. Electrochemical window

I/a.u

Usually, the electrochemical window of the electrolyte is determined at the platinum or glassy carbon electrode. All alkylcarbonates and lactones are strongly resistant to oxidation as well as reduction. Nevertheless, metallic lithium is deposited and re-oxidised with a bad faradic yield owing to the formation of a passive film on dendritic deposits. Using carbon as anode, the formation of the passive film (SEI) occurs at 0.7–0.8 V vs. Li/Li. The role of the salt in the formation of the SEI is not fully understood but the presence of LiF, fluorophosphates or fluoroborates indicates clearly that fluorinated salts like LiPF6 or LiBF4 are involved in the reduction mechanisms. As an example, the electrochemical window at the rotating platinum electrode of the electrolyte BL-EC (1:1 in mol)

0.01 0.009 0.008 0.007 0.006 0.005 0.004 0.003 0.002 0.001 0 730

LiPF6 BL+LiPF6 1.5M BL+LiPF6 0.5M

740

750

760

770

780

790

υ / cm-1

Fig. 6. Raman spectra of solid LiPF6 (as powder): 0.5 M LiPF6 and 1.5 M LiPF6 in BL.

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D. Lemordant et al.

 LiPF6 (1 M) or LiBF4 (1 M) is reported in Fig. 7. The electrochemical window limit in oxidation (5.7 V vs. Li/Li) has been determined by measuring the oxidation potential at a current density of 20 μA cm2. This electrochemical window is large enough for use in Li-ion batteries. It may be noted that when active materials (LiNiO2 or LiCoO2) are used instead of platinum, the electrochemical window is more limited due to the catalytic properties of the lithiated oxide. 3.3.2. Lithium insertion into graphite

The key to the working of the graphite anode is the chemistry of alkylcarbonates and lactones. In the presence of a restricted number of Li salts containing fluoroanions or oxalate derivatives, a thin organo-mineral layer is deposited on the electrode. This layer acts as a barrier towards the solvent and hence prevents the exfoliation of the graphite structure. This layer must also be ion-conductive to permit the intercalation of non-solvated Li ions into the graphite layers. In Fig. 8 the first charge–discharge cycle at a current density of 110 μA cm2 of a half C(graphite)/electrolyte/Li cell is reported. The electrolyte is a BL/EC mixture (1:1 or 9:1) containing LiPF6 or LiBF4. Lithium insertion does not occur during the charging process when the electrolyte contains LiPF6 and all other electrolytes (LiAsF6, LiClO4,…), with the exception of LiBF4 (1 M), exhibit the same behaviour. In the presence of LiPF6, no lithium insertion occurs and a high internal resistance is observed. A major part of the ohmic resistance arises from the insulating properties of the SEI film formed on graphite [25]. The SEI layer is then dependent on the nature of the anion. As LiPF6-based electrolytes cannot cycle on the graphite electrode in BL, the cycling ability of the graphite electrode is determined using LiBF4-based 300

j / mA.cm-2

250 200 150

LiPF6

100

LiBF4

50 0 2.5

3

3.5

4

4.5 5 6 E vs Li/Li+ / V

6.5

7

7.5

Fig. 7. Electrochemical window at the rotating platinum electrode of a BL/EC (1:1) mixture in the presence of 1 M LiPF6 or 1 M LiBF4; the rotating speed is 1000 rpm and the scan rate 5 mV/s.

Physicochemical properties of fluorine-containing electrolytes for lithium batteries

155

3.49 2.99

E / V vs. Li+/Li

2.49 1.99 1.49 0.99 0.49 -0.01 0:00

2:24

4:48

7:12

(a)

9:36 12:00 14:24 Times in hours

16:48

19:12

21:36

3.49 2.99

E / V vs. Li+/Li

2.49 1.99 1.49 0.99 0.49 -0.01 0:00

24:00

48:00

72:00

(b)

96:00 120:00 144:00 168:00 192:00 216:00 240:00 Times in hours

Fig. 8. Charge – discharge cycles at C/20 and D/20 (25°C) of a BL/EC (1:1) electrolytes with (a) 1 M LiBF4 and (b) 1 M LiPF6.

electrolytes. Irreversible capacities and reversible capacities vs. cycle number for LiBF4-based electrolytes at 25°C in BL/EC (1:1) and BL/EC (9:1) are represented in Fig. 9. The reversible capacity is higher in the eutectic mixture BL/EC (9:1) than in the equimolar mixture but exhibits more fading. The irreversible capacity is also lower in the 9:1 mixture and decreases to near zero at the 5th cycle. As the behaviour of these electrolytes at 60°C is similar this means that the SEI layer is stable at this temperature. 3.3.3. Morphology of the SEI layer

The morphology of the SEI layer formed on the negative electrode in LiPF6 and LiBF4 (purity 99.9% and 99.99%) solutions in BL:EC (1:1) has been studied by scanning electronic microscopy (SEM). The morphologies of the passivation film obtained with LiPF6 and LiBF4 (99.9%) are very different (Fig. 10a and b). A large number of nuclei of similar size are observed in Fig. 10b for LiBF4(99.9%),

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D. Lemordant et al.

BL-EC (9:1)

350 330 310 C/mAh g-1

290 270

BL-EC (1:1)

250 230 210 190 170 0

1

2

(a)

3 4 Cycle number

5

6

5

6

140 120

C/mAh g-1

100 80 60

BL-EC (1:1)

40 20

BL-EC (9:1)

0 0 (b)

1

2

4 3 Cycle number

Fig. 9. (a) Reversible capacity and (b) irreversible capacity vs. cycle number during charge–discharge cycles at charge and discharge rates of C/20 at 25°C for BL/EC (1:1) and BL/EC (9:1), and LiBF4 (1 M).

whereas the nuclei are fewer with LiPF6 (Fig. 10a). The SEI layer is thicker in the presence of LiPF6 because the shape of the graphite flakes is not clearly visible: the SEI layer coats the totality of the graphite surface. This morphology may be responsible for the insulating character of the SEI and the bad cycling results. Using battery-grade LiBF4 (99.99%), a homogeneous and dense passivation film is observed, as reported in Fig. 11. This result shows that impurities may have a strong influence on the morphology, and possibly on the quality of the SEI. In the next paragraph, the composition of the SEI layer by XPS will be studied.

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Fig. 10. SEM pictures of the morphology of the SEI layer obtained in BL/EC (1:1) in the presence of (a) LiPF6 (1 M, 99.9%) and (b) LiBF4 (1 M, 99.9%).

Fig. 11. SEM pictures of the morphology of the SEI layer obtained in BL/EC (1:1) in the presence of battery grade LiBF4 (1 M, 99.99%).

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3.3.4. XPS analysis of the SEI layer

In order to investigate the SEI layer formed on the graphite anode in BL:EC (1:1) in the presence of LiPF6 (1 M) and LiBF4 (1 M), samples have undergone one charge – discharge cycle with a constant current density of 220 μA cm2. The C1s, O1s, Li1s, F1s and P2p peaks are determined by this XPS analysis. The main component of the C1s peak at 284.6 eV is a result of contributions of the carbon contamination layer and graphitic carbon; the component at 286.5 eV and 288.9 eV indicates the existence of, respectively, C–O–C groups and alkylcarbonate (RCO3Li). The broad O1s peak at about 531.9 eV is consistent with the previous attributions. The analysis of the P2p peak reveals the existence of two distinct doublets that may be assigned to LixPOy species (133–134.7 eV) and LixPFy species (136.7–137.8 eV). The F1s peak is divided into two components, 685.3 and 687.4 eV, which are attributed to LiF and LixPFy species. In order to obtain information on the chemical composition of the SEI layer as a function of depth, high-resolution XPS spectra were recorded after 5, 20 and 40 min of sputtering. The results indicate a progressive increase in the carbon percentage, the disappearance of the alkylcarbonate and a decrease in the percentage of LixPFy species. Only LiF is present in the whole width of the SEI layer. These results are in fair agreement with the SEI model constituted by a mineral layer on the graphite surface and a complex organo-mineral layer over it. The results obtained show the existence of solvent reduction species and other species from complex reactions (LiF, LixPOy and LixPFy). These results are in agreement with those reported by Aurbach [26], concerning the analysis of the SEI layer on lithium with EC-DMC in presence of LiPF6 (1 M). Unlike with the previous results, the reproducibility of the XPS analysis with LiBF4(99.9%) was not observed. We can explain this phenomenon by the variable purity of the different samples. Nevertheless, a main trend has been obtained. On the graphite surface, we have observed a higher ratio of LiF and a low ratio of the other elements (C,O,B). These trends remain the same after 40 min of argon-ion sputtering as was noticed by Kanamura for an electrolyte containing BL and LiPF6 [27] . When the LiBF4 salt is of “battery grade” (99.99%), the same species as previously observed with LiPF6 are evidenced for the C1s peak but with a higher ratio of the oxygenated elements: components at 286.4 and 289.3 eV. Compared with LiPF6, the higher binding energy of this last component can be attributed to the existence of a small quantity of Li2CO3. The most important difference is observed for the F1s peak: a component at 685 eV attributed to LiF and another at 686.7 eV attributed to LixBFy. Moreover, less LiF is deposited than previously. The B1s peak shows the existence of oxygenated borate environment (193 eV) and fluorinated borate environment (195.5 eV) that may be assigned, respectively, to LixBOy and LixBFy species. From the depth profile of each element, it can be seen that the species arising from the solvent reduction (286.5 and 290 eV) forms a thicker layer than that

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Fig. 12. Scheme of the SEI film obtained from LiPF6 (purity  99.99%), LiBF4 (purity  99.9%) and LiBF4 battery-grade (purity 99.99%) based electrolytes.

for LiPF6. LiF is the only species in the SEI layer formed after 40 min of argonion sputtering. These results show that a different chemical composition and a higher thickness of the SEI layer are obtained with LiBF4 and LiPF6 electrolytes. From the XPS study, a schematic illustration of the composition and relative size of the SEI layers is shown in Fig. 12. As phosphates are known to polymerise easily, the presence of LixPFy and LixPOy in the SEI can be indicative of a polymer coating of the graphite electrode that may be responsible for its insulating properties. 4. PVDF–HFP–SiO2 PEG SYSTEM Previous works have revealed the various appealing properties of PVDF–HFP copolymer as a host for liquid electrolytes for applications in rechargeable lithium batteries [28–31]. The ionic conductivity of polymer gel electrolytes, prepared from PVDF–HFP copolymer, is due to the conductivity of the liquid electrolyte embedded in the pores [32]. As a consequence, it is largely affected by the porosity and the wetting properties of the membrane [33]. In this study, γ-valerolactone (VL) has been employed as a single liquid solvent or as a mixture with ethylenecarbonate (VL/EC, 80:20, in mol) and LiTFSI as lithium salt.

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The main advantages of LiTFSI are its high conductivity, its thermodynamic stability and safety characteristics, [34,35] although its practical uses are limited owing to its severe corrosion of Al current collector above 3.6 V vs. Li/Li. The solvent VL has a high solvating power towards lithium cation and is stable towards metallic lithium [36]. Furthermore, VL exhibits a low melting point (31°C), a high boiling temperature (207°C), a high permittivity (εr  32 at 25°C) and moderate viscosity (η  1.89 mPa s at 25°C). The addition of ethylene carbonate (EC) to VL improves the stability of the SEI layer at the carbon anode – electrolyte interface [36,37], but raises the melting point of the solution and limits its operating temperature range. For this reason, the molar percentage of added EC has been limited to 20%. The anodic electrochemical stability of the VL/EC(80:20)–LiTFSI (1 M) electrolyte, determined on platinum (Pt), nickel (Ni) and stainless steel (SS) electrodes, are respectively 5.7, 4.8 and 4.4 V vs. Li/Li. The main aim of this section is to expose the influence of the salt content in the liquid phase, of the silica content in the dry polymer, and of the temperature on the rate of the wetting process. 4.1. Preparation of microporous polymeric membranes

PVDF-HFP copolymer (92/8, “SOLEF” from SOLVAY) is first dissolved in acetone at 40°C, then dibutyl phatalate (DBP) and silica (AEROSIL R 972, from DEGUSSA-HULS), dried under vacuum at 260°C for 24 h before use, are added. The mixture is homogenised under argon by vigorous stirring. The resulting solution is cast on aluminium plates and the acetone is slowly evaporated. Next, the membrane is weighed and immersed in diethyl ether for about 24 h under continuous stirring in order to extract dibutyl phatalate (DBP, 99%), the initial plasticiser. After drying at 60°C under vacuum, the sample is weighed again to verify the complete extraction of DBP. 4.2. Absorption of the electrolyte

The dry copolymer membrane is punched into circular pieces, whose diameter and thickness are measured with an electronic calliper square. After being weighed accurately, the samples are immersed in the thermostated liquid electrolyte. At regular intervals, the samples are taken out, pressed lightly between two sheets of clean filter paper and weighed. For a given membrane and a fixed electrolyte composition, experiments were run in triplicate or more. The measurements of the sizes and weight of the samples allow one to calculate the weight of liquid electrolyte absorbed by the membrane with time and, for some samples, the variations in surface area with time. The measurement of the relative densities gives the volumetric fraction of liquid electrolyte at saturation if it is assumed that no compression of the liquid phase in the gel occurs.

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4.2.1. Homogeneity of the absorption

Prior to examining kinetic results and absorption abilities of the copolymers, the homogeneity of the absorption process must be verified. This means that the composition of the absorbed liquid phase is identical to the initial liquid electrolyte whose composition is chosen to provide the highest conductibility as possible. Two methods may be employed for this purpose: (i) conductivity and (ii) Raman spectroscopy. (i) It is found that the conductivity of the liquid phase measured before and after absorption of a significant fraction of initial liquid volume, is identical, taking into account experimental error. (ii) The intensity of the Raman line at 741 cm1, which is assigned to the δs(CF3) vibration characteristic of TFSI anion [13,14], is measured in VL/LiTFSI solutions and in gels. In liquid phases, the integrated intensity of δs(CF3) mode is proportional to the molar percentage of the salt, as seen in Fig. 13 (marks are black squares). If there is no strong interaction between the anion and the polymer network in the gel, the intensity of this Raman line must follow the same linear correlation. Integrated intensities of the 741 cm1 Raman line for solvent-saturated PVDF-HFP copolymers have been plotted on the same graph (open squares) at different percentages in salt. The results described in Fig. 13 show that the representative points are in alignment with the line drawn. This confirms the homogeneity of the absorption of the liquid electrolyte by the copolymer. 4.2.2. Amount of liquid electrolyte absorbed with time

We will first examine the theoretical model used to describe the absorption kinetic, then the influence of liquid phase, copolymer composition and temperature

Fig. 13. Integrated intensities of δs(CF3) Raman vibration band at 741 cm1 for VL/LiTFSI solutions and PVDF-HFP/VL/LiTFSI gels. The line drawn on the graph is the linear regression analysis for VL/LiTFSI experimental results.

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on the rate of absorption. In the last section, we will examine the absorption capacity in relation to the same factors. The advancement of liquid electrolyte in the PVDF-HFP membrane, αab(t), is defined as the ratio of the adsorbed weight at time t (m1,t) to the adsorbed weight at saturation (m1,sat): 4.2.2.1. Diffusion model

αab(t)  m1,t /m1,sat

(14)

For 0 αab(t) 0.6, the kinetics of absorption is well described by a diffusion model based on Fick’s second law. This model implies the hypothesis, supported by experimental evidence, that in the first stage of the process a fast wetting of the surface of the copolymer occurs. This fast wetting corresponds to the saturation of the surface pores by the liquid phase, leading to the formation of a thin film, where the weight concentration C1,0 of the liquid phase is equal to its final value C1,sat: C1,0  C1,sat  m1,sat/Vsat

(15)

Starting from this saturated layer, the solution diffuses into the polymer. The mass flux of diffusing liquid, Jt , is given by [15,16]: Jt  (dm1,t /dt)/St  C1,sat D1/2 π1/2 t1/2  K t1/2

(16)

where D is the average diffusion coefficient of particles and St is the cross-sectional area of the sample at time t. The variations of St with time, obtained by measurement of the sample sizes at different times, are represented on the graph reported in Fig. 14, where the normalised area of the sample St /S0 has been

Fig. 14. Evolution of the surface area of PVDF-HFP copolymer wetted by a VL/LiTFSI (98:8) solution at 20°C.

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163

plotted against the advancement αab(t). From this correlation, the following relation is deduced: St  S0 Β m1,t , where B  (Ssat  S0)/m1,sat. By integrating Eq. (16) with respect to time, the weight of liquid phase absorbed at time t is obtained:

冕

t

m1,t  St Jt dt

(17)

0

After integration, and applying the initial condition: m1,t  m1,0 at t  0, ln[1  (B/S0) m1,t]  ln[1  (B/S0) m1,0]  2 Kt1/2

(18)

In Fig. 15, the variations of ln[1  (B/S0) m1,t] are plotted against t. Far from the saturation (t 45 min), a linear regression analysis in t1/2 permits one to determine the slope of the correlation and hence the diffusion coefficient as D  4 π [2 BK V1,sat /(Ssat  S0)]2

(19)

For the sample under study, the initial weight of dry polymer m0,P  0.1011 g and the values of the geometric and kinetic parameters at 20°C are reported in Table 10. 4.2.2.2. Taking the saturation into account In order to describe the whole absorption

process, it is assumed now that the rate of absorption is proportional to both the solution – copolymer interface area and the fraction of empty pores in the material. At time t, the fraction of empty pores is given by 1  (m1,t /m1,sat), which leads to (dm1,t /dt)  k(S0 B m1,t)(1  (m1,t /m1,sat))

(20)

where k is the kinetic constant.

Fig. 15. Experimental (markers) and simulated (drawn lines) evolutions of the absorbed weight of a liquid-phase VL/TFSI (88:12) by the copolymer PVDF-HFP/SiO2 (73/27) at 2.5°C and 30°C.

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Table 10 Values of geometric and kinetic parameters of the PVDF-HFP membranesa 20

2.5

30

S0 (cm2)

2.891

3.01

7.32

Ssat (cm2)

4.925

4.30

10.74

V0,P (cm3)

0.145

0.128

0.306

Vsat (cm3)

0.312

0.262

0.684

m1,0 (g)

0.049

0.051

0.075

m1,sat (g)

0.2806

0.245

0.628

2.2

1.92

3.53

Temperature (°C)

102kexp : mn2 a

Data from Ref. [44].

Integration of Eq. (20) with respect to the initial conditions gives ln [(a  m1,t) /(b  c m1,t )]  Y2,0  kexp t

(21)

In this expression a  S0 m1,t /(Ssat  S0); b  (Ssat  S0)/ m1,t; c  (Ssat  S0)/m1,sat2; kexp  k Ssat/m1,sat and Y2,0 , the integration constant, is equal to ln [(a  m1,0) /(b  c m1,0 )]. The linear regression analysis of Eq. (21) gives both m1,0  0.049 g and kexp  2.20 102 min1. The absorbed weight at time t may be deduced from Eq. (21) m1,t  [(b exp(Y2,0) exp(kexp t))  a]/[1  (c exp(Y2,0) exp(kexp t))]

(22)

In Fig. 16, the variations of m1,t with time at 30°C and 2.5°C have been plotted using the values of the parameters reported in Table 10. The validity of Eq. (22) is established by the fact that the experimental points fit well with the calculated curves. 4.2.3. Discussion of the kinetic results

The influence of the salt content in the liquid phase on the rate of the absorption has been investigated for the copolymer PVDF-HFP/SiO2 (73:27) and

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Fig. 16. Evolution of the advancement of absorption αab with time for copolymers PVDF-HFP of various silica contents at 20°C. The liquid phase is VL/LiTFSI (92:8). Markers are αab values from experimental determinations of m1,t and m1,sat for the samples: Xo(SiO2)  0 (䉫); Xo(SiO2)  0.157 (䉱); Xo(SiO2)  0.270 (ⵧ).

Table 11 Diffusion coefficient (D) and kinetic constants (k) at various temperatures for the absorption of VL/TFSI solutions in PVDF-HFP membranesa T (°C) 0

% LiTFSI

0

1

105 k (g cm2 s1)

2.36

1.82b

2.06

(cm2 s1)

0.75

0.68b

0.55

105 k (g cm2 s1)

3.02

2.58

3.12

(cm2 s1)

1.09

1.04

0.84

105 k (g cm2 s1)

3.69

3.02

2.23

2.54

3.72

4.23

(cm2 s1)

1.49

1.45

1.12

1.11

1.40

1.33

105 k (g cm2 s1)

4.83

3.50

5.99

(cm2 s1)

1.94

1.73

1.65

105 D 10

105 D 20

105 D 30

105 D

2

4

8

12

a

Data from Ref. [44]. Measured at 2.5°C.

b

VL/LiTFSI solutions, with 0 XLiTFSI 0.12 in the temperature range 0°C T 30°C. The values of the kinetic parameters k and D are reported in Table 11. The following conclusions can be inferred: (i) As expected, kinetic absorption constants and diffusion coefficients are closely related.

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(ii) Increasing the salt content with respect to the pure solvent first reduces the rate of absorption and, for XLiTFSI  0.04, enhances the rate of the process. This behaviour can be explained by taking into account two competitive effects: solvated ions are much larger entities than solvent molecules and steric hindrance for their transport in the pores increases the viscosity. As a consequence, the average mobility is reduced. Furthermore, the penetration of solvated ions into the smaller pores may be forbidden. Despite these considerations, salt polymer interactions favour liquid-phase migration inside the largest pores where no steric hindrance occurs. (iii) The diffusion coefficients and kinetic absorption constants exhibit an Arrhenius behaviour given by D  D0 exp(Ea,D /RT ) and k  k0exp(Ea,c /RT ) where Ea,D and Ea,c are the activation energies of transfer, and D0 and k0 are frequency parameters. The values of these parameters are given in Table 12. A small addition of salt slightly depresses both Ea,D and Ea,c but at higher salt contents these values are greater than those obtained for the pure solvent. The origin of this observation is probably linked to ion – ion interactions in concentrated organic electrolytes leading to an increase in the energy of the transition state with respect to the initial state in all transport processes described by Eyring’s theory [38]. 4.2.4. Absorption ability

The diffusion coefficients and kinetic absorption constants have been determined for two kinds of liquid electrolytes (VL and VL/EC (80:20) and XLiTFSI  0.08) to which SiO2 is added. The molar percentages of silica in the dry copolymer are 15.7 and 27%, which corresponds to final weight percentages approximately equal to 4 and 8%. The results are reported in Table 13. For reasons that will be

Table 12 Arrhenius activation energies for the diffusion coefficients D and the kinetic absorption constants ka D0 (cm2 s)

Ea,D (kJ mol1)

k0 (g cm2 s1)

Ea,D (kJ mol1)

Pure VL

0.113

21.8

2.9 102

16.2

VL-TFSI (1%)

0.051

20.0

8.0 103

13.6

VL-TFSI (12%)

0.513

25.9

0.87

24.2

Liquid phase

a

Data from Ref. [44].

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Table 13 Diffusion coefficients and kinetic absorption constants with respect to the liquid-phase composition and silica percentage in the copolymer at 20°Ca VL X0(SiO2)b

VL/EC (80/20)

0

0.157

0.270

0

0.157

0.270

2 c

m1,0 *( mg cm )

⬇2

⬇9

⬇ l7

⬇8

⬇ 2l

⬇ 20

105 k (g cm2 s1)

0.67

3.57

3.72

0.53

3.64

⬇ 3.91

0.30

1.51

1.40

0.22

1.54

–

5

2

1

10 D ( cm s ) a

Data from Ref. [44], for liquid phases, the salt molar percentage is 8%. Initial molar fraction of silica in the dry copolymer. c Absorbed weight of liquid phase in the initial stage of superficial wetting of the copolymer by unit surface area of the sample: m1,0*  m1,0 /S0. b

discussed in the next section, the addition of 15.7% SiO2 increases the rate of absorption by a factor 5 (pure VL) or 7 (VL/EC mixture), but beyond this limit, this effect seems to be levelled. The solvent composition has no striking influence on the absorption rate. The results reported in Table 13 clearly show that, at a given temperature, the greater the salt content in the electrolyte, the lower is the final content of liquid phase in the gel. At constant composition of the liquid phase, an increase in the temperature enhances the absorption ability of the copolymer. A plausible explanation of these experimental results is that ion – polymer interactions may prevent the complete swelling of the material [39]. As the temperature rises, the intensity of these interactions decreases, leading to an easier expansion of the polymer network. 4.2.5. Influence of the silica content

To complete this study, Vab, the percentage per volume of absorbed liquid phase in the gel, ΦL, sat, the volumetric percentage of liquid phase in the gel at saturation and Wab, the percentage of absorbed liquid by weight are determined for various silica contents. For this study, the liquid phase is VL/EC (80/20) and X(LiTFSI)  0.08. The results at 20°C are compared in Table 14. As for kinetic constants, the nature of the solvent has little influence on the absorption ability. The most important parameter is the initial volumetric weight, which is directly related to the porosity of the copolymer. The addition of silica increases the absorption speed and ability because it increases the porosity. When maximum porosity is reached, further addition decreases the absorption ability and conductivity by restricting the porous volume (silica itself is not a porous material) and by fixation of some fraction of the lithium salt [40].

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Table 14 Absorption ability of PVDF-HFP/SiO2 (73:27) with respect to the salt molar percentage in VL/LiTFSI solutions and temperature. Vab, percentage of absorbed liquid phase in the gel by volume (deduced from the weight and ρ liq); Φl,sat,volumetric percentage of liquid phase in the gel at saturation; Wab, percentage of absorbed liquid phase by weight a % LiTFSI 0

1

2

4

8

12

Vab (cm3)

233

187

174

170

140

127

Φl,sat

89

86.5

83

80

76

68

Wab at 0°C

302

260

177

10°C

305

283

181

20°C

317

292

30°C

375

302

264

241

226

202 220

a

Data from Ref. [44].

Table 15 Time (in min) of wetting at 20°C required to reach a given advancement of absorption (αab): the liquid is VL/LiTFSI (92:8) X0 (SiO2)a

αab

0

0.157

0.270

0.50

63

12

8.25

0.95

252

67

53

0.99

376

102

83

a

Initial molar fraction of silica in the dry copolymer.

To determine the importance of silica addition to polymer and to permit the comparison of results, it is useful to calculate the advancement of the process at different times, αab, as presented in Fig. 16 for the VL/LiTFSI (92/8) solution at 20°C in copolymer with initial molar fractions of SiO2: X0(SiO2)  0.157 and 0.270. Silica enhances the initial superficial absorption and hastens copolymer saturation. There is no significant difference when the initial silica content is raised from X0(SiO2)  0.157 to X0(SiO2)  0.270. It is useful to know the time required to obtain αab values equal to 0.5, 0.95 and 0.99. The results are given in Table 15. In conclusion, a value of 95% of absorption at saturation provides an

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adequate conductivity for an industrial application and complete saturation, which requires a longer time, has not to be effectively reached. 5. CONCLUSION Fluorine compounds are widely used in the field of Li batteries, most of them are fluoroanions and fluoropolymers. Fluoroanions are components of lithium salts and, at present, cannot be readily replaced by non-fluoro compounds. Fluorine is also involved in the formation of the SEI layer at the carbon electrode and protects the current collectors. Lithium salts with fluoroanions of superacids are soluble in dipolar aprotic solvents, largely dissociated in high permittivity solvents and hence exhibit a high conductivity in solution. PVDF is a fluoropolymer, which presents very interesting properties for use as a conductive membrane in Li batteries and as a binder for electrode materials. From the absorption kinetics and absorption ability of a PVDF-HFP copolymer towards VL/LiTFSI and VL-EC/LiTFSI electrolytes, with or without silica, some important features have been revealed: (i)

In addition to providing better mechanical properties to the conducting membrane, silica increases the porosity of the dry copolymers. Consequently, the volume of the liquid phase in the gel becomes higher and the conductivity increases. This effect is levelled off by the fixation of a fraction of salt on silica, and an addition of 15% (in mol) in the PVDF-HFP copolymer leads to the highest conductivity by the incorporation of liquid electrolytes with a salt molar percentage of about 8%. The faster wetting observed in the presence of silica is another advantage for industrial purposes as the time required for saturation by the liquid electrolyte is reduced four times at 20°C. (ii) Any increase in salt content in the liquid electrolyte leads to the lower absorption ability of the copolymer. This behaviour is related to bonding interactions between ions and polymer chains. As a result, the maximum conductivities for the liquid electrolyte and the gel are not correlated. (iii) Finally, increasing the temperature of wetting increases both the rate and ability of absorption. The temperature of liquid electrolyte absorption can be raised advantageously within the limits of thermal stability of the components. As a matter of fact, this is an irreversible process, meaning that on cooling back to ambient temperature, the gel will not lose some fraction of the retained liquid electrolyte. REFERENCES [1] A.S. Best, P.J. Newman, D.R. MacFarlane, K.M. Nairn, S. Wong, and M. Forsyth, Solid State Ionics, 126 (1999) 191–196. [2] D.R. MacFarlane, J. Huang, and M. Forsyth, Nature, 402 (1999) 792–794.

170 [3] [4] [5] [6] [7] [8] [9] [10] [11] [12]

[13] [14] [15] [16]

[17] [18] [19] [20] [21] [22] [23] [24] [25] [26] [27] [28] [29] [30] [31] [32] [33] [34] [35] [36] [37] [38]

D. Lemordant et al. J. Sun, D.R. MacFarlane, and M. Forsyth, J. Polym. Sci., Part A, 34 (1996) 3465–3470. A. Chagnes, Ph.d. dissertation, Université de Tours, 2000. A. Chagnes, B. Carré, P. Willmann, R. Dedryvère, D. Gonbeau, and D. Lemordant, J. Electrochem. Soc., 150(9) (2003) A1255–A1261. R. Michel, Ph.d. dissertation, Université de Tours, 1996. D. Aurbach, Proc. 12th Int. Meeting Lithium Batteries, ECS, Nara, Japan, 2004, Abstract No. 41. W. Lange and E. Muller, Ber. B, 63 (1930) 1058. K. Syed Mohamed, D.K. Padma, R.G. Kalbandkeri, and A.R. Vasudeva Murthy, J. Fluorine Chem., 23 (1983) 509. R.A. Wiesboeck, US Patent No. 3,654,330, 1972: and US Patent No. 3,907,977, 1975. D.W. Davidson and S.K. Garg, Can. J. Chem., 50 (1972) 3515. D. Lemordant, B. Montigny, A. Chagnes, M. Caillon-Caravanier, F. Blanchard, G. Bosser, B. Carré, and P. Willmann, in: Materials Chemistry in Lithium Batteries, N. Kumagai and S. Komaba (Eds.), Research Signpost, Trivandrum, India, 2002, pp. 343–367. A. Einstein, Ann. Phys., 19 (1906) 289; A. Einstein, Ann. Phys., 34 (1911) 591. M. Ue, J. Electrochem. Soc., 141 (1994) 3336. A. Chagnes, B. Carré, P. Willmann, and D. Lemordant, Electrochim. Acta, 46 (2001) 1783–1791. M. Morita, M. Ishikawa, and Y. Matsuda, in : Lithium Ion Batteries, Fondamental and Performance, M. Wakihara and O. Yamamoto (Eds.), Wiley-VCH, Weinheim, 1998, p. 156, Chap. 7. G. Murphy, J. Chem. Soc. Faraday Trans. (2), 79 (1983) 1607. C. Ghosh, J. Chem. Soc., 113 (1918) 449. I. Ruff, G. Palinkas, and K. Combos, J. Chem. Soc. Faraday Trans. (2), 77 (1981) 1189. K.Hayamizu, Solid State Ionics, 107 (1998) 1. A. Chagnes, S. Nicolis, B. Carré, P. Willmann, and D. Lemordant, Chem. Phys. Chem., 4 (2003) 559–566. N. Bjerrum, Kon. Danske.Videnk. Selskab, 9 (1926) 7. L. Doucey et al., Electrochim. Acta, 44 (1999) 2371. D. Battisti, G.A. Nazri, B. Klassen, and R. Acorsa, J. Phys.Chem., 92 (1993) 5826. M. Lanz and P. Novak, J. Power Sources, 102 (2001) 277–282. D. Aurbach, J. Power Sources, 68(1) (1997) 91. K. Kanamura, H. Tamura, S. Shiraishi, and Z-.Takehara, Electrochim. Acta, 40(7) (1995) 913–921. M. Watanabe, M. Kanba, H. Matsuda, K. Mizogusbi, I. Shinoshara, E. Tscuchida, and K. Tsunemi, Makromol. Chem. Rapid. Commun., 2 (1981) 741. A.S. Godz, C.N. Schmutz, and J.M. Tarascon, US Patent No. 5,296,318, 1994. J.M. Tarascon, A.S. Godz, C.N. Schmutz, F. Shukoki, and P.C. Warren, Solid State Ionics, 49 (1996) 86. J.Y Song, Y.Y. Wang, and C.C. Wan, J. Power Sources, 77 (1999) 183. J.Y. Song, Y.Y. Wang, and C.C. Wan, J. Electrochem. Soc., 147 (2000) 3219. K.M. Abraham, Electrochem. Acta, 38 (1993) 1233. A. Webber, J. Electrochem. Soc., 138 (1991) 2586. S. lylia, J.Y. Sanchez, and M. Arsnand, Electrochim. Acta, 37 (1992) 1699. R. Jasinski, Electrochem. Technol., 6 (1968) 28. B.E. Blomgren, Lithium Batteries, Academic Press, New York, 1983. J.F. Kincaid, H. Eyring, and A.E. Stearn, Chem. Rev., 28 (1941) 301.

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[39] F.M. Gray, Solid Polym. Electrolytes, VCH Publishers , Canibridge, 1991. [40] K.H. Lee, Y.G. Lee, J.-Y. Park, and D.-Y. Seung, Solid State Ionics, 133 (2000) 257. [41] K. Xu, S. Zhang, U. Lee, J. Allen, and R. Jow, Proc. 12th Int. Meeting Lithium Batteries, ECS, Nara, Japan, 2004, Abstract No. 42. [42] Y. Marcus, Ion Solvation, Wiley, New York, 1985, p. 135. [43] D. Brouillette, J. Sol. Chem., 27 (2) (1998) 151. [44] M. Caillon-Caravanier, B. Claude-Montigny, D. Lemordant, and G. Bosser, J. Power Sources, 107 (2002) 125–132.

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 8

Fluorinated anions and electrode/electrolyte stability in lithium batteries Rachid Yazami* and Audrey Martinent† Laboratoire d’Electrochimie et de Physico-Chimie des Matériaux et des Interfaces LEPMI- UMR CNRS-INPG 5631, BP 75, 38402 St Martin d’Hères, France 1. INTRODUCTION Lithium-ion batteries (LIBs) are now the most advanced power sources for portable device applications such as cellular phones, laptop computers, pagers, organizers, etc. Since their first commercialization in Japan by Sony in 1991 [1], LIBs have improved in performance, reliability and safety. This has opened up opportunities for new applications in areas such as transportation (electric and hybrid cars, electric 2-wheel systems), space (satellites, planetary missions’ landers and rovers) [2] and medicine (implantable devices) [3,4]. Fluorine-containing compounds are commonly used in commercial batteries and in those under development. Polyvinylidene fluoride (PVdF) in its homopolymer or copolymer with hexafluoropropene (HFP) forms is used as the binder in the composition of most of the negative (anode) and positive (cathode) electrodes [5]. Lithium hexafluorophosphate (LiPF6) is the most widely used solute in liquid and gelled-type electrolytes [6]. In addition to LiPF6, a sizeable number of inorganic and organic fluorinated lithium salts have been investigated, particularly LiAsF6, LiBF4, LiCF3SO3 and LiN(CF3SO3)2 [7–20]. In polymer membranes, the perfluorinated sulfonate acid membrane (Nafion 117) shows improved film formation ability with PVdF [21]. Other applications of fluorinated compounds in LIBs include electrolyte solvents [22,23], additives [15,24,25] and flame retardants [16,26]. * Current address: California Institute of Technology (CALTECH), The International Associated Laboratory CNRS-CALTECH, Materials for Electrochemical Energetics (ME2), MC 138-78, Pasadena, CA 91125, USA. E-mail: [email protected] †

Current address: Commissariat à I’ Energie Atomique (CEA-Grenoble), 17, Avenue des Martyres, 38041 Grenoble, France. E-mail: [email protected]

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As the solute in the electrolyte composition, a lithium salt should comply with a series of requirements, which are part of a more general set of requirements of any electrolyte in LIB, namely: (i) high ionic conductivity (σi) in a wide temperature range, (ii) extremely low electronic conductivity, (iii) good thermal stability, (iv) good chemical stability versus the anode and the cathode, (v) nontoxicity, and (vi) low cost. The ionic conductivity of an electrolytic solution consisting of n (1  j  n) species each carrying charge zj is expressed by n Fez 2j cj σi  

(1)

1 6πη rj

where F is the Faraday number, e the electron charge, cj and rj are the molar concentration and the Stokes radius of the charged species, respectively, and η is the viscosity of the solution. Optimized conductivity is achieved from a trade-off between concentration and viscosity, as the latter increases with the former. Nevertheless, the solubility and the degree of ionic dissociation of the lithium salt, which determine cj , should be high enough to meet the conductivity requirement. The Stokes radius is fixed by the nature of the lithium salt anion and by the solvent(s), especially by their donor number. Thermal stability is a crucial parameter for battery performance because the battery may be exposed to high temperatures. In fact, the products resulting from the thermal decomposition of the lithium salt may interfere with the electrode reaction or alter the electrolytes physical properties [14,27–38]. For example, in inorganic fluorinated lithium salts with the general formula LiAFp (A P, As, Sb, B, etc.), ionic dissociation (Eq. (2)) is generally followed by thermal decomposition equilibrium as 

LiAFp y Li  AF p 

AF p y F  AFp1

(2) (3)

The presence, even in very small amounts, of the F and AFp1 species in the electrolyte will affect its ionic balance since insoluble LiF and mostly acidic HF may form. Moreover, Lewis, AFp1, either in solution or in gaseous form, may react with the electrodes or corrode the current collectors or the battery hardware in combination with HF [15,38]. It is therefore important that the anion salt be chosen so as to limit the degradation reaction as much as possible in Eq. (3). The primary role of the lithium salt in the electrolyte is to provide a continuum of lithium-ion diffusion and a migration path between the anode and the cathode during the battery charge and discharge operations. If the lithium-ion flux be discontinued, the LIB would stop immediately because the local electric

Fluorinated anions and electrode/electrolyte stability in lithium batteries

175

neutrality condition will no longer be fulfilled. The side reactions to the simple lithium transfer at either electrode will cause the loss capacity of the battery, a phenomenon called self-discharge. In LIBs with a graphitic carbon anode, a passivation film solid electrolyte interphase (SEI) grows on the lithiated graphite surface during the first lithium intercalation/de-intercalation cycles. The SEI plays a major role in maintaining battery stability as it significantly impedes thermodynamically predicted reactions between the graphite and the electrolyte. Such a reaction consumes lithium and results in the capacity loss [39,40]. The effective surface protection by the SEI depends on the electrolyte composition, namely the lithium salt and the solvents [9,16,17,20,23,32,35,38,40–44]. The electrolyte reacts readily with the cathode, particularly at high-voltage charge states and at high temperatures [25,38,45]. Despite intensive research on alternative lithium salts, LiPF6 remains the most widely used salt in commercial LIBs applications. The main reason is that LiPF6 gives the best compromise between all criteria cited above, especially high conductivity and an acceptable stability versus the anode and cathode. However, LiPF6 is expensive and highly sensitive to humidity, which poses some handling and storage issues. The way in which the spontaneous hydrolysis of LiPF6 occurs and has an impact on LIB behavior is an important scientific and technological topic, which has not been thoroughly covered in the literature, despite some good but limited studies [35,42,45–50]. This paper will attempt to fill the gap regarding the kinetics of LiPF6 hydrolysis, identification of the hydrolysis products and its effects on graphite electrode behavior [51,54,55]. 2. THE HYDROLYSIS OF LiPF6 2.1. Literature survey

The hydrolysis of LiPF6 occurs either in the solid state [52] or in solution, such as in an organic solvent [42,50–55]. In the solid-state reaction, assuming the hydrolysis follows the scheme: LiPF6  H2O y LiF  POF3  2HF

(4)

and is based on pressure vs. time measurements, Barlow found a good fit with the following kinetics law: 1 1 t   ln(a)   ln (a  0.5x) k k

(5)

where k is a time constant (⬃ 2  105 s1 at room temperature), a the initial number of moles of water and x the number of moles of HF [52]. The fit applies, however, only during the first 36 min. For longer reaction times, the hydrolysis

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reaction interfered with the HF attack of the silica glassware, resulting in the formation of gaseous SiF4: 4LiPF6  8H2O y 4LiF  4HO2PF2  12HF

(6)

12HF  3SiO2 y 6H2O  3SiF4

(7)

Accordingly, Barlow’s study suggests that in addition to HF and LiF, two hydrolysis products of LiPF6 form, namely POF3 and PF2OOH. However, no direct evidence of the formation of those products was shown in this study. When hydrolysis was carried out in a mixture of ethylene carbonate (EC) and dimethyl carbonate (DMC) solution, a different kinetics behavior was found from the pressure measurements. 19F- and 31P-nuclear magnetic resonance spectrometry (NMR) showed the formation of PF2OOH but not that of POF3. Another mechanism of LiPF6 hydrolysis in solution was proposed by Aurbach et al. [9], which involves thermal decomposition to LiF and PF5 [9] according to LiPF6 y LiF  PF5

(8)

PF5  H2O y 2HF  POF3

(9)

In the case of excess water in the solution, hydrolysis proceeds in steps to the formation of orthophosphoric acid (H3PO4) as suggested by Ishikawa et al. [56]: LiPF6  H2O y LiF  POF3  2HF

(4)

POF3  H2O y POF2 (OH)  HF

(10)

POF2 (OH)  H2O y POF(OH)2

(11)

POF(OH)2  H2O y H3PO4  HF

(12)

2.2. Recent studies on LiPF6 hydrolysis 2.2.1. Introduction

The focus in this section is on the changes in the physicochemical characteristics of LiPF6-based electrolyte solutions during hydrolysis and on their effects on the electrode behavior of the lithiated graphite anode for lithium-ion battery applications. The kinetics study includes Karl – Fisher chemical water analysis, conductivity measurements, 19F-NMR spectroscopy and electrochemical studies in lithium half-cells.

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177

2.2.2. Experimental procedure

Powder samples of LiPF6 were provided courtesy of Stella Company (former Hashimoto Co.), Japan. A molar solution was made using an equivolumic solvent mixture of EC and dimethyl carbonate DMC. The Karl – Fisher titration showed the initial water content of the solution to fall below 10 ppm. Table 1 gives the powder and solution composition determined by 19F-NMR analysis. Titrated amounts of distilled water were added to the solution so that the initial added water content was 100, 500, and 1000 ppm. The analysis before and after water contamination included the following: ●

●

●

●

Karl – Fisher water analysis using the Radiometer Copenhagen’s Aquaprocessor system in an argon-filled dry box. Analyses were performed on samples immediately after the addition of water and at regular times thereafter. Conductivity measurements in the temperature range 5° to 60°C using the Radiometer LF 330 conductimeter and a four-electrode Tetracon® 325 Pt cell. 19 F-NMR spectroscopy using the Brucker DRX 400 spectrometer operating at 376.4 MHz on CD3CN solutions using CF3COOH as the internal reference. Electrochemical tests on Li/electrolyte/graphite half-cells including galvanostatic charge/discharge cycling, slow scan voltammetry and electrochemical impedance spectrometry.

The conductivity and the 19F-NMR measurements were performed on water-contaminated solutions aged at 60°C for 1 week. For the electrochemical measurements, the solutions were used in the lithium half-cells aged for about 1 week after their preparation at ambient temperature. 2.2.3. The kinetics study

The time dependence of the water content in the 100, 200, and 1000 ppm water-precontaminated LiPF6 EC–DMC solutions is shown in Fig. 1. For the first 24 h the data points fit a first-order kinetic law expressed by [H2O](t) [H2O](0)exp(λt)

(13)

Table 1 Molar composition of LiPF6 in the solid state and in the electrolyte solution before hydrolysis Molar content (%)

LiPF6

POF2OH

HF

POF3

Solid state

99.93

0.06

0.00

0.01

EC–CDM solution

99.79

0.1

0.11

0.01

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1200

1000 ppm 500 ppm 100 ppm

[H2O]/ ppm

1000 800 600 400 200 0 0

48

96

144

192

240

288

336

t/ h

Fig. 1. Time dependence of the water content in 1 M LiPF6 in EC–DMC electrolyte solutions upon adding 100, 500, and 1000 ppm of water.

The time constant λ is close to 105 s1, half the value reported by Barlow in the case of solid-state reaction [51]. This supports Barlow’s conclusion that the LiPF6 hydrolysis reaction is slower in solution than in the solid state. The differences in kinetics should relate to the fact that the hydrolysis reaction is highly exothermal, which, in the case of the solid-state reaction, may affect the local temperature of the reactants, whereas in solution, the temperature equilibration is reached faster due to mechanical stirring. This may also be at the origin of the differences in the time during which the first-order law applies in the solid state and in solution (36 min and 24 h, respectively). The hydrolysis reaction may follow different mechanisms in solid state and solution as POF2OH forms only in the latter as discussed in the next section. Our finding of a first-order kinetics suggests that the mechanism of hydrolysis is governed by Eq. (4), as the reaction is limited by much smaller amounts of water compared with that of LiPF6 ( 1000 ppm vs. 1 mol.l1, respectively). For longer reaction times, Eq. (13) does not apply anymore, which indicates that other reactions should be involved such as Eq. (10). In fact, the 19F-NMR analysis performed after hydrolysis, does not show the presence of POF3, whereas POF2OH could be detected. Fig. 2 shows the details of the LiPF6, HF and POF2OH composition for each electrolyte. The HF/POF2OH ratio in the waterpre-contaminated electrolytes is slightly higher than 3, which supports the point that the hydrolysis does not progress beyond the formation of POF2OH in Eq. (10). Therefore, the progress of the reactions in Eqs. (11) and (12) should be very limited, which is the reason why POF(OH)2 and H3PO3 were not detected by the 19 F-NMR. Moreover, the amounts of POF3 remained below 0.03%, which indicates that the hydrolysis reaction of Eq. (10) is nearly complete.

Fluorinated anions and electrode/electrolyte stability in lithium batteries

179

Fig. 2. Effect of added water on the solutions composition from 19 F-NMR analysis after equilibriation at ambient temperature.

The results shown in Fig. 2 were achieved on electrolyte solutions maintained for several weeks at ambient temperature in hermetically sealed tubes filled with dry argon (⬃1 ppm H2O). To check the effect of temperature on the equilibrium state, the solutions were heated to 60°C in argon, maintained for 1 week, and then analyzed by 19F-NMR. The results are displayed in Table 2 together with those achieved before thermal aging. The solution composition changed significantly upon aging; the amounts of LiPF6 decreased and those of HF and POF2OH significantly increased. Yet, the composition in POF3 remained very low and no evidence of POF(OH)2 and H3PO3 formation was found from the 19P- and 31P-NMR analyses. The thermal decomposition of LiPF6 to LiF and PF5 according to Eq. (8), followed by the hydrolysis of PF5 as described in Eq. (9), should account for the changes in the solution composition. It suggests that PF5 hydrolyzes more readily than PF6 does, generating more HF and POF2OH as a result. It is worth noting that the largest relative changes in the HF and POF2OH composition occur in the water non-contaminated solution. HF generation is believed to be at the origin of the premature capacity fade in LiMn2O4-based cathodes [57]. The small amount of water in the non-contaminated solution suggests that a mechanism other than hydrolysis may account for HF formation, such as the PF5 acid – base-type reaction with the electrolyte solvents. The HF generation “in situ” is of great importance in lithium-ion batteries technology and deserves further investigation.

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Table 2 Thermal aging effects (60°C for 1 week) on composition of LiPF6 electrolyte solutions with and without added water Water added (ppm)

LiPF6 (mol%)

POF2OH (mol%)

HF (mol%)

POF3 (mol%)

0

94.07 (99.79)

1.13 (0.1)

4.77 (0.11)

0.03 (0.03)

100

91.71 (97.37)

1.65 (0.59)

6.61 (2.03)

0.03 (0.03)

500

86.82 (92.43)

2.64 (1.78)

10.52 (5.78)

0.02 (0.02)

1000

79.6 (87.33)

3.91 (2.96)

16.47 (9.69)

0.02 (0.02)

2.2.4. Effect on conductivity

Conductivity measurements were performed on the electrolyte solutions before and after water contamination between 5°C and 60°C. The results were fit with the Vogel – Tamman – Fulcher (VTF) equation Ea σ (T)  σ 0 exp  R(TT0)

冢

冣

(14)

Fig. 3 shows the temperature dependence of conductivity in the water-noncontaminated solution (a) and in those contaminated with 100 ppm (b), 500 ppm (c) and 1000 ppm (d). Table 2 shows the fit parameters σ0, B and T0 from Eq. (14). As expected, the preexponential term σ0 decreases, on average, with the water content. Concomitantly, T0, which is the equivalent to a glass transition temperature, increases and the activation energy Ea decreases. The latter result reflects the more pronounced effect of temperature on conductivity on adding water. The VTF equation derives from the relationship between relaxation times and 1/T until viscosity (the reciprocal of mobility) approaches infinity at T0 [58]. Higher T0 may accompany increased solution viscosity as more species form from LiPF6 hydrolysis. Therefore, water contamination is detrimental to the ionic conductivity of the LiPF6-based solution. The formation of neutral and poorly soluble species such as POF2OH, POF3 and LiF should be at the origin of the overall decay in the ion transport characteristics. A T0 of 157 K in the water noncontaminated electrolyte is in agreement with the result reported by Stallworth et al. of 160 K in their 1M LiPF6 EC-PC solution [59]. 2.2.5. Effect on electrochemical behavior 2.2.5.1 Introduction Lithium-ion batteries use graphite as the lithium intercalation

material for the negative electrode (anode) and mixed oxides such as LiMO2 (M  Co, Ni, Mn, V, etc.) for the positive electrode (cathode) [6]. During charge and discharge, lithium shuttles between the negative and positive electrodes via the

Fluorinated anions and electrode/electrolyte stability in lithium batteries

181

Fig. 3. Temperature dependence of the 1 M LiPF6 EC–DMC electrolyte solutions contaminated with 0 (a), 100 (b), 500 ppm (c), and 1000 ppm (d) of water . The data points were fitted with the VTF equation (14). The bold data points correspond to results at room temperature.

electrolyte. A high open-circuit voltage of around 4 V arises due to the difference in the chemical potential of lithium in graphite and in LiMO2. Actually, the graphite electrode operates at very low potentials that are only 100 to 200 mV above those of metallic lithium. Thermodynamics predicts that any organic solvent or most multicomponent anions will react with lithiated graphite to form reduction products. The “kinetics” stability of the anode is attributed to the formation of a passivation layer on the surface of the graphite grains called the SEI. The latter forms during the first charge/discharge cycle of the battery during which lithium intercalates and de-intercalates into or from graphite [40]. The SEI grows as a result of the reductive decomposition of the electrolyte solvent and lithium salt. Such a reaction consumes lithium irreversibly and should therefore be as limited as possible in order to achieve a high capacity of the lithium-ion battery.

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In the following section, we will evaluate the effect of adding water to the electrolyte solution with regard to the irreversible capacity loss, electrode kinetics and, more particularly, electrode/electrolyte interface properties. 2.2.5.2 Voltammetry The slow scan voltammetry technique consists of applying a linear sweep to the graphite electrode potential from the starting open-circuit potential (OCV) of ⬃3 V to 5 mV against metallic lithium. The rate of potential sweep is fixed at a low value of 5 μV. s1 to achieve full lithiation and delithiation of the graphite electrode. The cells consist of coin-type half-cells with metallic lithium, LiPF6 EC-DMC solutions with and without added water (as described above) as the electrolyte, a polypropylene-based microporous separator and a graphite composite electrode. The cells were operated at ambient temperature for several cycles. The voltammograms during the first cathodic polarization of the graphite electrode in electrolytes with 0 (a), 100 (b), 500 (c) and 1000 ppm (d) of added water are shown in Fig. 4. The voltage scan covers the 2.5–0.5 V area, which we will refer to as the “high-voltage” area. Fig. 5a–d shows the continuation of the previous voltammograms in the “low-voltage” area ranging between 0.3 and 0.005 V, followed by de-intercalation/re-intercalation cycles in the same voltage range. The voltammograms in Fig. 4 show a major current peak reduction in the 0.7–0.8 V range. The three cells with water-added electrolytes Figs. 4 (b–d) display an additional peak around 2.2 V and the cells with 500 and 1000 ppm water have a current shoulder at around 1.9–2.0 V (Fig. 4c and d, respectively). The major peak at 0.7–0.8 V is typical of the first cathodic reduction of graphite and is associated with SEI formation. The latter is caused by the electrolyte reduction, involving that of the solvents and LiPF6. This peak does not have a corresponding anodic peak and does not appear in the subsequent cycles, which shows its irreversible character. The fact that the major peak is present even in the watercontaminated electrolytes at nearly the same peak voltage and the same normalized intensity (mA g1 of graphite) indicates that electrolyte reduction is not affected by the presence of water or by LiPF6 hydrolysis. However, the peak at ⬃2.0 V and the shoulder when present at 1.9–2.0 are characteristic of water-contaminated electrolytes and should be associated with proton reduction (Table 3). In fact, a good correlation is found with the area under the peak at 2.0 V and with the amount of water as determined by the Karl – Fisher analysis. Slow scanning rate voltammetry proved to be a new convenient and accurate method to probe the water content in electrolyte solutions above ⬃20 ppm. The shoulder at 1.9–2.0 V is less well defined and is difficult to use for analytical purposes. It is most likely due to the HF reduction. The integration of current peaks in the “high voltage” area [i.e. from the initial rest voltage (OCV0 to 0.5 V) allows a quantitative evaluation of the amounts of charge used in each voltage domain. Table 4 reports the corresponding irreversible capacity losses together with the equivalent amount of protons consumed during the first high-voltage peak (OCV0 to 1.5 V).

0

0

-1

-1

0 ppm H2O

-2

-3

-3

I/mA g–1

-2

-4 -5

-6 -7 0.8

1.2

(a)

1.6

I/mA g–1

-8

2.4

0.5

0

0 -1

-2

-2 500 ppm H2O

-4 -5

-7

E/ V vs. Li+/Li

-8 0.5

2.5

(d)

1000 ppm H2O

-5

-7 2

2.5

-4

-6

1.5

2

-3

-6

1

1.5 E/ V vs. Li+/Li

-1

-8 0.5

1

(b)

E/ V vs. Li+/Li

-3

(c)

2

100 ppm H2O

-5

-7 -8

183

-4

-6

I/mA g–1

I/mA g–1

Fluorinated anions and electrode/electrolyte stability in lithium batteries

1

1.5

2

2.5

E/ V vs. Li+/Li

Fig. 4. First cathodic polarization voltammogram of Li/LiPF6 EC–DMC  x ppm H2O/graphite half-cells under 5 μ V s1 sweeping rate in the 2.5–0.5V range (a) x  0, (b) x 100, (c) x  500 (b), and (d) x 1000 ppm.

The analysis of results in Table 4 leads to the following conclusions: ●

● ●

Adding 100 ppm of water surprisingly seems to improve the electrode performances as far as irreversible capacity is concerned. The capacity loss due to the SEI formation is not affected by the water content. The excess in capacity loss is due to the proton reduction in water-containing electrolytes, especially above 100 ppm.

Since no significant current feature is observed in the 0.5 – 0.3 V voltage area, this area was omitted in the graphic representations. As shown in Fig. 5(a), a series of reduction and oxidation peaks appear in the voltammograms below 0.3 V. In the first cycle, four reduction peaks, three of which are well defined, appear in the non-contaminated electrolyte at ⬃ 0.200, 0.145 (weak), 0.100, and 0.0695 V, respectively. These peaks are characteristic of stage transitions in well-crystallized graphite materials [60]. At first dilute stage 1 is formed, which

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Table 3 Fits parameters with a range of the VTF equation (13) for the 1M LiPF6 EC–DMC electrolyte solutions contaminated with water H2O (ppm)

σ0 (mS cm1)

Ea (kJ.mol1)

T0 (K)

0

100

500

1000

0.17

0.14

0.13

0.11

0.14–0.20

0.11–0.16

0.12–0.15

0.04–0.16

3.23

2.75

2.65

2.13

2.83–3.64

2.37–3.13

2.41–2.90

1.01–3.25

157

167

169

184

148–166

158–176

163–175

154–213

Table 4 Effect of amounts of water added to the electrolyte on the irreversible capacity losses on the graphite electrode Water added (ppm)

Irreversible capacity (mAh g1) Proton reduction (OCV0–1.5 V)

82

6.4

75.6

200

100

79.5

6.5

73

200

500

99

26.5

72.5

966

1000

120

46

74

1650

0

SEI formation (1.5–0.6 V)

Equivalent protons consumption (ppm)

Total (OCV0–0.6 V)

consists of randomly distributed Li between every two graphene layers. Dilute stage 1 then transforms to the stage 4 compound (⬃Li0.13C6), which corresponds to the 0.2 V peak. Stage 4 in turn transforms to stage 3 (⬃Li0.18C6) and then to stage 2L (⬃Li0.23C6) in the 0.15– 0.14 V range. Then disordered stage 2L converts to ordered stage 2 (Li0.5C6) at ⬃0.1 V and finally stage 2 transforms to stage 1 LiC6, which gives rise to the largest peak at ⬃0.0695 V. When the voltage scan is reversed, oxidation peaks appear at 0.110, 0.145, 0.170 (small), and 0.227 V, respectively. The oxidation peaks are associated one to one with the reduction peaks discussed above as the stage transformations in graphite are highly reversible. In the water-contaminated electrolytes, the reduction peaks broaden significantly and shift to lower and higher voltages during reduction and oxidation,

Fluorinated anions and electrode/electrolyte stability in lithium batteries

25

30 0 ppm H2O 20

1st

cycle 2nd cycle

20

100 ppm H2O

15

1st cycle 2nd cycle

10

I/mA g–1

10

I/mA g–1

185

0 -10

5 0 -5

-20 -30

-10 -15 0

0.05

0.1

(a)

0.15

0.2

0.25

0.3

0

(c)

0.1

0.15

0.2

0.25

0.3

E/ V vs. Li+/Li

25

20

500 ppm H2O

20

15

1st cycle 2nd cycle

15

10 5 0

0 -5

-10

-10 0.05

0.1

0.15

0.2

E / V vs. Li+/Li

0.25

0.3

1st cycle 2nd cycle

5

-5

0

1000 ppm H2O

10

I/mA g–1

I/mA g–1

25

-15

0.05

(b)

E / V vs. Li+/Li

-15

(d)

0

0.05

0.1

0.15

0.2

0.25

0.3

E/ V vs. Li+/Li

Fig. 5. First and second cyclic voltammograms of Li/LiPF6 EC–DMC  x ppm H2O/graphite half-cells under 5 mV s1 sweeping rate in the 0.3–0.005 V range: (a) x  0 ppm, (b) x  100 ppm, (c) x  500 (b), and (d) x  1000 ppm.

respectively. The peak shift is the signature of higher overall cell polarization. Surprisingly, the oxidation peaks are better defined than the corresponding reduction peaks in the first cycle. This feature persists in the second cycle during both reduction and oxidation. Reduction peak polarization also decreases between the first and the second cycles. The enhanced definition of the current peaks and reduced cell polarization between the first reduction and oxidation scans and during the second cycle suggests that the SEI may not have fully formed in the “high-voltage” area shown in Fig. 4 a–c. It also suggests that some of the hydrolysis reaction products may adsorb on the graphite surface during lithium intercalation and then re-dissolve in the electrolyte accompanying the first lithium de-intercalation. This adsorption/dissolution mechanism occurs mainly during the first cycle in the water-contaminated solutions due to the delay in SEI formation. In the second cycle the SEI is better stabilized, which allows the lithium intercalation and de-intercalation to proceed more normally, giving rise to typical staging peaks and lower cell polarization.

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The effect of adding water to the electrolyte solution on the reversible capacity of the graphite electrode during the first and second cycles was also investigated. Table 5 displays the discharge (Li “intercalation”) and charge (Li “de-intercalation”) capacity in different electrolytes. During the first cycle, the discharge capacity increases with the water content in the solutions and the charge capacity decreases. However, the relative increase in discharge capacity is 6.5% between 0 and 1000 ppm added water, whereas the decrease in the corresponding charge capacity is only 2.85%. The difference probably is due to the irreversible character of proton reduction in the “high-voltage” area. During the second cycle, the discharge capacity is lower than in the first cycle, but it still stands higher than the charge capacity, which confirms that the SEI is still not stabilized. The charge capacity in the second cycle is remarkably high and is almost the same for all electrolytes (around 350 mAh g1). In conclusion, adding up to 1000 ppm of water to the electrolyte will negatively affect the coulomb efficiency of the first cycle and to a lesser extent that of the second cycle due to proton reduction. The presence of water and hydrolysis products delays SEI formation. Surprisingly it does not affect the discharge capacity of the graphite electrode, at least under slow scan voltammetry. 2.2.5.3. Electrochemical impedance spectrometry (EIS) EIS is an appropriate technique for studying electrode/electrolyte interfacial properties and their evolution with electrode potential. From the voltammetry study in the previous section, we can distinguish two voltage areas during the first cathodic polarization of the graphite electrode. The first one covers a voltage range between the initial OCV and 0.5 V, where the reactions at the electrode are irreversible (proton and electrolyte reduction). The second area ranges between 0.3 and 0 V and corresponds to the lithium intercalation. In the 0.5–0.3 V area no significant faradic process takes place.

Table 5 Effect of amounts of water added to the electrolyte on the discharge and charge capacity of the graphite electrode during the first two cycles

Water added (ppm)

First cycle OCV0 → 0.005 V → 0.3 V Qdis. (mAh g1)

Qchar. (mAh g1)

Second cycle 0.3 V → 0.005 V → 0.3 V Qdis. (mAh g1)

Qchar. (mAh g1)

0

538

350

370

350

100

548

340

376

350

500

558

316

370

353

1000

573

340

363

352

Fluorinated anions and electrode/electrolyte stability in lithium batteries

187

In order to further characterize the kinetic processes in these two areas and the effect of water in the electrolyte, we performed EIS measurements on specially designed three-electrode coin cells. The cells use two metallic lithium electrodes, the reference and counterelectrodes, the third electrode is the graphite-based working electrode. Fig. 6a–d shows a typical chart of the Nyquist plot evolution during the first cathodic polarization of the graphite electrode in the OCV0–0.5 V voltage area obtained with LiPF6 EC-DMC electrolytes having different amounts of added water. In the electrolyte with no added water, Fig. 6a shows a high-frequency (HF) semi-cycle appearing on the left-hand side of the Nyquist plot even at OCV0. At voltages below ⬃1 V, a depressed semi-circle at medium frequencies (MF) starts to show, decreasing in diameter with increased cathodic polarization. A low frequency arc is also present with decreasing phase angle from near 90° at

Fig. 6. Evolution of the Nyquist plot of the graphite electrode during the first cathodic sweep in the 2 V (or 1.5 V)-0.5 V range, in Li/LiPF6 EC–DMC  x ppm H2O/graphite half-cells, (a) x  0 ppm, (b) x  100 ppm, (c) x  500 (b), and (d) x  1000 ppm.

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OCV0 (blocking electrode) to close to 45° at 0.5 V (semi-infinite diffusion, Warburg behavior). The size of the HF semi-circle remains almost unchanged in the voltage range 1.5 – 0.5 V. In the water-containing electrolytes, the shape of the Nyquist plots is quite different from the previous one. In the 100 ppm water-containing electrolyte in Fig. 6b, the most visible difference is a larger HF semi-circle denoting higher impedance. In the 500 ppm water-contaminated electrolyte of Fig. 6c, a more depressed HF semi-circle at 2.5 and 2 V evolves into an inductive-type loop below 2 V. The inductive behavior may correlate with the first reduction peak shown in Fig. 4c, which we attribute to proton reduction. It is not ruled out that reduction involves a step where the proton adsorbs on the graphite surface. Adsorption is one known cause of an inductive loop in the Nyquist plot. This unusual behavior is also observed in the electrolyte solution with 1000 ppm added water as shown in Fig. 6d. Here, we extended the cathodic sweep to 0.3 V to show the continuous change in the inductive loop with cathodic polarization. At this point we have no general theory to deal with the Nyquist plot shape in water-contaminated electrolytes for quantitative analysis. The appearance of an inductive loop, however, is visible evidence of water contamination. Therefore, EIS confirms the differences in electrode behavior due to LiPF6 hydrolysis. In Fig. 7a–d, we show the evolution of Nyquist plots in the 0.2– 0 V area where lithium intercalation into graphite is expected to occur. With the water non-contaminated electrolyte, the overall shape of the Nyquist plots does not vary much, with the HF and MF semi-circles and the low frequency arc at about 45°. All the water-contaminated electrolytes show an inductive loop, including the one with 100 ppm water for which such a loop was not observed in the OCV0– 0.3 V voltage range (see Figs. 7b and 6b, respectively). In the 500 and 1000 ppm water-contaminated electrolytes, the inductive loop observed at higher voltages remains in the lower voltage area of 0.2– 0 V. The loop, however, decreases in size as the applied voltage approaches 0 V. Since in the 0.2 – 0 V areas, lithium intercalation peaks are observed (Fig. 5b–d), the persistence of the inductive behavior suggests that lithium intercalation occurs even when a proton is adsorbed on the graphite surface, assuming adsorption is at the origin of the inductive loop. It is therefore very likely that lithium deposition/insertion is concomitant with hydrogen adsorption/reduction in the “low-voltage” area. The following may account for the complex reaction mechanism involving the hydrolysis of LiPF6: proton adsorption (Eq. (17)), lithium adsorption and reduction (Eq. (18)), and proton reduction in the adsorbed state (Eqs. (19) and (20)). Hydrogen evolution and the intercalation of lithium into graphite is described in Eq. (21): 

LiPF6 y Li  PF6 



PF 6  2H2O y F  POF2 (OH)  3HF

(2ⴕ) (4  10)

Fluorinated anions and electrode/electrolyte stability in lithium batteries

189

Fig. 7. Evolution of the Nyquist plot of the graphite electrode during the first cathodic sweep in the 0.2 – 0 V range, in LiPF6 EC– DMC electrolytes with 0 ppm (a), 100 ppm (b), 500 ppm (c), and 1000 ppm (d) of the water.



HF y H  F 

(15)

Li  F y LiF

(16)

εH  C y C(εH)ads. x xε x C(ε H)ads.   Li   e y C(εH,  Li)ads... 6 6 6 x xε C(ε H,  Li)ads. y C(εLiH,  Li)ads. 6 6 1 xε  Lix C6 (3εH2)ads. C(ε LiH,  Li) y ads. 6 6

(17)

Lix C6 (3ε H2)ads. y Lix C6  3εH2 ↑

(21)

(18) (19) (20)

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The complex chemistry sketched in the above equations could be the origin of the observed delay in the SEI formation in water-containing electrolytes as evidenced by the voltammetry and the EIS studies. Actually, we found that the inductive behavior persisted during the lithium de-intercalation in all watercontaminated solutions. This is illustrated in Fig. 8a–d, which shows the evolution of the Nyquist plots during the first intercalation/de-intercalation cycle, starting from 0.4 V (a) down to 0 V (b), then back to 0.05 V (c) and 0.2 V (d). The electrolyte solution here is with the 1000 ppm H2O. The data point frequency in Fig. 8 is given as powers of 10 (i.e. the data point bearing number 4 corresponds to 104 Hz).

7

20

0 V vs. Li+/Li

6

+

0.4 V vs. Li /Li

15

5

2 3 -2

5

4

–Im(Z)/

–Im(Z)/

10

-2

3 3

2 -1

0

-5

-1

4

1

1

0

2 0

0 0

5

10

15

20

Re(Z)/Ω

(a)

1

25

2

3

4

6

7

8

Re(Z)/Ω

(b)

4

5

2 -2

+

0.05 V vs. Li /Li +

0.2 V vs. Li /Li

1.5 3

-2

–Im(Z)/

–Im(Z)/

1

2

-1 3

0.5

0

0 3

1

4

4

-1

2

-0.5

1

1 2

0

0 1

(c)

2

3

Re(Z)/Ω

4

-1 1.5

5

(d)

2

2.5

3

3.5

4

4.5

Re(Z)/Ω

Fig. 8. Evolution of the Nyquist plot of the graphite electrode during the first intercalation/deintercalation cycle in the in LiPF6 EC–DMC electrolyte with 1000 ppm added water. The spectra were taken at 0.4 V (a) and 0 V (b) during intercalation and at 0.05 V (c) and 0.2 V (d) during de-intercalation. The frequency of the selected data points is given in powers of 10 Hz.

Fluorinated anions and electrode/electrolyte stability in lithium batteries

191

The large inductive loop in Fig. 8a at 0.4 V decreases steadily in size as the electrode potential approaches 0 V (Fig. 8b), and almost disappears at 0.05 V. When the voltage sweep is reversed it reappears at 0.2 V, although with smaller size. The adsorption of hydrogen (Eq. (17)) takes place when the electrode potential becomes more positive as at a lower potential it competes with the lithium intercalation, the dominant electrode reaction mechanism. It is interesting that the inductive loop in the 100 ppm water-contaminated electrolyte appears at lower voltages than in the 500 and 1000 ppm solutions. This suggests that the inductive behavior becomes observable only when the amount of adsorbed hydrogen is sufficiently large. Driving the electrode potential toward more negative voltages may favor the accumulation of adsorbed hydrogen in the less contaminated solution of 100 ppm H2O. The EIS proves to be very sensitive to LiPF6 hydrolysis and could be used as a technique to probe electrolyte purity. Moreover, since the major source of proton after LiPF6 hydrolysis comes rather from HF than from unreacted H2O (compare Figs.1 and 2), the EIS is a good probe for HF impurity in the solution. 2.3. Conclusion

The kinetics of the LiFP6 hydrolysis was followed by 19F-NMR analysis, which makes it possible to analyze the hydrolysis reaction products both qualitatively and quantitatively. The amounts of POF3 in the water-added solutions remained very small compared with HF and POF2OH, which indicates that the hydrolysis reaction proceeds beyond Eq. (4) to cover Eq. (10). Since we have found no spectrometric evidence of POF(OH)2, it is the author’s argument that POF2OH is stable in water; therefore, the hydrolysis reaction described in Eq. (11) does not take place. Moreover, the HF/POF2OH molar ratio is close to 3, which supports the conclusion that hydrolysis does not proceed beyond POF2OH. The conductivity measurements show evidence of the decreased ion transport properties of water-contaminated LiFP6 electrolyte solutions. We suggest that the hydrolysis products, which very likely are poorly soluble, will increase the solution viscosity and therefore decrease the ionic conductivity. Our electrochemical measurements show very high sensitivity to water contamination even at low (100 ppm) water addition. The irreversible capacity increases with the water content, which delays the SEI formation. Slow scan voltammetry proved to be an alternative probe for HFH2O content in the electrolyte. Surprisingly, the reversible capacity reached after the first cycle was not affected by the addition of up to 1000 ppm water to the electrolyte. EIS measurements were found to be more sensitive to LiPF6 hydrolysis. The inductive behavior that we associate with proton reductive adsorption gives the strongest evidence for electrolyte water contamination. 3. GENERAL CONCLUSION In this paper, we focused on one of the important issues related to the use of LiPF6 as the solute for electrolytes in lithium batteries. There is abundant scientific

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literature on the effects of water contamination on half- or full-cell behavior. However, the kinetics of hydrolysis has not been covered comprehensively enough, a gap this paper attempts to fill. The literature stresses the need for very high-purity LiPF6 when considering lithium battery applications. Water and HF are the most important impurities in LiPF6-based electrolytes. Many patents were awarded in the United States to cope with the purification issue, including LiPF6 synthesis in anhydrous organic solvents [61–63], removing HF with a weak base resin [64,65], highly anhydrous HF [66,67], and preparation of solvate complexes of LiPF6 [61,68,69]. Indeed, as we have shown in this work, the electrode/electrolyte interfacial properties are significantly affected by the presence of impurities. This has been discussed in the literature regarding the spinel LiMn2O4, which is considered as one of the best candidates to replace the currently used LiCoO2 positive electrode in rechargeable lithium batteries [70]. It is surprising that, despite very intensive research among the lithium battery community, no substitute for LiPF6 has been implemented so far in commercial lithium-ion batteries. This highlights the outstanding properties of LiPF6 as the best compromise material in terms of conductivity and stability with anode and cathode. However, the cost and thermal stability still need to be improved. ACKNOWLEDGEMENTS The authors acknowledge the scientific and experimental support to carry out the electrochemical measurements from their colleagues at LEPMI, Prof. Claude Montella, Prof. Gérard Le Gorrec, and Prof. Jean-Paul Diard. They also acknowledge the financial and technical support (19F-NMR) of AtoFina, S. A. (former Elf-Atochem S. A.), France. REFERENCES [1] [2]

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

195

Chapter 9

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy) aluminate superweak anions Benjamin G. Nolan,a Shoichi Tsujioka,b and Steven H. Straussa a

Department of Chemistry, Colorado State University, Fort Collins, CO 80523, U.S.A. b

Central Glass Company, Tokyo, Japan

1. INTRODUCTION Many lithium salts of molecular anions have been considered as potential ionic components of electrolytes for primary and secondary lithium-ion batteries. Simple salts such as LiClO4, LiBF4, LiPF6, LiAsF6, and LiCF3SO3 fail to meet one or more of the accepted performance criteria [1,2]. Much more promising (although more costly) are various perfluoroalkylsulfonyl salts such as LiN(CF3SO2)2 and LiC(CF3SO2)3 [3]. A growing class of new lithium salts contain a variety of B(O២O)2 anions, where the chelating O២O2 dioxo dianions impart hydrolytic and thermal stability to the borate anions and, hence, to their salts [4–7]. Many of these new borate anions contain two or more fluorine atoms, a strategy that improves the electrochemical stability and generally increases the conductivity at a given salt concentration. The seminal work in this area, a study of the electrochemical properties of LiB(1,2-C6H4O2)2, was reported by Barthel and co-workers in 1995 [4]. In this chapter, we review our recent work on the electrochemical properties of nine electrolyte salts based on the formula LiB(OC(2-O-C6H4nFn)(CF3)2)2 (n  0–3) [5] and, for comparison, the six related electrolyte salts LiAl(OCR(CF3)2)4 (R  H, Me, CF3, Ph), LiAl(OCH2CF3)4, and LiAlF(OCPh(CF3)2)3 [6]. Some of the aluminate anions are among the most weakly coordinating (i.e. superweak [7–9]) anions studied to date [10–12], and the salt LiAl(OCH(CF3)2)4 has recently been used as a component of a nanocomposite electrolyte [13]. (The simplest definition of a superweak anion is that it is the conjugate base of a neutral superacid.)

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The search for a practical replacement for LiPF6 in lithium batteries is one of the goals of our work. Equally important is the understanding of the steric and electronic factors that lead to (i) high conductivity at high salt concentrations, (ii) high thermal stability, and (iii) improved electrochemical stability (both kinetic and thermodynamic). For example, a common strategy, in many laboratories, to increase the conductivity of an electrolyte salt has been to convert as many anion C–H bonds into C–F bonds. This is not always the best approach for two different reasons: perfluorination is not necessarily better than polyfluorination. In addition, we used 1,2-dimethoxyethane (DME) for many of our experiments because of its lower viscosity, even though this solvent is unsuitable for secondary lithium-ion batteries. 2. LITHIUM SALTS OF AL(ORF)4 SUPERWEAK ANIONS Data for the six lithium salts studied are listed in Table 1 [6]. In addition, results for LiPF6 have been added for comparison. The structures of the alkoxide substituents and the abbreviations used are also shown in Table 1. 2.1. Thermal stability

Differential scanning calorimetry experiments revealed that two of the aluminate salts, LiAl(HFIP)4 and LiAl(HFPP)4, undergo thermal decomposition at 100°C, significantly higher than the 40°C thermal decomposition point of solid LiPF6 [1a]. In addition, when ethylenecarbonate/dimethylcarbonate (EC/DMC) solutions of these two lithium aluminates were heated to 100°C for 1 day, the room temperature conductivities, 19F-NMR spectra, and appearance (colorless solutions) were unchanged. Solutions of LiPF6 in EC/DMC are reported to decompose at 85°C [1b]. An EC/DMC solution of LiPF6 decomposed when heated to only 70°C for 1 day (the evidence for decomposition was a color change and the formation of a precipitate) [6]. 2.2. Conductivity

The DME σmax values for the six lithium aluminate electrolytes and for LiPF6 are listed in Table 1, along with 50:50 mol% EC/DMC σmax values for LiAl(HFIP)4, LiAl(HFPP)4, and LiPF6, and propylenecarbonate (PC) σmax values for LiAl(HFIP)4 and LiPF6. Also listed are 0.1 and 0.2 M DME σ values for the seven electrolytes. Fig. 1 displays a plot of molar conductivity (Λ) vs. the square root of the concentration for DME solutions of LiAl(HFIP)4. Similar plots were obtained for LiAl(HFTB)4, LiAl(HFPP)4, and LiAl(PFTB)4. The shape of the plot, with a local minimum in Λ at ca. 0.01 M and a local maximum in Λ at ca. 0.1 M, is common for electrolytes in low-dielectric solvents (a similar curve for LiBF4 in DME has been reported [14]) and has been interpreted in terms of significant triple-ion

Compound Abbreviation Alkoxide structure LiPF6

LiAl(OCH(CF3)2)4 LiAl(HFIP)4

Thermal decomposition temperature

Eox (solvent, conc.) (V vs. Li/0)

σmax (solvent, concentration) (mS cm1)

Emax for negligible aluminum corrosion (EC/DMC, V vs. Li/0)b

40°C (solid)c 70°C (EC/DMC)

 5.0 (EC/DMC, 1.0 M)

5.3 (DME, 0.3 M)d 3.0 (DME, 0.2 M)d 1.3 (DME, 0.1 M)d 11.5 (EC/DMC, 1 M) 5.8 (PC, 1 M)

5.0 V

100°C (solid) 100°C (EC/DMC)

 5.2 (DME, 0.1 M)  5.0 (EC/DMC, 0.5 M)

11.2 (DME, 0.5 M) 6.2 (DME, 0.2 M)d 3.2 (DME, 0.1 M)d 6.3 (EC/DMC, 0.6 M) 3.4 (PC, 0.6 M)

5.0 V

100°C (solid) 100°C (EC/DMC)

 5.2 (DME, 0.1 M)  5.0 (EC/DMC, 0.3 M)

4.2 (DME, 0.3 M) 3.5 (DME, 0.2 M)d 2.1 (DME, 0.1 M)d 3.2 (EC/DMC, 0.3 M)

5.0 V

–

O C CF 3 H CF3

LiAl(OCPh(CF3)2)4 LiAl(HFPP)4 O– C CF 3 CF3

LiAl(OC(CF3)3)4 LiAl(PFTB)4 O– C CF F3 C CF3 3

6.4 (DME, 0.2 M)d 3.6(DME, 0.1 M)d

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 197

Table 1 Thermal and electrochemical stabilities and maximum conductivities of LiPF6 and LiAlFn(ORF)4–n electrolytesa

Compound Abbreviation Alkoxide structure

198

Table 1 (Continued ) Thermal decomposition temperature

LiAl(OC(CH3)(CF3)2)4 LiAl(HFTB)4

σmax (solvent, concentration) (mS cm1)

Emax for negligible aluminum corrosion (EC/DMC, V vs. Li/0)b

9.6 (DME, 0.5 M) 6.2 (DME, 0.2 M)d 3.4 (DME, 0.1 M)d

–

O H3 C

Eox (solvent, conc.) (V vs. Li/0)

C CF 3 CF3

 80°C (solid)

2.6 (DME, 0.5 M) 1.3 (DME, 0.2 M)d 0.5 (DME, 0.1 M)d

 60°C (solid)

5.4 (DME, 0.8 M) 1.6 (DME, 0.2 M)d 0.6 (DME, 0.1 M)d

O– C CF 3 CF3

LiAl(OCH2CF3)4 LiAl(TFE)4 O– C CF 3 H H

Abbreviations: HFIP  OCH(CF3)2, HFTB  OC(CH3)(CF3)2, HFPP  OCPh(CF3)2, TFE  OCH2CF3, DME, 1,2-dimethoxyethane, EC/DMC, 50:50 mol% EC/DMC, PC, 1,2-propylene carbonate, Eox, potential at which the aluminate anion is oxidized in the indicated solvent; σmax, maximum conductivity in the solvent and at the concentration (conc.) indicated in parentheses. a All data from this work unless otherwise indicated. b Duration of the experiment was 1 h. c From Tsujioka and co-workers [6]. d The conductivity values in italics do not represent maximum conductivities.

Benjamin G. Nolan et al.

LiAlF(OCPh(CF3)2)3 LiAlF(HFPP)3

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 199

molar cond., Λ, S cm2 mol–1

60

40

20

0.0

0.2

0.4

0.6 ]1/2,

[LiAl(HFIP)4

0.8

M1/2

Fig. 1. Molar conductivity vs. the square root of the concentration for DME solutions of LiAl(HFIP)4. The lines drawn through the data points are guide to the eye. The molecular species shown is the [Al(HFIP)4–Li–Al(HFIP)4] triple ion in the structure of [1-Et-3-Me-1,3C3H3N2]-[Li(Al(HFIP)4)2] (redrawn from Ivanova and co-workers [10]). The unlabeled shaded gray spheres are oxygen atoms. The unlabeled unshaded and shaded white spheres represent carbon and fluorine atoms, respectively. Reproduced with permission from J. Electrochem. Soc., 151, A1418 (2004). Copyright 2004, The Electrochemical Society.

formation at concentrations above the local minimum [15]. The X-ray structure of the Li(Al(HFIP)4)2 inner-sphere triple ion is also shown in Fig. 1 [10]. The Li ion is coordinated to two polyfluoroalkoxide oxygen atoms from each of the two Al(HFIP)4 anions (and, interestingly, is also weakly coordinated to four C–F bonds from four different CF3 groups). This is a rare example of the isolation and structural characterization of a bonafide triple anion containing a central Li moiety. The conductivity of a lithium salt of a molecular anion in a given solvent is dependent, to varying degrees at different concentrations, on (i) the coordinating and ion-pairing ability of the anion and (ii) the mobility of the anion, which in turn is closely correlated with the size of the anion. Fig. 2 shows σ vs. DME concentration curves for five of the six lithium aluminate electrolytes in this study. The 0.1 and 0.2 M DME σ values for LiAl(PFTB)4, which are not shown, are virtually the same as the corresponding σ values for LiAl(HFIP)4 and LiAl(HFTB)4.

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12

conductivity, σ, mS cm–1

LiAl(HFIP)4

LiAl(HFTB)4

8

LiAl(TFE)4 LiAl(HFPP)4 4

LiAlF(HFPP)3 0 0

0.2

0.4 concentration, M

0.6

Fig. 2. Conductivities of DME solutions of LiAl(HFIP)4, LiAl(HFTB)4, LiAl(HFPP)4, LiAl(TFE)4, and LiAlF(HFPP)3. Reproduced with permission from J. Electrochem. Soc., 151, A1418 (2004). Copyright 2004, The Electrochemical Society.

Solutions of LiAl(PFTB)4 in DME, which are more concentrated than 0.2 M, could not be prepared because of the limited solubility of this perfluorinated electrolyte. Therefore, the lower solubility of the perfluorinated salt LiAl(PFTB)4 relative to the polyfluorinated salt LiAl(HFTB)4 has resulted in a significantly lower DME σmax value for LiAl(PFTB)4 (6.4 mS cm1 at 0.2 M) than for LiAl(HFTB)4 (9.6 mS cm1 at 0.5 M). It is significant that the σ values for LiAl(HFPP)4 are lower at all concentrations than the corresponding σ values for LiAl(HFIP)4, LiAl(HFTB)4, or LiAl(PFTB)4. Since all four anions were very weakly coordinating (see below), we did not think that differences in anion basicity or ion-pairing ability were responsible for the differences in σ values. We proposed that the difference is due to the significantly larger size and, hence, significantly increased solution viscosity and concomitant decreased mobility of the Al(HFPP)4 anion relative to the other three Al(OCR(CF3)2)4 anions (R  H, CH3, CF3). This is consistent with the fact that σmax occurs at only 0.3 M for LiAl(HFPP)4, instead of at 0.5 M for LiAl(HFIP)4 and LiAl(HFTB)4. It is also consistent with the observation that

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 201

σmax for LiAl(HFIP)4 is 17% higher than σmax for LiAl(HFTB)4, which contains the (slightly) larger anion. Not surprisingly, anion mobility does not affect σ as much at low concentrations as it does at high concentrations: the 0.05 M DME σ values of LiAl(HFIP)4, LiAl(HFTB)4, and LiAl(HFPP)4 are very similar because their ionpairing abilities, as well as their coordinating abilities towards Li (i.e. Lewis basicities), are very similar [10]. We independently determined the coordinating abilities of several tetrakis(polyfluoroalkoxy)aluminate anions by determining the position of the following equilibrium by NMR spectroscopy in dichloromethane solution for ORF  PFTB, HFPP, HFTB, HFBuPP (OC(pC6H4(t-Bu))(CF3)2), HFCP (OC(C6H11)(CF3)2), and DPTE (OC(C6H5)2(CF3)): LiAl(HFPP)4  N(n-Bu)4Al(ORF)4 y LiAl(ORF)4  N(n-Bu)4Al(HFPP)4 The results are shown in Fig. 3 [10]. Note that the 0.1 and 0.2 M DME σ values for LiAl(PFTB)4 and LiAl(HFTB)4 are virtually the same even though the equilibrium constants for the above reaction in dichloromethane are significantly different for

104

LiAl(HFPP)4 + Al(ORF)4−

Al(ORF)4–Lewis basicity (Keq)

Keq 102

CH2Cl2 DPTE

Al(HFPP)4− + LiAl(ORF)4 HFCP HFBuPP

100

HFTB

HFPP

10–2

10–4

PFTB

6

7

8 9 10 pKa of fluoroalcohol

11

12

Fig. 3. Plot of log Keq for the exchange reaction LiAl(HFPP)4  Al(ORF)4 y LiAl(ORF)4  Al(HFPP)4 vs. aqueous pKa value for the corresponding parent fluoroalcohol. The log Keq values for Al(PFTB)4 and Al(DPTE)4 represent upper and lower limits, respectively. The straight line is a linear least-squares fit to the four central data points. Note that log Keq is defined as 0 for the Al(HFPP)4 anion.

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ORF  PFTB (Keq  103) and ORF  HFTB (Keq  101). These results underscore the important, although perhaps less appreciated, distinction between the ionpairing ability of an anion toward Li in a strongly coordinating solvent (DME) and the coordinating ability (Lewis basicity) of the anion toward Li in a weakly coordinating solvent such as dichloromethane. The LiAl(TFE)4 electrolyte, which contains the smallest of the five tetrakis(polyfluoroalkoxy)aluminates, also has the lowest DME σ values at 0.1–0.4 M. The anion in this electrolyte is electronically and sterically different than the four Al(OCR(CF3)2)4 anions (R  H, CH3, CF3, Ph). It has only one CF3 group per alkoxide substitutent and is the smallest of the five Al(ORF)4 anions. For both these reasons, we conclude that the Al(TFE)4 anion is ion-paired with Li to a much greater extent than are the four Al(OCR(CF3)2)4 anions listed in Table 1, leading to relatively low σ values for LiAl(TFE)4. On the other hand, the small size of the Al(TFE)4 anion is responsible for the fact that σ continues to increase at concentrations between 0.5 and 0.8 M, presumably because the viscosity of a 0.5 M DME solution of LiAl(TFE)4 is lower than the viscosity of a 0.5 M DME solution of either LiAl(HFIP)4 or LiAl(HFTB)4. The LiAlF(HFPP)3 electrolyte is unique in that it contains a very polar Al–F bond with a strongly coordinating fluorine atom (in contrast to the weakly coordinating C–F bonds). The X-ray structures of LiAl(HFPP)4 [16] and LiAlF(HFPP)3 [17] are compared in Fig. 4. The latter structure contains a diamond-shaped Li2F2 core involving the fluorine atoms that are bonded to the Al atoms. One of the two Li–F(Al) bonds is 1.821(8) Å, shorter than any of the Li–O(Al) bonds, that range from 1.978(8) to 2.017(8) Å in either structure. Therefore, the presence of the Al–F bond in the AlF(HFPP)3 anion renders this anion much more strongly coordinating (and possibly more strongly ion-pairing) than the Al(HFPP)4 anion. This explains why the conductivity of LiAlF(HFPP)3 in DME is so much lower than that of LiAl(HFPP)4, despite the fact that the AlF(HFPP)3 anion is smaller than Al(HFPP)4. Consistent with this, σmax for LiAlF(HFPP)3, while lower than σmax for LiAl(HFPP)4, occurs at a higher concentration than for LiAl(HFPP)4. The 0.1, 0.2, and 0.3 M DME σ values for LiPF6, which also are not shown, are lower than the corresponding values for four of the six lithium aluminates (see Table 1). We suggested that the PF6 anion, like the AlF(HFPP)3 and Al(TFE)4 anions, is more strongly associated with Li in DME than are the four Al(OCR(CF3)2)4 anions [6]. In contrast, the EC/DMC σ values for LiPF6 are higher than for LiAl(HFIP)4 and LiAl(HFPP)4, as shown in Fig. 5. This is clearly due to both the higher viscosity and the higher dielectric constant of the EC/DMC mixture relative to DME. Nevertheless, the EC/DMC σmax for LiAl(HFIP)4, 6.3 mS cm1, is high enough for this thermally stable electrolyte to be considered as a replacement for LiPF6 in secondary lithium-ion batteries.

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 203

Fig. 4. Drawings of the structures of the monomeric structure LiAl(HFPP)4 (right, [18]) and the centrosymmetric dimeric structure of LiAlF(HFPP)3 (left, [17]). The unlabeled shaded circles are fluorine atoms and the unlabeled open circles are carbon atoms. Hydrogen atoms have been omitted for clarity. The unlabeled solid gray spheres are oxygen atoms. The unlabeled unshaded and shaded white spheres represent carbon and fluorine atoms, respectively. Reproduced with permission from J. Electrochem. Soc., 151, A1418 (2004). Copyright 2004, The Electrochemical Society.

We investigated whether 50:50 mol% EC/DMC was the optimum blend of these solvents for the LiAl(HFIP)4 electrolyte. The results show that a blend between 20:80 and 30:70 mol% EC/DMC gave a marginally higher σ value than the standard 50:50 mol% mixture [6]. 2.3. Electrochemical stability

The electrochemical stabilities of the Al(HFIP)4 and Al(HFPP)4 anions were investigated using cyclic voltammetry [6]. Negligible faradaic current was observed between 0 and 5.2 V vs. Li/0 (conditions: 0.1 M Li salt in DME, glassy carbon working electrode, 5 mVs1). At potentials higher than 5.2 V, irreversible DME oxidation, and possible anion oxidation, occurred. Below 0 V, plating of lithium was observed. The same results were obtained for a 0.1 M DME solution of LiPF6. In addition, a 0.5 M EC/DMC solution of LiAl(HFIP)4 did not undergo oxidation at potentials  5 V vs. Li/0. At potentials 5 V, irreversible oxidation of EC, DMC, and/or the Al(HFIP)4 anion occured.

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conductivity, σ, mS cm–1

LiPF6

8

LiAl(HFIP)4 4

LiAl(HFPP)4

0 0

0.4 0.8 concentration, M

1.2

Fig. 5. Conductivities of 50:50 mol% EC/DMC solutions of LiAl(HFIP)4, LiAl(HFPP)4, and LiPF6. Reproduced with permission from J. Electrochem. Soc., 151, A1418 (2004). Copyright 2004, The Electrochemical Society.

2.4. Lack of Reactivity with MCMB Carbon, LiCoO2, and Aluminum

Cyclic voltammograms (CVS) of EC/DMC solutions of LiAl(HFIP)4 (0.5 M) and LiPF6 (1 M) using an MCMB carbon electrode were similar between 0 and 2 V vs. Li/0 (5 mVs1 scan rate) [6]. In both cases, efficient reductive intercalation of Li was observed. There were no other reactions of the electrolyte with the carbon electrode. The same was true for both electrolyte solutions when a LiCoO2 electrode was used instead of the carbon electrode (the potential range investigated was 2.4–4.8 V vs. Li/0). The lack of any tendency of LiAl(HFIP)4 and LiAl(HFPP)4 to promote the corrosion of aluminum was investigated by cyclic voltammetry, chronoamperometry, and scanning electron microscopy (SEM) [6]. CVs for these two electrolytes and for LiPF6 in PC using an aluminum working electrode are shown in Fig. 6. In each experiment, potentials between 2 and 5 V vs. Li/0 were scanned five times. Passivation of the electrode surface was evident for all three electrolytes during the first scan. By the fifth scan, the current density at 5 V was

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 205

current density J, μA cm–2

10 5 0 10 5 0 10 5

LiAl(HFPP)4 scan 1

LiAl(HFPP)4 scans 2–5

LiAl(HFIP)4 scan 1

LiAl(HFIP)4 scans 2–5

LiPF6 scan 1

LiPF6 scans 2–5

0 0

2

4

0

potential, E, vs.

2

4

Li+/0

Fig. 6. Cyclic voltammograms of 0.3 M propylene carbonate solutions of LiAl(HFIP)4, LiAl(HFPP)4, and LiPF6 (conditions: aluminum working electrode; lithium foil counter and reference electrodes; 10 mV s1). Reproduced with permission from J. Electrochem. Soc., 151, A1418 (2004). Copyright 2003, The Electrochemical Society.

 2.2 μA cm2. After the fifth scan, each aluminum electrode was examined by SEM, and no evidence of corrosion (i.e. pitting) was observed. In addition, new aluminum working electrodes were held at 5 V vs. Li/0 in each electrolyte solution for 1 h. Again, no pitting of any of the three electrodes was observed by SEM. 3. LITHIUM SALTS OF B(OC(2-O-C6H4nFn)(CF3)2)2 SUPERWEAK ANIONS Data for the nine LiB(OC(2-O-C6H4nFn)(CF3)2)2 salts studied are listed in Table 2 [5]. Also shown are the structures of the diolate substituents, the abbreviations used, and, for comparison, data for a 10th salt with alkyl-group substituents on the phenyl rings instead of fluorine-atom substituents. Despite the superweak nature of the Al(HFPP)4 anion, the DME σmax value for LiAl(HFPP)4 was only 4.2 mS cm1 for a 0.3 M solution. Since the DME conductivity of this compound was significantly higher than that of LiPF6 at 0.01 and 0.1 M, the low σmax value, as discussed above, can be attributed to the large size of the Al(HFPP)4 anion causing a sufficiently high viscosity to offset the larger number of charge carriers at concentrations  0.3 M. In addition to the problem of low σmax values, most of the Al(ORF)4 anions are rapidly hydrolyzed by small amounts of water. To offset the problems of low σmax and hydrolytic instability, we designed the borate anions discussed in this section of the chapter [5]. Our rationale was that bis(diolato)borates would be more stable to hydrolysis than the homologous

206

Table 2 Thermal and electrochemical stabilities and maximum conductivities of LiB(O២O)2 saltsa (O២O)2 Formula Abbreviation Structure

Minimum thermal stability of LiB(O២O)2b

Eox (DME), Eox (EC/DMC) (V vs. Li/0)

σmax (DME), σ (EC/DMC)c (mS cm–1)

Passivation of Pt electrode at 4–6 (V vs. Li /0)d

OC(2-O-C6H4)(CF3)22 F02–

200°C (18 h)

4.32 4.67

5.88 4.15c (3.97)



140°C (18 h)

4.48 4.81

5.39



130°C (2 days)

4.54 4.73

6.57



138°C (4 days)

4.40 4.78

6.89



O

–

O– C CF 3 CF3

–

–

O C CF 3 CF3

O F

OC(2-O-4-F-C6H3)(CF3)22 4-F12– O–

O– C CF 3 CF3

F

OC(2-O-5-F-C6H3)(CF3)22 5-F12 O

F

–

–

O C CF 3 CF3

Benjamin G. Nolan et al.

OC(2-O-3-F-C6H3)(CF3)22 3-F12–

–

O F

152°C (18 h)

4.57 4.92

6.39



145°C (2 days)

4.64 4.93

7.87



135°C (48 h)

4.70 4.92

7.55



150°C (2 days)

4.84 5.16

7.79



–

O

C CF 3 CF3 F

OC(2-O-4,5-F2-C6H3)(CF3)22 4,5-F22 O–

O– C CF 3 CF3

F F

OC(2-O-4,6-F2-C6H3)(CF3)22 4,6-F22 O–

F

O– C CF 3 CF3 F

OC(2-O-3,4,5-F3-C6H)(CF3)22 3,4,5-F32 O F

–

O

–

C CF 3 CF3

F F

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 207

OC(2-O-3,5-F2-C6H2)(CF3)22 3,5-F22

208

Table 2 (Continued) (O២O)2 Formula Abbreviation Structure

Minimum thermal stability of LiB(O២O)2b

OC(2-O-4,5,6-F3-C6H)(CF3)22 4,5,6-F32 O–

152°C (2 days)

Eox (DME), Eox (EC/DMC) (V vs. Li/0)

σmax (DME), σ (EC/DMC)c (mS cm–1)

Passivation of Pt electrode at 4–6 (V vs. Li /0)d

4.74 5.05

8.39



4.15 Not determined

4.31



–

O C CF 3 CF3

F

F

OC(2-O-3-tBu-5-Me-C6H2)(CF3)22 105°C (6 days) 3,5-R2-F02 O– t-Bu

–

O C CF 3 CF3

Me

Note: For Abbreviations see Table 1. The symbol O២O2 represents the diolate(2–) anions shown in the table. b These data are from Nolan and co-workers [18]. Decomposition temperatures were not determined. The temperatures listed were used to dry the lithium salts; no decomposition was observed for the time period shown in parentheses. The decomposition temperatures are probably much higher than the values listed. c The maximum conductivity of LiB(F0)2 in EC/DMC of 4.15 mS cm1 was for a 60:40 mol% mixture. The 50:50 mol% EC/DMC value is shown in parentheses. d The  sign indicates that passivation was observed for these three lithium salts; the  sign indicates that passivation was not observed. The same results were observed for both 0.1 M DME and 0.1 M 50:50 mol% EC/DMC solutions. a

Benjamin G. Nolan et al.

F

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 209

aluminates, and that each borate would bear two OC(2-O-C6H4nFn)(CF3)22 dianions instead of four OC(C6H5)(CF3)2 monoanions. The latter feature would lead to an anion of about half the mass and, hence, about half the volume of Al(HFPP)4. In fact, the crystallographically determined formula unit volumes are different by nearly a factor of 2,977 Å3 for LiAl(HFPP)4 [16] and ca. 490 Å3 for LiB(F0)2 [18]. The strategy was successful in that B(F0)2 is stable to hydrolysis (in fact, it is synthesized using water as the solvent [18]) and its lithium salt has a σmax value 40% higher than that of LiAl(HFPP)4, as shown in Fig. 7. The two σ vs. concentration plots in Fig. 7 also show that the DME conductivity of LiAl(HFPP)4 is higher than that of LiB(F0)2 at concentrations of 0.2 M or less, an observation that strongly suggests that the Al(HFPP)4 anion is more weakly coordinating and more weakly ion pairing than the B(F0)2 anion (the important distinction between the coordinating ability of an anion and its ion-pairing ability was discussed above). Therefore, the lower conductivity of LiAl(HFPP)4 at higher concentrations must be due to ion mobility and viscosity consequences of the larger size of the Al(HFPP)4 anion relative to the B(F0)2 anion. Starting with a variety of mono-, di-, and trifluorinated phenols, it was possible to prepare eight fluorinated derivatives of HO(2-OH–C6H4)(CF3)2 [19], from which the eight fluorinated derivatives LiB(Fn)2 were prepared [18]. The 6 LiAl(HFPP)4

conductivity, mS cm–1

LiB(F0)2

4

1.2 0.8

2

0.4 0 0

0.05

0 0

0.2 0.4 concentration of lithium salt, M

0.6

Fig. 7. Conductivities of DME solutions of LiB(F0)2 and LiAl(HFPP)4. The conductivity of the lithium borate electrolyte is the lower of the two at low concentrations and the higher of the two at high concentrations. Reproduced with permission from J. Electrochem. Soc., 150, A1726 (2003). Copyright 2003, The Electrochemical Society.

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availability of these nine lithium salts allowed the change in conductivity as a function of the number and positions of phenyl-ring electron-withdrawing groups to be measured with minimal change in the size of the borate anion (the van der Waals radii of hydrogen and fluorine atoms are 1.20 and 1.47 Å, respectively [20]). Thus, we were in a position to compare and contrast a series of nine electronically different but sterically similar lithium electrolytes. 3.1. Thermal Stability

Table 2 lists the formulas, abbreviations, and structures of the Fn2 and 3,5-R2 (O២O)2 ligands, the highest temperatures at which the 10 LiB(O២O)2 salts were dried, and the drying times [5]. These temperatures can be considered as minimum thermal stabilities. They are all 135°C or higher except for the 105°C value for LiB(3,5-R2)2. The criteria used to judge thermal stability were (i) no change in 19F-NMR spectra, (ii) no discoloration, and (iii) no further weight change once all of the moisture was removed by drying at the indicated temperatures. The temperatures at which the compounds decompose, which were not measured, are probably much higher. 3.2. Conductivity

The DME σmax values for the 10 new lithium borates are also listed in Table 2, along with the 0.5 M 50:50 mol% EC/DMC σ value for LiB(F0)2 [5]. The EC/DMC value, 3.97 mS cm1, is only 68% of the DME σmax value. We attributed the decrease to the greater viscosity of EC/DMC relative to DME. The maximum conductivity of LiB(F0)2 at 0.5 M in EC/DMC did not occur for a 50:50 mol% mixture. The 0.5 M σ values ranged from 0.95 mS cm1 for 100% DMC to a maximum of 4.15 mS cm1 for 40:60 mol% EC/DMC to 3.97 mS cm1 at 50:50 mol%. It is interesting to compare the ca. 4 mS cm1 σmax values for LiB(F0)2 with those for LiPF6, which are typically  10 mS cm1 [21]. At 0.01 M in DME, the σ values for LiB(F0)2 and LiPF6 are 0.137 and 0.073 mS cm1, respectively. These results demonstrate that the B(F0)2 anion is more weakly coordinating and more weakly ion pairing than PF6, but that the borate anion is sufficiently large and the solution viscosity restricts σ at high concentrations. Fig. 8 shows a plot of equivalent conductivity (Λ) vs. the square root of the concentration for 0.001–0.6 M DME solutions of LiB(F0)2. A similar plot was obtained for LiB(3,4,5-F3)2. As discussed above, the shape of this plot, with a local minimum in Λ at 0.01 M and a local maximum in Λ at 0.1 M, suggests a significant triple-ion formation at concentrations above the local minimum. The X-ray structure of [LiB(F0)2]2.DME [18], also shown in Fig. 8, shows the possible nature of an anionic triple ion. The Li1 ion is coordinated to a DME molecule and to the two alkoxy oxygen atoms of one of the borate anions. If we imagine the dissociation of Li1 from its borate, the [borate–Li2–borate]  triple

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 211

50

2

equiv. cond., S cm mol

–1

40

30

20

10

0 0

0.2

0.4

0.6

0.8

[LiB(F0)2]1/2, M1/2

Fig. 8. Equivalent conductivity vs. the square root of the concentration for 0.0001 to 0.6 M DME solutions of LiB(F0)2. The lines drawn through the data points are guide to the eye. The structure of [LiB(F0)2]2 DME (redrawn from Nolan and co-workers [18] is shown (the unlabeled atoms are carbon atoms). For the borate anion on the right, one of the fluorine atoms is hidden behind the boron atom. Reproduced with permission from J. Electrochem. Soc., 150, A1726 (2003). Copyright 2003, The Electrochemical Society.

.

ion would remain. Note that Li2 is coordinated to the two phenoxy oxygen atoms from each borate. The phenoxy oxygen atoms are almost certainly more strongly basic than the alkoxy oxygen atoms. One piece of evidence that supported this conclusion was found in the X-ray structure of [Li(acetone)B(4,5,6-F3)2]2 [18], in which the Li–O(phenoxide) bond is shorter (1.964(7) Å), and presumably stronger, than the Li–O(alkoxide) bond (2.038(7) Å; Δ(Li–O)  0.04 Å). Fig. 9 shows σ vs. concentration plots for DME solutions of the nine LiB(Fn)2 electrolytes [5]. All nine plots exhibit σmax at 0.5 M, an observation that supports the hypothesis that the nine B(Fn)2 anions have nearly the same size. The σmax values are listed in Table 2. All of the electrolytes have DME σmax values greater than 5 mS cm1, and four of them have σmax values greater than 7 mS cm1.

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Benjamin G. Nolan et al.

conductivity, mS cm–1

8

6

4

4,5,6-F3 3,4,5-F3 4,5-F2 4-F1 5-F1 F0 3-F1

2

4,6-F2 3,5-F2

O– Fn2– =

O– C CF 3 CF3

3 4

6 5

0 0

0.3

0.6

0 0.3 0.6 0 0.3 0.6 concentration of LiB (Fn)2, M

Fig. 9. Conductivities of 0.1 to 0.6 M DME solutions of nine LiB(Fn)2 salts. The structure of the generic Fn2 chelating diolate ligand is shown at the bottom right. The ligand abbreviations for each compound are shown in the rectangular boxes. The dashed lines drawn through the 0.2 M data points for LiB(3-F1)2, LiB(3,5-F2)2, and LiB(3,4,5-F3)2 demonstrate the observation that, in general, increasing the number of fluorine atom substituents on the diolate phenyl ring increases the conductivity. Note, however, that the conductivity of LiB(3-F1)2 is lower, not higher, than the conductivity of LiB(F0)2. Reproduced with permission from J. Electrochem. Soc., 150, A1726 (2003). Copyright 2003, The Electrochemical Society.

Barthel and co-workers [22] found that progressive fluorine-for-hydrogen substitution in LiB(1,2-C6H4O2)2 resulted in an increase in DME conductivity. The 25°C conductivities of 1.08 0.08 M solutions of LiB((1,2-C6H4O2)2, LiB((3-F-1,2-C6H3O2)2, and LiB((1,2-C6F4O2)2 are 2.2, 4.1, and 11.1 mS cm1, respectively. Similarly, the σmax values of LiB(F0)2, LiB(5-F1)2, LiB(4,5-F2)2, and LiB(4,5,6-F3)2 are 5.9, 6.9, 7.6, and 8.4 mS cm1, respectively. The dashed lines in Fig. 9 connecting the 0.2 M σ values for LiB(3-F1)2, LiB(3,5-F2)2, and LiB(3,4,5-F3)2 show that this is generally true for all concentrations from 0.1 to 0.6 M, not just for 0.5 M solutions. There is an interesting subset of the data that shows the opposite behavior. For the three cases of fluorine substitution of the two hydrogen atoms ortho to the two phenoxy oxygen atoms in each borate, the conductivity always decreases. The σmax values decrease from 5.9 for LiB(F0)2 to 5.4 mS cm1 for LiB(3-F1)2, from 6.9 for LiB(5-F1)2 to 6.4 mS cm1 for LiB(3,5-F2)2, and from 7.9 for LiB(4,5-F2)2 to 7.8 mS cm1 for LiB(3,4,5-F3)2. This is consistent with the results for 0.01 M DME

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 213

solutions of these compounds [18]. Note that the gas-phase proton basicity of the 2-fluorophenoxide ion is not only lower than the basicity of the phenoxide ion, but is also lower than the basicities of the 3- and 4-fluorophenoxide ions [23]. Therefore, B(3-F1)2 is a more strongly ion-pairing anion than B(F0)2, even though B(3-F1)2 is probably the more weakly coordinating of the two. To our knowledge, this is the first example showing that substituting hydrogen atoms for fluorine atoms in a molecular electrolyte salt can, in some cases, result in a decrease in conductivity at a given concentration. Clearly, conductivity can be influenced by the positions of the fluorine atoms as well as by the number of fluorine atoms. 3.3. Electrochemical Stability

The electrochemical stabilities of the borate anions were investigated using CV. None of the electrolytes was reduced at 0 V vs. Li/0 (conditions: 0.1 M in 50:50 mol% EC/DMC, stainless-steel working electrode; 0.1 M in DME, platinum electrode). A typical set of three voltammograms is shown in Fig. 10. The currents observed between 0 and 0.5 V and between 0.5 V and 0.5 V correspond to plating and stripping of metallic lithium, respectively. Since the couloumbic efficiency for the plating/stripping process is clearly less than 100%, the current between 0 and 0.5 V could potentially include reduction of the borate anions. However, a significant amount of anion reduction is unlikely because (i) no discoloration of the solutions occurred even after multiple cycles and (ii) the shapes 2000 F 3 C CF3 O C B C O O F3 C CF3 O

current, i, μA

1000

0

-1000

-2000 -1

0

2

1 E, V vs. Li

3

+/0

Fig. 10. Voltammogram showing the reductive stability over three scans from 3 to 0.5 V vs. Li/0 for a 0.1 M 50:50 mol% EC/DMC solution of LiB(F0)2. The current responses correspond to the plating and stripping of lithium over the three scans. A stainless-steel working electrode and lithium counter and reference electrodes were used. Reproduced with permission from J. Electrochem. Soc., 150, A1726 (2003). Copyright 2003, The Electrochemical Society.

214

Benjamin G. Nolan et al.

of the plating/stripping portions of the voltammograms were virtually the same for all of the salts used. Note that there is also a loss in cycling efficiency as evidenced by a decreasing stripping current over the three cycles. We did not further investigate the nature of the loss in plating/stripping efficiency. The potentials at which the borates are oxidized (Eox) in DME and in 50:50 mol% EC/DMC (same conditions as above) are also listed in Table 2. The Eox values were determined by linear extrapolation of the increasing oxidation current back to the horizontal zero-current line. With the exception of LiB(3,5-R2)2, all of the electrolytes are oxidized above 4.2 V, and the two trifluoro derivatives are oxidized above 5 V. Without exception, increasing the number of fluorine atoms on the borate phenyl rings increased Eox. This reflects a lowering of the borate HOMO energy with increasing fluorine content. This effect was reported by Barthel and co-workers [22c] for the B(C6H4nFnO2)2 anions discussed above (n  0, 1, 4). In contrast to conductivity, the potential at which a B(Fn)2 anion was oxidized was not significantly affected by the positions of the fluorine atoms on the borate phenyl rings. On average, Eox increased by 0.16 V per hydrogen-atom/fluorine-atom phenyl-ring substitution. 3.4. Residual Current, Electrode Passivation, and Aluminum Corrosion

Fig. 11 shows a typical set of three voltammograms from 0–3 V vs. Li/0 [5]. Similar voltammograms with similarly small maximum residual currents were

residual current, i, μA

30 F 3 C CF3 O C B C O O F3 C CF3 O

20 10 0 −10 −20 0

2

1 E, V

3

vs.Li+/0

Fig. 11. Voltammogram showing residual current over three scans from 0 to 3 V vs. Li/0 for a 0.1 M 50:50 mol% EC/DMC solution of LiB(F0)2. The structure of the borate anion is shown on the upper right. Virtually identical voltammograms were recorded for 0.1 M solutions of the nine other lithium borates as well as for a 0.1 M solution of LiPF6. A stainless-steel working electrode (area  0.48 cm2) and lithium counter and reference electrodes were used. Reproduced with permission from J. Electrochem. Soc., 150, A1726 (2003). Copyright 2003, The Electrochemical Society.

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 215

observed for 0.1 M 50:50 mol% EC/DMC solutions of all 10 lithium borates as well as for LiPF6. Therefore, the small currents are probably due to the reduction/oxidation of solvent impurities and not of the borate anions. It is possible that part of the residual current is due to underpotential plating and stripping of metallic lithium, as proposed by Malik and co-workers [24]. Not surprisingly, the residual currents increased slightly as the conductivity of the electrolyte increased. An interesting difference among the LiB(Fn)2 electrolytes is whether or not oxidation of the borate anion resulted in electrode passivation. Of the nine LiB(Fn)2 salts studied, only LiB(F0)2, LiB(4-F1)2, and LiB(4,6-F2)2 exhibited passivation in either DME (0.02 cm2 platinum working electrode) or 50:50 mol% EC/DMC (0.48 cm2 stainless-steel working electrode) [5]. Evidently, the borates with fluorine atoms ortho or para to the phenoxy oxygen atoms do not decompose in a way that leaves a coherent solid film on the electrode surfaces. As with conductivity, the passivating potential of at least some fluorinated organic electrolytes depends on the positions of the fluorine-atom substitutions. The tendency of the 10 lithium borate electrolytes, we studied to promote the corrosion of aluminum, was investigated by chronoamperometry at 4.2 V vs. Li/0 for 1 h [5]. The final current density for LiB(5-F1)2, and for the other nine lithium borates we studied, was ca. 0.1 μA cm2 (a 2.0-cm2 aluminum foil was used as working electrode). In contrast, the final current density for LiCF3SO3 was  104 times. The electrolyte LiCF3SO3 was previously shown to promote the corrosion of aluminum even at 3.0 V [2d]. 3.5. Comparison with Other Lithium Bis(diolato)Borates

Table 3 lists the thermal stabilities, Eox and Ered values, σmax values, and tendencies to promote aluminum corrosion for LiB(F0)2, LiB(4,5,6-F3)2, and 15 other lithium salts of B(O២O)2 anions [5, 22, 25–35]. In addition to LiB(F0)2 and LiB(4,5,6-F3)2, only two of the other 15 electrolytes have entries in all four data columns in Table 3. In most cases, the tendency to promote the corrosion of aluminum at high potential has not been reported. In four cases, no conductivity results have yet been published. In one case, only Eox is known. The most promising lithium borate electrolyte, and the one that should now be considered as the standard against which all others are measured, is LiB(C2O4)2, first reported in a German patent [25] and studied in detail by Angell and co-workers [26], who have given it the nickname LiBOB. Taking into account all of the accepted performance criteria except cost, its possible rivals for replacement of LiPF6 in large-scale primary and secondary lithium-ion batteries, for which thermal stability is a critical issue, are LiB(1,2C2(CF3)4O2)2 (depending on its tendency to promote aluminum corrosion as well as its tendency to hydrolyze to highly toxic perfluoropinacol), LiB(1,2-C6H4O2)2 (although its Eox value of 4.1 V may be problematic), and several of the LiB(Fn)2

(O២O)2 Formula Structure

Thermal stability of Eox (solvent) LiB(O២O)2 in the solid statebEred (solvent) (V vs. Li/0)

OC(2-O-C6H4)(CF3)22 O

–

O

σ (solvent, conc.) (mS cm1)

216

Table 3 Thermal and electrochemical stabilities and conductivities of various LiB(O២O)2 electrolytesa Aluminum corrosion (E vs. Li/0)

Ref.

200°C

4.67 (EC/DMC)

–0.5 (EC/DMC)

5.88 (DME, 0.5 M) 4.15 (EC:DMC, 0.5 M )

No (4.2 V)

[5]

152°C

5.05 (EC/DMC)

0.5 (EC/DMC)

8.39 (DME, 0.5 M)

No (4.2 V)

[5]

302°C

5.0 (PC)

0.0 (PC)

14.9 (DME, 1 M) 25.2 (AN, 1 M)

No (5.75 V)

[25–27]

245°C

Not reported

Not reported

Not reported

[25–27]

130°C

ca. 5 (DME or PC)

0.0 V (DME or PC)

11.1 (DME, 0.6 M) 2.1 (PC, 1 M)

Not reported

[28]

–

C CF 3 CF3

OC(2-O-4,5,6-F3-C6H)(CF3)22 –

O C CF 3 CF3

F

F F

C2O42 O– O–

O O

C3H2O42 O

–

O O

–

O

1,2-C2(CF3)4O22 O– CF 3 CF3

O

–

CF3 CF3

Benjamin G. Nolan et al.

O

–

1,2-C6H4O22

3.6 (PC)

0.5 (PC)

2.24 (DME, 1.3 m) 1.74 (PC, 0.5 m) 5.0 (EC/DME, 0.5 M)

not reported

[22,29]

150°C

3.7 (PC)

–0.5 (PC)

4.09 (DME, 1.2 m) 2.40 (PC, 0.4 m)

Not reported

[22b]

270°C

4.1 (PC)

0.5 (PC)

11.1 (DME, 1 M)

Not reported

[22c]

280°C

3.7 (PC)

0.5 (PC)

4.2 (EC/DME, 0.5 M)

Not reported

[29,30]

100°C

3.95 (EC/DMC/PC) Not reported

0.8 (EC/DMC/PC, 0.71 m)

No (4.5 V)

[31]

O–

3-F-1,2-C6H3O22 O

–

O–

F

1,2-C6F4O22 –

O

O–

F

F

F F

2,3-C10H6O22 O– O–

2,3-C5H3NO22 –

O

–

N

O

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 217

O

290°C

–

218

Table 3 (Continued) (O២O)2 Formula Structure 2,2-C12H8O22

Thermal stability of Eox (solvent) LiB(O២O)2 in the solid statebEred (solvent) (V vs. Li/0)

σ (solvent, conc.) (mS cm1)

Aluminum corrosion (E vs. Li/0)

Ref.

4.1 (PC)

0.0 (PC)

1.0 (EC/DME, 0.5 M)

Not reported

[32,33]

300°C

4.4 (PC) ca. 0.6 (PC)

3.0 (EC/DME, 0.5 M)

Not reported

[29,32]

300°C

4.2 (PC) Not reported

2.5 (EC/DME, 0.5 M)

Not reported

[29]

Not reported

4.1 Not reported

Not reported

Not reported

[34]

O–

C6H4(CO2)O2 –

O

–

O C

O

2-O-3-CH3-C6H3CO22 O

–

O

–

C

CH3

O

2-O-4-CF3-C6H3CO22 O–

O– C O

CF3

Benjamin G. Nolan et al.

270°C

–

O

2-O-3,5-Cl2-C6H2CO22 O

Cl

300°C

4.4 (PC) Not reported

3.9 (EC/DME, 0.3 M)

Not reported

[29]

300°C

4.4 (PC) Not reported

Not reported

Not reported

[29]

70°C

4.6 (EC/DMC)

Not reported

No (4.5 V)

[34,35]

–

C O

Cl

2-O-3,5,6-Cl3-C6HCO22 –

–

O

O C

Cl

O Cl Cl

2-O-5-F-C6H3SO32 –

O

–

O

S O O F

Abbreviations: DME, 1,2-dimethoxyethane; EC, ethylene carbonate; DMC, dimethylcarbonate; PC, propylene carbonate; DMSO, dimethylsulfoxide; AN, acetonitrile; Eox, potential at which the borate anion is oxidized; Ered, potential at which the borate anion is reduced; σ, conductivity in the solvent and at the concentration (conc.) indicated in parentheses. a Temperatures at which decomposition, or significant weight loss, is observed. For values with  signs, these are minimal thermal stabilities.

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 219

O

–

220

Benjamin G. Nolan et al.

electrolytes (although they are limited to 0.5 M solutions, at least in DME and EC/DMC, and their toxicity is unknown). ACKNOWLEDGMENTS This research was supported by U.S. National Science Foundation and by Central Glass Company. REFERENCES [1] [2]

[3]

[4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20]

(a) T.C. Ehlert and M.M. Hsia, J. Chem. Eng. Data, 17 (1972) 18; (b) S.E. Sloop, J.K. Pugh, S. Wang, J.B. Kerr, and K. Kinoshita, Electrochem. Solid–State Lett., 4 (2001) A 42. (a) D. Linden and T.B. Reddy (Eds.), Handbook of Batteries, 3rd edn., McGraw-Hill, New York 2002; (b) J. Barthel and H.J. Gores, Handbook of Battery Materials, J.O. Besenhard (Ed.), Wiley, Weinheim, 1999; (c) M. Morita, M. Ishikawa, and Y. Matsuda, Lithium Ion Batteries: Fundamentals and Performance, M. Wakihara and O. Yamamoto (Eds.), Wiley, Weinheim, 1998, p.156; (d) G.E. Blomgren, Lithium Batteries, J.P. Gabano (Ed.), Academic Press, New York, 1983, Chap. 2. (a) C.W. Walker, Jr., J.D. Cox, and M. Salomon, J. Electrochem. Soc., 143 (1996) 154; (b) M. Ue, J. Electrochem. Soc., 143 (1996) L270 and references therein; (c) L.A. Dominey, V.R. Koch, and T.J. Blakeley, Electrochim. Acta, 37 (1992) 1551; (d) A. Webber, J. Electrochem. Soc., 138 (1991) 2586. J. Barthel, A. Schmid, and H.J. Gores, J. Electrochem. Soc., 147 (2000) 21 and references therein. B.G. Nolan and S.H. Strauss, J. Electrochem. Soc., 150 (2003) A1726. S. Tsujioka, B.G. Nolan, H. Takase, B.P. Fauber, and S.H. Strauss, J. Electrochem. Soc., 151 (2004) A1418. I. Krossing and I. Raabe, Angew. Chem. Int. Ed., 43 (2004) 2066 and references therein. C.A. Reed, K.C. Kim, E.S. Stoyanov, D. Stasko, F.S. Tham, L.J. Mueller, and P.D.W. Boyd, J. Am. Chem. Soc., 125 (2003) 1796 and references therein. A.J. Lupinetti and S.H. Strauss, Chemtracts - Inorg. Chem., 11 (1998) 565 and references therein. S.M. Ivanova, B.G. Nolan, Y. Kobayashi, S.M. Miller, O.P. Anderson, and S.H. Strauss, Chem.-Eur. J., 7 (2001) 503 and references therein. S.H. Strauss, B.G. Nolan, T.J. Barbarich, and J.J. Rockwell, Weakly Coordinating Anions Containing Polyfluoroalkoxide Ligands, U.S. Patent 6,221,941 B1, April 24, 2001. I. Krossing and L. Van Wullen, Chem.-Eur. J., 8 (2002) 700 and references therein. H. Tokuda and M. Watanabe, Electrochim. Acta, 48 (2003) 2085. J. Barthel, R. Gerber, and H.J. Gores, Ber. Bunsenges. Phys. Chem., 88 (1988) 616. R.M. Fuoss and C.A. Kraus, J. Am. Chem. Soc., 55 (1933) 2387. T.J. Barbarich, S.T. Handy, S.M. Miller, O.P. Anderson, P.A. Grieco, and S.H. Strauss, Organometallics, 15 (1996) 3776. B.G. Nolan, T.J. Barbarich, S. Tsujioka, S.M. Ivanova, B.P. Fauber, S.M. Miller, O.P. Anderson, and S.H. Strauss, in preparation. B.G. Nolan, S.M. Miller, O.P. Anderson, and S.H. Strauss, in preparation. B.G. Nolan, S. Tsujioka, and S.H. Strauss, J. Fluorine Chem., 118 (2002) 103. (a) A. Bondi, J. Phys. Chem., 68 (1964) 441; (b) L. Pauling, The Nature of the Chemical Bond, Cornell University Press, Ithaca, NY, 1960.

Electrochemical properties of lithium electrolytes based on bis(polyfluorodiolato)borate and tetrakis(polyfluoroalkoxy)aluminate superweak anions 221 [21] C.A. Angell and W. Xu, PCT Intl. Appl. WO 0199209 A2, 2001. [22] (a) J. Barthel, M. Wühr, R. Buestrich, and H.J. Gores, J. Electrochem. Soc., 142 (1995) 2527; (b) J. Barthel, R. Buestrich, E. Carl, and H.J. Gores, J. Electrochem. Soc., 143 (1996) 3565; (c) J. Barthel, R. Buestrich, E. Carl, and H.J. Gores, J. Electrochem. Soc., 143 (1996) 3572. [23] J.J. Urban, R.L. von Tersch, and G.R. Famini, J. Org. Chem., 59 (1994) 5239. [24] Y. Malik, D. Aurbach, P. Dan, and A. Meitav, J. Electroanal. Chem., 282 (1990) 73. [25] U. Lishka, U. Wietelmann, M. Wegner, German Patent DE 19829030 C1, 1999. [26] (a) W. Xu and C.A. Angell, Electrochem. Solid–State Lett., 4 (2001) E1; (b) K. Xu, S. Zhang, T.R. Jow, W. Xu, and C.A. Angell, Electrochem. Solid–State Lett., 5 (2002) A 26. [27] (a) S. Tsujioka, H. Takase, and M. Takahashi, European Patent Appl. EP 2000-115578, 2001; (b) U. Heider, M. Schmidt, A. Kuehner, and A. Schmenger, U.S. Patent Appl. 2001-758546, 2001; (c) J.-C. Panitz, U. Wietelmann, and M. Scholl, German Patent DE 10111410 C1, 2002. [28] W. Xu and C.A. Angell, Electrochem. Solid–State Lett., 3 (2000) 366. [29] Y. Sasaki, M. Handa, K. Kurashima, T. Tonuma, and K. Usami, J. Electrochem. Soc., 148 (2001) A999. [30] J. Barthel and R. Bustrich, German Patent DE 19633027 A1, 1998. [31] J. Barthel, A. Schmid, and H.J. Gores, J. Electrochem. Soc., 147 (2000) 21. [32] J. Barthel, R. Buestrich, H.J. Gores, M. Schmidt, and M. Wühr, J. Electrochem. Soc., 144 (1997) 3866. [33] Y. Sasaki, S. Sekiya, M. Handa, and K. Usami, J. Power Sources, 79 (1999) 91. [34] J. Barthel, H.J. Gores, R. Neuder, and A. Schmid, Pure Appl. Chem., 71 (1999) 1705. [35] J. Barthel, M. Schmidt, and H. J. Gores, J. Electrochem. Soc., 145 (1998) L17.

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Chapter 10

Fluorinated electrolytes based on lithium salts of strong Brønsted acids Olt E. Geiculescu, Stephen E. Creager, and Darryl D. DesMarteau Department of Chemistry, Clemson University, Clemson, South Carolina 29634, U.S.A. 1. INTRODUCTION Ionic conductivity in solvent-free solid polymer electrolytes (SPEs) has been extensively studied because of the potential applications for such materials in electrochemical power sources and devices [1–3], particularly in high-energydensity rechargeable lithium batteries [4,5]. The SPEs have many advantageous properties for such applications including good dimensional and thermal stability, a wide electrochemical stability window, better shape flexibility and manufacturing integrity, and improved safety. Since the pioneering studies of materials based on alkali metal-salt complexes with poly(ethylene oxide) (PEO) reported by Wright [6] and Armand et al. [7] in the 1970s, the SPEs based on polyethers have undergone rapid growth both in academia and in industrial research and development, mostly relating to applications in secondary lithium batteries [8–13]. For a polymer to be used as a successful host for a salt it should provide coordinating sites capable of facilitating ion separation by solvating the cations of the salt, which will compensate for the lattice dissolution energy of the salt. Also, for achieving good salt solubility and to facilitate ion motion, appropriate distances between coordinating sites and low barriers to bond rotation are necessary[1]. Such potential polymeric hosts could be polyethers, polyesters, polyimines, and polythiols, all of which have available electron pairs that can solvate cations easily. Among these polymers, electrolytes based on polyethers were tested on a large scale, with PEO showing the best spacing for maximum solvation, while poly(methylene oxide) (PMO) and poly(propylene oxide) (PPO) were much weaker solvents [1]. PEO-based complexes of alkali metal salts (especially of lithium) have been extensively studied due to the extremely strong solvating properties of PEO for a wide variety of salts [14–18]. In particular, PEO is of interest

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because of the ability of the polar ether groups in the main chain to coordinate easily with salt cations, in a manner similar to crown ethers, forming a homogenous solution. This implies that the spacing and conformational flexibility provided by the PEO unit, (–CH2CH2O–), are optimal for coordination with the cation, the PEO unit acting as a Lewis base and the cation as a Lewis acid. These electrolytes commonly exhibit ionic conductivities that range from 108 to 104 S/cm at temperatures between 40 and 100°C, which precludes, for the time being, their practical applications in battery technology at ambient temperature [19]. On the one hand, the low conductivity is due to the high degree of crystallinity of PEO (60–80% depending on the molecular weight [20]), which is not favorable for batteries since ion conduction proceeds only through the elastomeric amorphous phase and, on the other, to the low solubility of the salt in the amorphous phase [21,22]. The stoichiometric complexes of PEO with salts are crystalline, welldefined materials that do not participate in ionic conductivity due to the total occupancy of the solvation sites [23,24]. Several researchers have clearly shown that ion transport preferentially occurs in the amorphous elastomeric phase of the PEO – salt compounds, such that the morphological structure of PEO – salt compounds plays an important role in determining ion transport [21,22]. Also, the PEO host has sufficient local dynamic mobility to allow for ion transport despite the relatively high macroscopic viscosity that is characteristic of such materials [25,26]. The polymer motions relevant to ionic conductivity are the side-chain segmental motions rather than the diffusion of the entire polymer backbone [25]. Unfortunately, PEO tends to crystallize or form crystalline complexes with an increase in salt concentration, resulting in a sharp decrease in ionic conductivity. Therefore, reducing the crystalline regularity of the PEO – salt complex and creating a system that remains amorphous over the entire temperature range of interest proved to be a challenging task. Generally, the requirements for solid polymer electrolytes used for rechargeable lithium batteries involve conductivities between 104 and 103 S/cm and a good dimensional and thermal stability, all in the temperature range of 40 to 70°C. Additionally, an electrochemical stability window spanning the potentials between 0.0 and  4.0 V vs. Li/Li, chemical compatibility with both the Li anode and the cathode, and an ability to afford Li cycling at an efficiency of  99% are also desired [11]. The research has primarily concentrated in two directions: increasing the segmental mobility of the PEO polymer through copolymerization [27], grafting [10], cross-linking [15,28,29], modification of polymer by pendant PEO segments [30,31] and plasticization of matrix polymers [32,33] or by using new lithium salts with a lower lattice energy that makes them easier to be solvated by the host polymer [10,34]. Of all the alkali metals used in SPEs and in battery formulation, lithium was by far the most popular choice due to its highest negative standard electrode potential (Li/Li couple is 3.045 V vs. standard hydrogen electrode, or SHE), highest specific capacity (4 A h/g), and lowest density of all

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225

metals (0.53 g/cm3). Generally, the SPEs are formed when the solvating energy of the polymeric host (exothermic process) is higher than the lattice energy of the salt (endothermic process), which favors the use of salts consisting of a polarizing cation and a large anion of delocalized charge. The salt affects the ionic conductivity through crystalline complex formation, intra- or intermolecular cross-linking of the polymer chains, and degree of salt dissociation. Also, the electrostatic interactions between cations and anions are more important in aprotic liquids (including polymer solvents) than in protic solvents. The anion has a major influence on both phase composition and ion pairing, which affects the ionic conductivity, with the anions being barely solvated by the aprotic solvents [1]. In the late 1970s, numerous alkaline and alkaline earth salts of Brønsted nitrogen superacids were synthesized, known especially as bis[(perfluoroalkyl)sulfonyl]imide salts (or sulfonimides for short) [35]. The simplest member of this family is the bis[(trifluoromethyl)sulfonyl]imide acid, (CF3SO2)2NH (or protonated form of the bis ((trifluoromethane sulfonyl) imide anion), HTFSI), which has a pKa value of 7.8 in acetic acid compared with 10.2 for HNO3 (acidity increases as pKa decreases). It is truly a superacid in the gas phase, according to measurements using pulsed FT ion cyclotron resonance mass spectrometry (FTICR-MS) method performed by Koppel et al. [36], and is several orders of magnitude more acidic than traditional strong acids such as CF3SO3H, FSO3H, HI, etc. [36,37]. Even more, the superior homologs of HTFSI such as (C2F5SO2)2NH (or the protonated form of the bis ((perfluoroethane sulfonyl) imide anion), HBETI), CF3SO2NHSO2CnF2n1 (n  2–4) and (C4F9SO2)2NH exhibit an increasing acidity with an increase in the size of the perfluoroalkyl groups in the molecular formula, the last member having the highest acidity known in gas phase [36]. The remarkable acidity of these sulfonimides is due, in part, to the resonance stabilization of the conjugate base anions of the acids. Extensive delocalization of the charge over the SO2–N–SO2 framework, assisted by the withdrawing inductive effect of the perfluoroalkyl groups, makes the conjugate base anion of the acid very weak and highly resonance-stabilized [37]. Crystal structures of some of these sulfonimide salts confirmed the presence of such a delocalization [38]. In these salts, the negative charge is centered mainly on the oxygen atoms rather than on the central nitrogen atom. Cation contacts with oxygen in the solid are generally observed but, depending on the nature of the cation, nitrogen–cation contacts are also noticed. In addition, even in the acidic form, the single S–N bond has a length of 1.644 Å, which is significantly shorter than that of the normal S–N single bond (1.74 Å ) [39]. Many lithium salts, including sulfonimide salts, were tested for the preparation of SPEs, but Armand and Elkadiri [40] proposed criteria for solubility, ionic conductivity, and redox stability for the polymer electrolytes that imposed certain restrictions on the salts. So far, of all the salts used, only LiClO4, CF3SO3Li (LiTf), (CF3SO2)2NLi (LiTFSI), and (CF3SO2)2CHLi (LiTFSM) satisfied the

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electrochemical stability criteria, although new perfluoroalkyl compounds with different acidic functionalities and larger monovalent (CnF2n1SO3Li with n  4, 6, 8; (CF3CO2)2NLi) or divalent (LiOOC(CF2)3COOLi, LiSO3(CF2)3SO3Li) anions have also been tested [10]. Recent research on new lithium salt-based SPEs has focused on lithium salts with low lattice energies, especially lithium sulfonimides. Salts consisting of a large anion with delocalized negative charge are particularly desired, and among such salts, those based on the bis[(trifluoromethyl)sulfonyl]imide anion (e.g., (CF3SO2)2NLi or LiTFSI) have been particularly emphasized [17,29,34,41–46]. As discussed earlier, the TFSI anion has a highly delocalized negative charge and very low basicity, which leads to good salt dissociation, less ion pairing, and high ionic conductivity in PEO-based electrolytes [34,43,47]. It is also thought that the TFSI anion exhibits a plasticizing effect on the PEO host, which could contribute to higher SPE conductivity [48]. Most prior work on lithium salt SPEs has concentrated on monoanionic salts [11,14,49,50]. Also, some work has been done using etheric dilithium salts based on LiTFSI motifs such as [CF3SO2N(Li)SO2(CF2)2O(CF2)2, 4]2 [51] or aromatic [(polyfluoroalkyl)sulfonyl] dilithium salts such as C6H4(COTFSMLi)2 [52]. Dimeric and oligomeric anions in particular are attractive, particularly in battery technology, because the contribution of the anion to the overall ionic conductivity is diminished in such salts, thereby increasing the lithium transference number [10,52]. High lithium transference is desirable in a battery because it helps to diminish concentration polarization of the salt, which can lead to diminished performance and premature device failure. A more recent approach involves the dispersion of selected ceramic nanofillers such as fumed silica, alumina, or titania into the polymer host to increase both its stability and mechanical properties, especially for the gel-type electrolytes [13,53]. It has been shown that ceramic fillers, when selected properly on the basis of their nature and particle size, may greatly influence the characteristics and properties of polymer electrolytes of different types [54–56]. The present chapter focuses both on the synthesis of novel lithium salts based on polyanions with structures similar to that of LiTFSI [57–60] and the structural, thermal, and electrochemical characterization of SPEs prepared using these salts in polyether hosts [51,61,62]. The effect of cross-linking on ionic conductivity is also explored for several of the new lithium salts. The polyanion in each of the new salts consists of two discrete end-units based on LiTFSI motifs that are connected together by a polyanionic [(perfluoroalkylene)disulfonyl]imide oligomeric chain of variable length, as can be seen from Scheme 1. As polymer hosts, either high-molecular-weight PEO (Mw  4 106 Daltons) or low molecular weight polyethylene glycol (PEG) (Mw  2 103 Da) cross-linked with a solution of 27 wt% 4,4,4-methyllidynetrisphenylisocyanate in ethyl acetate (commercial name Desmodur RE) was used.

Fluorinated electrolytes based on lithium salts of strong Brønsted acids

F F

O

O

C

S

F

O

NLi+

F

S

C

O

F

O S x

n=0

O

-

N Li+

O

F

S

C

O n

F

227

F

LiTFSI

n=1

(x = 2, 4, 6, 8)

Dimer

n=3

(x = 4, 6)

Tetramer

n≈5

(x = 4, 6)

Hexamer

n ≈ 17

(x = 4, 6)

Octadecamer

n ≈ 225

(x = 4)

Polymer

Scheme 1. The new lithium polyanionic salts used in SPEs preparation.

2. NEW LITHIUM SALT SYNTHESIS Scheme 2 describes the general synthetic route for obtaining the new dilithium salts (n  1 and x  2, 4, 6 ,8), in which the two necessary reactants, an ,ωdisulfonylfluoride compound, FSO2(CF2)xSO2F, and CF3SO2N(Na)Si(CH3)3, are prepared in parallel [37,57,59–64]. First, the commercially available ,ωdiiodoperfluoroalkylene, I(CF2)xI (x  2, 4, 6, 8), was dissolved in a 1:1 (v/v) mixture of acetonitrile and distilled-deionized (DI) water; this mixture was then added slowly into an aqueous solution of Na2S2O4 and NaHCO3 at 10°C such that the molar ratio of I(CF2)xI/Na2S2O4/NaHCO3 was 1:3.2:5.5. The reaction was completed after stirring at room temperature for 2 days with the formation of a disulfinate sodium salt, NaSO2(CF2)xSO2Na. deionized (DI) water was poured into another flask, cooled to 0°C using an ice – salt mixture, and chlorine gas was bubbled through the solution until saturation was reached. The reaction mixture was added slowly to the flask containing the chlorine-saturated water while chlorine gas was still bubbled vigorously throughout the addition. A white precipitate was formed. This mixture was filtered and the white solid was dried in air at room temperature for 1 h. The solid was further purified by sublimation at 60°C under dynamic vacuum leaving a white solid, ClSO2(CF2)xSO2Cl. The next step was to stir this dichloride with an excess KF in dry acetonitrile at room temperature for 4 days, then the temperature was increased to 90°C for another 2 days, during 19F-NMR showed the reaction to be complete. The reaction mixture was filtered through a Celite layer and the product was isolated from filtrate by the addition of excess DI water. The product was dried over P4O10 and then distilled under dynamic vacuum at room temperature.

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CF3SO3H P2O5 I - (CF2)x - I

(CF3SO2)2O NH3

Na2S2O4, NaHCO3/CH3CN/H2O CF3SO2 - NH2

NaSO2 - (CF2)x - SO2Na

NaOH/H2O

Cl2/H2O/CH3CN

CF3SO2 - NHNa

ClSO2 - (CF2)x - SO2Cl 1. KF/CH3CN 2. H2O

xs HMDS/CH3CN CF3SO2 - N(Na)Si(CH3)3

FSO2 - (CF2)x - SO2F 1.0 equiv CH3CN

2.5 equiv - (CH3)3SiF

CF3SO2N(Na)-SO2(CF2)xSO2N(Na)-SO2CF3 1. Nafion/H2O 2. Li2CO3/H2O CF3SO2N(Li)-SO2(CF2)xSO2N(Li)-SO2CF3

Scheme 2. General synthetic route for the new dimeric lithium salts (n  1) used in SPEs preparation (x  2, 4, 6, 8).

The starting material for the end-capping reactant was the triflic acid, CF3SO3H, to which was added P4O10 in small portions at 0oC until a molar ratio of triflic acid/P4O10 of 2:1 was achieved; the stirring was continued until room temperature was attained. Next, the triflic anhydride, (CF3SO2)2O, was distilled off the mixture under dynamic vacuum at 120°C. Liquid ammonia (excess) was condensed in a flask at 80°C and the triflic anhydride was added slowly. After the reaction was complete (about 3 h), excess ammonia was removed by allowing the flask to warm to room temperature. The resulting sulfonylimide, CF3SO2NH2, was obtained by sublimation under dynamic vacuum, by increasing the temperature from 80 to 100°C during a 1 h period. The solid sulfonylimide was titrated at room temperature with 0.1 M NaOH in DI water, until a pH of

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8.4 was reached. The water was removed by rotary evaporation under reduced pressure at 40°C and the solid sodium salt, CF3SO2NHNa, was dried further under dynamic vacuum at 120°C for 16 h. The next step was to stir the dried sodium salt with an excess 1,1,1,3,3,3-hexamethyldisilazane (CH3)3SiN(H)Si(CH3)3 (known as HMDS) in anhydrous acetonitrile for 2 days at 90°C. Both the solvent and unreacted HMDS were removed under dynamic vacuum at ambient temperature, resulting in the end-capping agent, CF3SO2N(Na)Si(CH3)3. For the coupling step, the quantity of FSO2(CF2)xSO2F obtained previously was dried over P4O10 for 2 h and condensed through the vacuum line into a chilled stainless-steel cylinder, followed by condensation in anhydrous acetonitrile. Next, an excess quantity of CF3SO2N(Na)Si(CH3)3 was transferred into the same cylinder, which was sealed and placed on a shaker and heated at 120°C for 6 days. The end of the reaction and the structure of the final product (including the x-value) were checked by 19F-NMR based on end-group analysis. The solvent was removed and the resulting sodium salt, CF3SO2N(Na)SO2(CF2)xSO2N(Na)SO2CF3, was dried under dynamic vacuum at 100°C for 1 day. Further, ion exchange was used to obtain the acid form by dissolving the sodium salt in DI water and passing it through a Nafion®-H beads column for 8 h. The product was then washed with DI water and collected. The solvent was removed by rotary evaporation under reduced pressure, and the resulting acid form was first dried at 100°C for 1 day, then sublimed at 130°C under dynamic vacuum for 2 days, to remove any excess of CF3SO2NH2. In the final step, the dimeric acid was dissolved in DI water and titrated at room temperature, using a pH meter, with a saturated aqueous solution of Li2CO3 to the endpoint of 7.10. The solvent was removed by rotary evaporation under reduced pressure and dried at 100°C under dynamic vacuum for 2 days. The Li salt was obtained without further purification. Scheme 3 presents a step-growth polymerization that uses an exact stoichiometry of two difunctional monomers, bis[(perfluoroalkylene)sulfonylfluoride] (FSO2–(CF2)x–SO2F) and disodium bis[(perfluoroalkylene)sulfonylimide trimethylsilane] ((CH3)3Si(Na)NSO2–(CF2)x–SO2N(Na)Si(CH3)3) to yield the desired oligomeric salts for n  3 and x  4, 6 [57,58,62,64]. To obtain the latter of the two monomers, the appropriate quantity of the first monomer was reacted with excess liquid NH3 at 80°C and the mixture was allowed to warm up to room temperature. The resulting product was obtained by sublimation under the same conditions as described previously for the monofunctional (trifluorometanesulfonyl)imide, CF3SO2NH2. Next, it was reacted with an aqueous solution of HCl to remove the ammonia completely, then the water was removed by rotary evaporation under reduced pressure at 50°C, obtaining a white solid bis[(perfluoroalkylene)sulfonylimide] acid, H2NSO2–(CF2)x–SO2NH2. The solid sulfonylimide was then titrated at room temperature with 0.1 M NaOH in DI water, until a pH of 8.4 was reached. Water was removed by rotary evaporation under reduced pressure at 40°C and the solid sodium salt,

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FSO2 - (CF2)x - SO2F xs NH3/-NH4F H4N(H)NSO2 - (CF2)x - SO2N(H)NH4

I - (CF2)x - I

HCl/-NH4Cl

Na2S2O4, NaHCO3/CH3CN/H2O

H2NSO2 - (CF2)x - SO2NH2

NaSO2 - (CF2)x - SO2Na

NaOH/CH3OH

Cl2/H2O/CH3CN ClSO2 - (CF2)x - SO2Cl

H(Na)NSO2 - (CF2)x - SO2N(Na)H

1. KF/CH3CN 2. H2O

xs HMDS/CH3CN (CH3)3Si(Na)NSO2 - (CF2)x - SO2N(Na)Si(CH3)3

FSO2 - (CF2)x - SO2F 2.5 equiv CH3CN

1.0 equiv

(A)

- (CH3)3SiF

FSO2(CF2)xSO2 - N(Na)SO2 - (CF2)x - SO2N(Na) - SO2(CF2)xSO2F

(B)

xs CF3SO2N(Na)Si(CH3)3/CH3CN CF3SO2N(Na) - [SO2(CF2)xSO2N(Na)]3 - SO2CF3 1. Nafion/H2O 2. Li2CO3/H2O CF3SO2N(Li) - [SO2(CF2)xSO2N(Li)]3 - SO2CF3

Scheme 3. Synthetic scheme for tetrameric lithium salts (n  3 and x  4, 6) preparation.

(Na)HNSO2–(CF2)x–SO2NH(Na), was dried under dynamic vacuum at 120°C for 16 h. The next step was to stir the dried sodium salt with an excess HMDS in anhydrous acetonitrile for 2 days at 90°C. Both the solvent and unreacted HMDS were removed under dynamic vacuum at ambient temperature, producing the monomer, (CH3)3Si(Na)NSO2–(CF2)x–SO2N(Na)Si(CH3)3. For the coupling step, some of the FSO2(CF2)xSO2F obtained previously was dried over P4O10 for 2 h and condensed through the vacuum line into a stainlesssteel cylinder, followed by condensing in anhydrous acetonitrile. A mass of

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231

(CH3)3Si(Na)NSO2–(CF2)x–SO2N(Na)Si(CH3)3 (1:2.5, mol/mol) was transferred to the same cylinder, which was sealed and placed on a shaker and heated at 120°C for 6 days. The end of the reaction and the structure of the final product were checked by 19F-NMR based on end group analysis. For end capping, the quantity of FSO2(CF2)xSO2–(Na)NSO2(CF2)xSO2N(Na)–SO2(CF2)xSO2F obtained in the previous step was reacted with an excess quantity of CF3SO2N(Na)Si(CH3)3 (molar ratio of 1:2.2) following the same procedure as in Scheme 2. After solvent removal, the resulting sodium salt, CF3SO2N(Na)[SO2(CF2)xSO2N(Na)]3SO2CF3, was dried under dynamic vacuum at 100°C for 1 day. Further, to get to the acid form, the tetrameric sodium salt was treated in the same way as the dimeric salt in Scheme 2. The synthesis of the superior homologues (n  3, x  4, 6) is presented in Scheme 4 [62,64] As can be seen, these homologues of the polyanionic series were prepared by a stoichiometric step-growth polymerization that alternates consecutive reactions with (A) and (B) from Scheme 3 ((CH3)3Si(Na)NSO2– (CF2)x–SO2N(Na)Si(CH3)3 and F[SO2(CF2)xN(Na)]2–SO2(CF2)xSO2F respectively) until the desired n-value was reached. After end capping, the resulting salts (in Na form) were acidified by several passes (as aqueous solutions), through a Nafion® column in acid form. 19F-NMR and thermogravimetric analysis (TGA) were used to monitor complete conversion to the acid form. The acidic form was converted into the Li form by titration with an aqueous solution of Li2CO3 to an endpoint of 7.4. Finally, the water was removed by rotary evaporation under reduced pressure and the salt was dried at 100°C under dynamic vacuum for 2 days. The Li salt was obtained without further purification. Prior to use, LiTFSI was dried for 24 h at 150°C and then 1 h at 170°C under dynamic vacuum (2 102 torr), while all the other lithium salts were dried for only 24 h at 100°C under the same vacuum. All CH3CN F-[SO2(CF2)xSO2N(Na)]4m+2-SO2(CF2)x-SO2F

m (A) + (m+1) (B) - (CH3)3SiF

CF3SO2N(Na)Si(CH3)3/CH3CN 2 equivalent

CF3SO2N(Na)-[SO2(CF2)xSO2N(Na)]4m+3-SO2CF3 1. Nafion/H2O 2. Li2CO3/H2O CF3SO2N(Li)-[SO2(CF2)xSO2N(Li)]4m+3-SO2CF3

Scheme 4. Generalized synthetic scheme for oligomeric and polymeric lithium salts preparation.

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reactions were carried out under rigorous dry conditions until there seemed to be no further conversion.The progress of each reaction was monitored by 19F-NMR. For n  1, 3, 5 the compound is monodisperse but for n ⬇ 17 and 225, the materials are polydisperse with an expected low polydispersity for n ⬇ 17 and a higher polydispersity for n ⬇ 225. Determining Mn and Mw values to obtain a polydispersity index is very difficult for charged polymers of this type due to uncertainties regarding polymer conformation in solution and a lack of suitable standards, and was not attempted. The average n values were determined by endgroup 19F-NMR by comparing the integral of the terminal CF3SO2– groups to the –SO2CF2– groups, and the –CF2CF2– internal groups. The structure and the purity of each of the new dilithium salts were checked by 19F- and 1H-NMR. The presence of any impurities such as –CF2SO2NH2 endgroups (after acidification) or –CF2SO2F endgroups is easily detected by their different chemical shifts in 19F-NMR. Similarly, the presence of –CF2SO2NH2 in the presence of –SO2–N(H)–SO2– groups is easily detected by 1H-NMR under dry conditions. The 19F-NMR spectrum corresponding to the dimeric salt (n  1, x  4) is presented in Fig. 1 as a representative spectrum. It shows three signals; the two –CF3 endgroups (78.9 ppm, s) have a ratio of 6:4 to both the two

Fig. 1. 19F-NMR spectrum for the dimeric lithium salt (n  1) with x  4.

Fluorinated electrolytes based on lithium salts of strong Brønsted acids

233

–SO2CF2– groups (112.3 ppm, m) and the two –CF2CF2–internal groups (119.4 ppm, m). For the octadecamer (n ⬇ 17, x  6) the 19F-NMR is presented in Fig. 2 and shows four signals: –CF3 endgroups (78.9 ppm, s), –SO2CF2– groups (112.3 ppm, m), –CF2CF2– groups (119.4 ppm, m) and –CF2CF2CF2– groups (120.8 ppm, m) with a ratio of the endgroups to the internal groups of approximately 6:68:68:68. 3. SPE PREPARATION PEO-based electrolytes were fabricated by casting from N, N-dimethylformamide solutions. DMF was used in place of acetonitrile (the more commonly used solvent for casting PEO-based SPEs of lithium salts) because many of the salts in Scheme 1 exhibited low solubility and precipitated during evaporation from acetonitrile solutions. Casting solutions were prepared by dissolving the appropriate amounts of high-molecular weight PEO, Mw  4 106 Da (with the required fraction of polymer to salt to achieve the desired EO/Li ratio of 30:1) in DMF solvent. The solutions were mixed for 48 h at 50°C and then poured into a poly(tetrafluoroethylene-co-perfluoropropylvinylether) (PFA) dish. The solvent was slowly removed by evaporation

Fig. 2. 19F-NMR spectrum for the octadecameric lithium salt (n ⬇ 17) with x  6.

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with the dish placed in a vacuum oven held at 70°C under dynamic vacuum. Mechanically and dimensionally stable SPE membranes (usually 100–200 μm) were obtained by peeling the dried material from the dish after solvent removal. Cross-linked PEG-based electrolytes were prepared using LiTFSI and some of the new salts in concentrations corresponding to an EO/Li ratio of 30:1. First, the necessary amounts of low-molecular weight PEG (Mw2103 Da) and Li salt were mixed together in the dry box for 1 h at 70°C on a hot plate. A stoichiometric quantity of cross-linker (4,4,4-methyllidynetrisphenylisocyanate) was added to PEG and the mixture was stirred for ½ h at the same temperature. Next, for the cross-linking to take place, the mixture was pressed between two Teflon sheets, and the entire assembly was introduced for 2 h into a vacuum oven at 90°C. Dimensionally stable elastomeric SPE membranes (usually 400–500 μm in thickness) were obtained by removing the Teflon sheets. 4. RESULTS AND DISCUSSION 4.1. PEO-based solid polymer electrolytes

Thermal properties of SPEs, such as glass transition temperature (Tg), melting temperature (Tm), specific fusion enthalpy (Hf), and differential molar heat capacity (Cp) at the glass transition temperature were determined from modulated differential scanning calorimetry (MDSC) thermograms for all SPE samples using previously described methods [51,61,62]. The values for specific fusion enthalpy and differential molar heat capacity at the glass transition temperature are reported after correction for the salt content of the SPEs (Hfcorr and Cpcorr); the salt is considered entirely dissolved in the polymeric host and therefore completely amorphous. The degree of crystallinity (χDSC) was calculated for each SPE as a ratio of the corrected heat of fusion (Hfcorr) to the heat of fusion for pure PEO. This value was calculated to be 165 J/g by extrapolating linearly from an XRD-determined value of 80% crystallinity, corresponding to a DSC-determined heat of fusion of 135 J/g, to a hypothetical 100% crystallinity for PEO only. The value of 165 J/g compares well with a heat of fusion of 185 J/g obtained by Armand et al. [47] for mixtures of LiTFSI and high-molecular weight PEO, though it is somewhat smaller than the value of 210 J/g reported by Booth et al. [16] for perfectly crystalline poly(oxyethylene), or the value of 207 J/g calculated based on the polymer-melting temperature according to an empirical correlation used by Kim et al. [65]. This difference could be due to the fact that our PEO sample was first heated to 120°C, slowly cooled to room temperature, following the same thermal treatment as the SPEs subjected to electrochemical impedance spectroscopy (EIS) analysis and afterwards used in a DSC measurement. Tables 1 and 2 list the values of Tg, Tm, Hfcorr, and χDSC obtained from DSC thermograms for SPEs prepared using either dilithium salts (n  1) or polyanionic lithium salts (x  4) compared with the LiTFSI-based electrolytes. The

Fluorinated electrolytes based on lithium salts of strong Brønsted acids

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Table 1 DSC characteristics and the degree of crystallinity (χ) for SPEs made of LiTFSI or a dimeric lithium salt (n  1) and PEO, EO/Li  30:1 (Reprinted from Geiculescu et al. [61]. With permission from Elsevier.) ΔHfcorr (J/g)

χDSC (%)

χXRD (%)

4.600

105

64

56

54

4.925

84

51

49

29

59

4.151

104

63

60

6

30

62

3.587

114

69

61

8

33

57

3.158

132

80

57

Salt type (x)

Tg (°C)

Tm (°C)

mPEO/mLi salt (g/g)

LiTFSI

54

59

2

36

4

corresponding corrected value of the differential molar heat capacity, Cpcorr σ, was calculated to be 11.6 0.7 J/mol K for all SPEs. As has been observed by Cowie et al. [66], variations in the thermal treatment of the sample did not appear to alter the Cp values significantly in the low-temperature region of interest, although some changes were noted in the Tg region. Ionic conductivity is strongly dependent not only on temperature but also on the structure of the polymer electrolyte. The PEO/lithium salt SPEs in general are found to have a complicated structure composed of complex crystalline phases with high melting points (PEO  salt), a PEO crystalline phase with a melting point around 60°C, and a PEO amorphous phase in which the inorganic salts are soluble. Furthermore, the composition of these phases changes with temperature [1]. The phase transitions in mixtures of PEO with LiTFSI were studied by Lascaud et al. [43], who constructed the phase diagram for this system, and involves a series of three intermediate crystalline compounds corresponding to LiTFSI weight fractions of 0.52, 0.68, and 0.76 or EO/Li ratios of 6:1, 3:1, and 2:1, respectively. In addition, as reported previously by different authors [34,67], the compound corresponding to a EO/Li ratio of 6:1 crystallizes only very slowly in the presence of an excess of PEO, which leads to a crystallinity gap over the range of 6:1 EO/Li 12:1. Also, the LiTFSI liquidus curve above 106°C in the phase diagram shows that at high temperatures LiTFSI is miscible in all proportions with PEO. Using this phase diagram, the EO/Li ratio for the new SPEs has been chosen to be 30:1 (diluted electrolytes). For each SPE X-ray diffraction (XRD) analysis was performed and the diffractogram for 2θ values between 10° and 30° was used to calculate the degree of crystallinity, χXRD, by comparing the magnitude of the Bragg peaks generated by the crystalline part with the proportion of diffuse scattering from the

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amorphous regions [51]. The XRD-measured degree of crystallinity was expressed as percentage of crystalline phase within the total sample, and its values for all SPEs are presented in Table 1 (dilithium salts-based electrolytes) and Table 2 (polyanionic lithium salts-based electrolytes). Comparing the crystallinity data from DSC with those from XRD for all the SPEs studied, we can observe that though they are close, the diffraction method always produces lower extents of crystallinity than the DSC method. This is probably because the XRD method rigorously considers only the regions of longrange order as being crystalline [68]. Also, as can be seen in Tables 1 and 2, all SPEs exhibited a degree of crystallinity of at least 50%. Ionic conductivity is the most significant property of these SPEs relating to their use in electrochemical power devices. The reproducibility of ionic conductivity measurements using EIS and our experimental setup has been established previously [51,62]. Figs. 3 and 4 present Arrhenius plots of log(ionic conductivity) vs. (reciprocal of absolute temperature) for dilute SPEs (EO/Li of 30:1) made from LiTFSI and some of the lithium salts illustrated in Scheme 1. The data are presented in two sets, one for a series of SPEs prepared from dilithium salts (n  1, x  2, 4, 6, 8; Fig. 3) and the other for a series of SPEs prepared from polyanionic lithium salts (x  4, n  1, 3, 5, 17, 225; Fig. 4). In each case, the data are presented both for the monomeric salt LiTFSI, and either the four dilithium salts in which two sulfonyl imide anions are linked by a perfluoroalkylene chain of variable length (–(CF2)x–), or the five polyanionic salts in which two LiTFSI motifs are connected by a variable number (n) of anionic repeating units (–[SO2(CF2)4SO2N(Li)]n–).

Table 2 DSC characteristics and the degree of crystallinity (χ) for SPEs made of LiTFSI or a polyanionic lithium salt (x  4) and PEO, EO/Li  30:1 (Reproduced from Geiculescu et al. [62]. With permission from The Electrochemical Society, Inc. ΔHfcorr (J/g)

χDSC (%)

χXRD (%)

4.600

105

64

56

59

4.151

104

63

60

24

66

3.952

125

76

54

5

27

67

3.881

134

81

63

17

28

66

3.815

100

61

54

225

25

69

3.777

101

61

50

Salt type (n)

Tg (°C)

Tm (°C)

0

54

59

1

29

3

mPEO/mLi salt (g/g)

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-2

log [κ κ (S/cm)]

-3 -4 -5 -6 -7

LiTFSI x=2 x=4 x=6 x=8

-8 2.6

2.8

3.0 1000/[T (K)]

3.2

3.4

Fig. 3. Arrhenius plots for SPEs made of LiTFSI or a dimeric lithium salt (n  1) and PEO, EO/Li  30:1. (Reprinted from Journal of Fluorine Chemistry with permission from Elsevier [61].) -2

log [κ (S/cm)]

-3 -4 -5 -6 -7

n=0 n=1 n=3 n≈5 n ≈ 17 n ≈ 225

-8 2.6

2.8

3.0

3.2

3.4

1000/[T (K)]

Fig. 4. Arrhenius plots for SPEs made of LiTFSI or a polyanionic lithium salt (x  4) and PEO, EO/Li  30:1. (Reproduced by permission of The Electrochemical Society, Inc. [62].)

As expected, all of the Arrhenius curves in Figs. 3 and 4 exhibit an abrupt change in conductivity near 60oC. This change is commonly seen in dilute PEObased SPEs [69,70] and is attributed to a crystalline melting/freezing transition of the PEO host, as seen from the melting temperatures in Tables 1 and 2. The transition is suppressed in all of the high-salt SPEs, as has previously been

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observed for LiTFSI SPEs in a EO/Li range near 10:1 [34,43]. This fact is undoubtedly related to the lack of crystallinity of all the TFSI-based SPEs (both the dilithium and polyanionic series ) in this salt concentration regime, less for the highest member of the polyanionic series (x  4, n ⬇ 225). Close inspection of the Arrhenius curves in the region between 60 and 120°C reveals a slight but significant curvature [71]. This curvature is also commonly observed in SPEs and other glassy ionic conductors, and is indicative of mechanical coupling between the motions of the charge carriers and the motions of the matrix. Curved Arrhenius plots may be fit using the semi-empirical relation of Vogel, Tammann, and Fulcher (VTF), written in the following form:

冤

B κ  A T 1/2 exp   (TT0)

冥

(1)

where A and B are phenomenological fitting parameters; A is related in a general way to the number of charge carriers, and B is related to either the apparent activation energy of ion transport [72,73], the expansivity of the polymer – salt mixture [25,74], or the renewal time of the matrix [75–78], depending upon the model. The term T0 is called the “ideal” or “equilibrium” glass transition temperature; it corresponds to the temperature at which the free volume disappears or the temperature at which the configurational entropy approaches zero, again depending upon the model. It is generally regarded as having a value of 20–50 K below the glass transition temperature [68]. In the case of LiTFSI-derived polyanionic salts we used a value 25 K below the glass transition temperature, which has been utilized previously for sulfonate- and sulfonimide-based SPEs [43,79]. The curved lines in Figs. 3 and 4 correspond to nonlinear least-squares fits of the data to the VTF equation using a value of T0 equal to 25° lower than the Tg values determined by DSC. For all of the SPEs studied, the T0 values together with the best-fit values of A and B are presented in Tables 3 (dilithium series) and 4 (polyanionic lithium series). Shriver et al. [80,81] have applied the configurational entropy model [72,73] to polymer electrolytes, expressing the term B in the VTF equation by a combination of terms as T0Sc* Δμ B   kB Tg ΔCpcorr

(2)

where kB is the Boltzmann constant, Sc* the minimum configurational entropy required for a cooperative rearrangement of a polymer chain segment involved in ion transport in the matrix (generally taken as kBln 2), Cpcorr the corrected molar heat capacity change at temperature Tg as the system moves from the glassy to the rubbery state, and μ the apparent activation energy opposing the rearrangement of the polymer segmental unit. This model has been applied to convert the

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Table 3 VTF parameters and the apparent activation energy of the segmental polymer motion (Δμ) for SPEs made of LiTFSI or a dimeric lithium salt (n  1) and PEO, EO/Li  30:1 (Reprinted from Geiculescu et al. [61]. With permission from Elsevier.) B (K)

Δμ σ (kJ/mol)

Salt type (x)

T0 (°C)

A (K1/2 S/cm)

Li salt concentration (wt%)

LiTFSI

79

3.03

808

15.2 1.0

17.9

2

61

0.25

559

10.5 0.7

16.9

4

54

0.14

520

9.7 0.6

19.4

6

55

0.65

573

10.5 0.7

21.8

8

58

1.41

667

12.6 0.8

24.1

Table 4 VTF parameters and the apparent activation energy of the segmental polymer motion (Δμ) for SPEs made of LiTFSI or a polyanionic lithium salt (x  4) and PEO, EO/Li  30:1 (Reproduced from Geiculescu et al. [62]. With permission from The Electrochemical Society, Inc.) Salt type (n)

T0 (°C)

A (K1/2 S/cm)

B (K)

Δμ σ (kJ/mol)

Li salt concentration (wt%)

0

78.5

3.03

808

15.4 1.4

17.9

1

53.9

0.14

520

9.8 0.9

19.4

3

48.7

0.14

475

8.9 0.8

20.2

5

51.5

0.15

437

8.2 0.7

20.5

17

53.3

0.17

390

7.3 0.6

20.8

225

49.5

0.02

449

8.4 0.7

20.8

B term of the VTF equations into apparent activation energies, μ σ (in kJ/mol), which are listed in Table 3 (for the dilithium series) and Table 4 (for the polyanionic lithium series). The configurational entropy model has been used to describe ionic conductivity in polymers prepared from oligo(ethylene oxide) macromers [82], polymer electrolytes based on phosphate and polyether copolymers [83], polymers based on high molecular weight PEO [51,80], in poly(itaconates) with poly(propylene glycol) side chains [84], etc.

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A central observation regarding the data in Figs. 3 and 4 and the calculated parameters in Tables 3 and 4 is that SPEs from all of the new lithium salts exhibit conductivities that are suppressed relative to that of the monolithium salt LiTFSI at all temperatures. This finding is probably a consequence of the fact that the conductivity in most lithium salt-based SPEs is dominated by anion transport [69,85–88], and anion motion is expected to be diminished for these larger anions due to increased friction or entanglement with the matrix. This effect would cause diminished conductivity in two ways: first, because the anions themselves are less effective in transporting charge, and second because the slow-moving anions could act to retard cation transport by coulomb trapping [51,62]. Ionic conductivity increases with anion size for the dilithium salt-based SPEs with the notable exception of the electrolyte prepared using the dimeric salt in which the linker group is a perfluoroethylene chain (x  2). In the latter case, conductivity is intermediate between that of SPEs made from salts with x  4 and 6. This behavior may be due to the fact that the shorter linker between the two sulfonimide anions allows them to overcome the electron-withdrawing inductive effects of fluorine atoms belonging to the linker, thereby forming two separate, non-interacting anions linked together. All of the other dianionic salts are expected to have more of a delocalized dianion behavior. On the other hand, the smaller overall size of the dianion in salt with x  2 may mean that the diminished electron-withdrawing effects of the perfluoroethylene linker (lower delocalization) are partially offset by the fact that the smaller anion should experience less viscous drag and thus contribute more to the ionic conductivity. An analogous observation regarding the influence of anion size upon conductivity was noted in our earlier publication, albeit only for some of the aforementioned dilithium salts and some dilithium etheric salts based also on LiTFSI motifs [51]. Comparing the conductivity data for polyanionic salt-based SPEs in the region between 60 and 120°C (Fig. 4 and Table 4), where the SPEs are melted and in a single amorphous phase, reveals some unexpected trends with respect to the number of anionic repeating units (n). The ionic conductivity increases with increasing n up to n ⬇ 17 for EO/Li of 30:1 after which it decreases with n, with the lowest conductivity value corresponding to n ⬇ 225. We believe that this behavior is the result of several competing factors. The increase in n generates larger anions with a lower mobility, which make a diminished contribution to the overall ionic conductivity. At the same time, the larger anions have a better charge delocalization, which increases conductivity by reducing the ion pairing for some of the higher members of the polyanionic series. Ion transport is also favored for the high-n members of the polyanionic series through the decrease of the apparent activation energy of the segmental motion (up to n ⬇ 17 for EO/Li  30:1). Apparently, the balance of these effects yields optimal conductivities with respect to anion size for the 30:1 EO/Li SPEs. In the case of the SPEs made from salts with n ⬇ 225, there is the additional factor of phase segregation,

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resulting in salt-rich regions near each individual anion chain with salt-poor regions between them. Ionic conductivity in such a material will be limited by the conductivity between the salt-rich domains and is expected to be low relative to that of an otherwise similar SPE that is not phase-segregated. Consideration of the A and B values and apparent activation energies in Tables 3 and 4 from the VTF fits also lends some insight. The primary difference between the various salts lies in the A values, which are highest for LiTFSI but are substantially diminished for all of the dilithium salts. The lower A values suggest that the major cause of the diminished conductivity is that there are fewer mobile charge carriers present in both the dilithium and polyanionic lithium salt SPEs, which in turn is probably a consequence of the fact that the anions have lower mobility and can partially trap the cations. A secondary effect is that in all of the new lithium salt SPEs, the B term from the VTF fits and the apparent activation energy of the segmental motion, μ, are diminished relative to those for LiTFSI, which may indicate that these polyanions have affected the local microstructure of the host polymer, acting as plasticizers, so as to diminish the apparent activation energy for ion transport and the local expansivity of the matrix. We note that diminished ionic conductivity due to suppressed anion motion is not necessarily a problem for lithium battery technology; in fact, it can be advantageous since it will probably lead to higher lithium transference numbers and diminished problems with cell polarization caused by electrolyte “pooling” on one side of the battery. Furthermore, measurements of cationic transference number (t) and apparent lithium salt diffusion coefficient (Ds) have been performed at the same temperature of 90°C for the diluted SPEs (EO/Li  30:1) based on either dimeric lithium salts (n  1, Table 5) or polyanionic lithium salts (x  4, Table 6). The ionic conductivities at 90°C were calculated for all SPEs using the VTF equation and parameters A, B and T0 from Tables 3 and 4. The salt diffusion coefficients were calculated from the slopes of the linear portions of open-circuit potential (OCP) relaxation curves after applying the restricted diffusion method (galvanostatic polarization/current interruption) to the electrolytes, using the following equation, which assumes that the potential difference is proportional to the concentration difference [70,89,90]:

π 2Ds 1n(ΔΦ)   t  A2 δ2

(3)

where Ds is the differential lithium salt diffusion coefficient (cm2/s), t the relaxation time (s), δ the film thickness (cm), and A2 a constant. The aim of the current interrupt experiments is to establish concentration gradients at the electrode surface without allowing the concentration boundary layers to propagate to the center of the cell. In order to fulfill this requirement, the upper limit for the

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Table 5 Transport properties determined at 90°C for SPEs made of LiTFSI or a dimeric lithium salt (n  1) and PEO, EO/Li  30:1 Salt type (x)

(Ds σ)108 (cm2/s)

κ103 (S/cm)

t

Li salt concentration (wt%)

LiTFSI

4.2 0.8

1.33

0.15

17.9

2

2.9 0.5

0.32

0.15

16.9

4

1.9 0.7

0.20

0.16

19.4

6

3.1 0.5

0.66

0.16

21.8

8

5.3 0.4

0.82

0.16

24.1

Table 6. Transport properties determined at 90°C for SPES made of LiTFSI or a polyanionic lithium salt (x  4) and PEO, EO/Li  30:1 Salt type (n)

(Ds σ)108 (cm2/s)

κ103 (S/cm)

t

Li salt concentration (wt%)

0

4.2 0.8

1.33

0.15

17.9

1

1.9 0.7

0.20

0.16

19.4

3

1.6 0.3

0.24

0.16

20.2

5

1.3 0.3

0.34

0.17

20.5

17

0.5 0.06

0.58

0.17

20.8

225

0.4 0.08

0.04

–

20.8

polarization time (ti) is limited by the well-known condition for one-dimensional semi-infinite diffusion, ti

δ 2/Ds. Lithium transference numbers were determined using the EIS method [8,91,92]. It used blocking electrodes for the anion and non-blocking electrodes for the lithium cation, so that the experimental cell had the form Li⏐(polymer)nLiX⏐Li. In this case, the complex impedance spectrum at low temperatures consisted of three arcs [8,93], which, using the equivalent Randles circuit theory for a SPE with non-blocking electrodes, could be attributed to the various physical/chemical phenomena taking place. For all samples, the complex impedance spectrum at 90°C consisted only of the intermediate and lower frequency arcs, while the high-frequency arc vanished due to the fact that at

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temperatures above the melting point of the polymer host (PEO), the capacitive component of the electrolyte is apparently negligible or not visible at frequencies lower than 1 MHz [94]. Although battery SPEs generally have high salt concentration with non-ideal behavior and present strong ionic interactions, the theory for dilute electrolytes has been used either for its simplicity or, as is the case now, when dealing with low salt concentrations. Therefore, for the dilute electrolytes t was calculated according to the following equation [93,95]: Rb t   RbZd

(4)

where Rb corresponds to the bulk resistance of the electrolyte and Zd a diffusioncontrolled impedance called the Warburg impedance. From Tables 5 and 6 it can be seen that the cationic transference numbers are practically the same for all of the salts tested, varying only between 0.15 and 0.17, almost independent of the number of perfluoromethylene repeating units in the dianionic salts and the number of anionic repeating units in the polyanionic salts. The apparent salt diffusion coefficients Ds follow the same trend observed for ionic conductivities κ at 90°C, which seems to suggest that the degree of nonideality introduced by the mean molar activity coefficient corresponding to SPEs prepared from the dimeric salts does not differ too much from one salt to another. These effects are somewhat unusual in the sense that one might expect that increasing anion size should further suppress the anion contribution to the overall ionic conductivity, thereby increasing the cationic transference number and diminishing the conductivity. The apparent salt diffusion coefficients Ds decrease with an increase in the number of anionic repeating units (which is to be expected) but does not follow the same trend observed for ionic conductivities at 90°C, which seems to suggest that the degree of non-ideality, introduced by the mean molar activity coefficients ( f ) corresponding to SPEs prepared from the polyanionic series lithium salts, has a serious influence (but not one easy to quantify) upon the transport properties of these electrolytes. 4.2. Cross-linked PEG-based solid polymer electrolytes

The conductivity of high-molecular-weight PEO-based electrolytes at room temperature is reduced by crystallization of the polymer, especially for the lowsalt-content SPEs [96]. Several routes have been investigated to prevent crystallization, such as block copolymerization [97,98], grafting [99,100] or cross-linking [15,101], which allows one to incorporate PEO into a macromolecular sequence that will resist host crystallization. PEG cross-linking was reported to give the polymer a rubber-like texture and a large amorphous phase that enhances its ionic conductivity and prevents the material from creeping [15]. Cheradame et al. [102] have concentrated on

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forming chemically cross-linked SPEs, which both inhibited crystallization and exhibited good mechanical properties to be fabricated as strong films or membranes. Several other groups have also studied network SPEs based on PEO linked by isocyanates [29,103,104]. Consistently, conductivities of the network systems were some two orders of magnitude lower than those recorded for salts just dissolved in the non-cross-linked polymeric host. This observation reflects the fact that a cross-linked matrix can impede ion transport more than a noncross-linked one. LiTFSI, dilithium salts (n  1, x  2, 4, 6, 8), and some of the polyanionic lithium salts (x  6, n  1, 3, 17) were used to prepare SPEs by dissolving the salts in the low-molecular-weight PEG (Mw  2000 Da) and then cross-linking the PEG with 4,4,4-methylidynetrisphenylisocyanate. The same EO/Li ratio as for the PEO-based SPEs was used for the preparation of these electrolytes (30:1). Following the procedures already described, Arrhenius curves (Figs. 5 and 6) were obtained for all the samples by impedance spectroscopy and thermograms by modulated differential scanning calorimetry. Thermal properties such as Tg, Tm, Hfcorr and χDSC were determined from MDSC thermograms and both Hf and Cp values were corrected for the salt content. As expected, LiTFSI-based electrolytes exhibited the highest ionic conductivity for the entire temperature and concentration ranges. Also, the cross-linked PEG-based electrolytes exhibited lower conductivities (up to one and a half order of magnitude) when compared with the PEO-based SPEs over the entire temperature range, due to the reduction of the contribution of segmental motion of the polymer to the overall ionic conductivity. -3.0 -3.5

log [κ (S/cm)]

-4.0 -4.5 -5.0 -5.5 -6.0

LiTFSI x=2 x=4 x=6 x=8

-6.5 2.6

2.8

3.0 1000/[T (K)]

3.2

3.4

Fig. 5. Arrhenius plots for SPEs made of cross-linked PEG and LiTFSI or a dimeric lithium salt (n  1), EO/Li  30:1.

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245

-2

log [κ (S/cm)]

-3 -4 -5 -6 -7

LiTFSI n =1 n =3 n = 17

-8 2.6

2.8

3.0 1000/[T (K)]

3.2

3.4

Fig. 6. Arrhenius plots for SPEs made of cross-linked PEG and LiTFSI or a polyanionic lithium salt (x  6), EO/Li  30:1.

From Fig. 5, for diluted SPEs, we can see that the electrolyte based on LiTFSI still has the highest conductivity followed in order by those prepared using the dilithium salts with x  8, 6, 2 and 4, in the same order as seen for the PEObased SPEs using the same salts and the same EO/Li ratio (Fig. 3). In this case, no transition in conductivity is observed at around 58°C, the melting temperature of PEG, which indicates that all these SPEs are probably completely amorphous, which is confirmed by the corresponding DSC thermograms. This is a major difference compared with the behavior of non-cross-linked PEO-based SPEs using the same salts [51,105], where such a transition was observed at around 60°C and was attributed to a crystalline melting/freezing transformation of the PEO host. Therefore, for dilute SPEs, the increasing ionic conductivity with the higher fluorine content in the salt molecule indicates an increase in anion size, while decreasing the anion contribution to the overall conductivity [106] increases the same conductivity by decreasing the anion basicity through electron delocalization, which leads to a good salt dissociation and less ion pairing. SPEs prepared using the polyanionic lithium salts exhibited ionic conductivities that declined monotonically with an increase in the number of anionic repeating units, n, from 1 to 3 to 17, over the entire temperature range (Fig. 6). Again, all of the Arrhenius curves reveal a slight curvature over the temperature region of interest (60 to 120°C), common to polyether-based electrolytes, which is indicative of coupling between ion transport and polymer host mobility. The curved Arrhenius plots were fitted using the VTF equation (1) in which the parameter B was further employed to compute the apparent activation

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energy of the segmental polymer motion (μ σ) using the configurational entropy model (2). The curved lines in Figs. 5 and 6 correspond to non-linear least-squares fits of the data to the VTF equation using again, for the equilibrium glass transition temperature (T0), a value 25 K lower than the glass transition temperature (Tg). By examining A, B and μ values, it is evident that the primary difference between the various salts lies in the A values, they are less diminished with respect to LiTFSI when compared with PEO-based SPEs using the same salts. This suggests that the major cause of the diminished conductivity is the fact that there are fewer mobile charge carriers present in the dilithium salt SPEs, but their concentration is lower due to the cross-linking, especially for LiTFSI. With regard to the B term from the VTF fits and the apparent activation energy, μ, generally they are not diminished relative to those for LiTFSI, which may indicate that these salts did not affect the local microstructure of the host polymer. Due to the cross-linking, which makes the matrix more rigid, both LiTFSI and the new lithium salts no longer have such a strong plasticizing effect as in PEO. This becomes obvious looking at the corrected differential molar heat capacity values (Cpcorr), which measure the change in polymer molar enthalpy when the electrolyte is going from an amorphous to a glassy state at the glass transition temperature. Therefore, for PEO-based SPEs, Cpcorr increased from 11.6 0.7 J/mol K (for EO/Li  30:1) to 52.0 1.8 J/mol K (for EO/Li  10:1), an almost fivefold increase, whereas for cross-linked PEG-based SPEs it just increased negligibly from 32.4 1.6 J/mol K (for EO/Li  30:1) to 40.0 1.2 J/mol K (for EO/Li  10:1). The apparent activation energy values, μ, slowly increase with the linker chain increase for the dilithium series and are, on average, fourfold higher than those corresponding to PEO-based SPEs; this increase signifies a strong inhibition of the segmental motion of the polymer segments due to cross-linking. For the polyanionic series, μ values are in the same range for all samples, two and five times, respectively, higher than PEO-based values, which is most probably due to a much lower plasticizing effect of these salts on the polymeric host. Again, the behavior of the SPEs made using the dilithium salt with x  2 shows the distinctiveness of their properties that have been discussed earlier in the chapter. Also, for the polyanionic series, the A values diminish monotonically with an increase in the number of anionic repeating units n, which could be attributed to the slower movement of the increasing anions, due to cross-linking, and which could more easily entrap the lithium cations, increasing the ion pairing. 5. CONCLUSIONS SPEs have been prepared from a series of new bis[(perfluoroalkyl) sulphonyl]diimide dilithium salts based on LiTFSI motifs (n  1, x  2, 4, 6, 8;

Fluorinated electrolytes based on lithium salts of strong Brønsted acids

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Scheme 1) using either high-molecular-weight poly(ethylene oxide) (PEO) or cross-linked low-molecular-weight poly(ethylene glycol) (PEG) as polymeric hosts. Ionic conductivities for the SPEs were measured over a temperature range between ambient and 120°C. Conductivities of SPEs made using the dimeric salts were consistently lower than those for SPEs prepared using the monomeric salt LiTFSI over the entire temperature range, which probably reflects a diminished contribution of the anions in the dilithium salts to the overall conductivity. An unexpected finding of increasing ionic conductivity with an increase in the content of fluorine in the dianions is thought to be the result of two opposing trends: one reflecting an increase in anion size with an increased content of fluorine, which diminishes anion transport and conductivity, and another reflecting an increase in anion basicity with increased fluorination (due to the withdrawing inductive effect of fluorine atoms), which results in diminished ion pairing and an enhancement in the number of charge carriers, thereby increasing the conductivity. Although cross-linking decreased ionic conductivity, it improved the dimensional stability and the mechanical properties for all electrolytes. Still, the LiTFSI-based SPE exhibited the highest ionic conductivity for the entire temperature range. For the cross-linked SPEs based on dilithium salts, ionic conductivity exhibited the same trend as the PEO-based electrolytes: conductivity values increased over the entire temperature range with the size of the dianion. The apparent activation energy of the segmental motion (μ) and the concentration of charge carriers (A values) increased with anion size for the dilithium series for both non-cross-linked and cross-linked SPEs. This fact confirms the increase in the friction/entanglements of the polymeric segments with the increasing anion size, reducing both the segmental motion and anion contribution to the overall ionic conductivity. At the same time, the increase in ionic conductivity is due to an increase in the number of charge carriers. Also, solid polymer electrolytes have been prepared from two series of novel lithium polyanionic salts (x  4, n  1, 3, 5, 17, 225 and x  6, n  1, 3, 17) based on the LiTFSI motifs connected together by [(perfluorobutylene)disulfonyl]imide or [(perfluorohexylene)disulfonyl]imide oligomeric chains of variable length using, as polymeric hosts, high-molecular-weight PEO for the first series and cross-linked low-molecular-weight PEG for the second series. LiTFSIbased electrolytes always exhibited a higher ionic conductivity than for comparable SPEs prepared using the polymeric Li salts, which probably reflects a decrease in the anion contribution to the overall conductivity in the polymeric salts. Trends for ionic conductivity with respect to the oligomeric anion chain length (n) were noted. In particular, the existence of an optimum regarding the size of the polyanion for n ⬇ 17 was noted. This optimum value was rationalized in terms of the cumulative effects of anion mobility, ion pairing, host plasticization by the anions, and salt phase segregation on the conductivity. The use of the

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new polyanionic lithium salts resulted, on the one hand, in diminished ion pairing and an enhancement in the number of charge carriers in the SPEs, thereby increasing ionic conductivity, while on the other hand, the use of bigger polyanions produced a decrease in ionic conductivity due to the friction/entanglement with the polymeric matrix. For the cross-linked PEG-based electrolytes using the second polyanionic series (x  6, n  1, 3, 17), conductivities decreased monotonically with the number of anionic repeating units (n). This was to be expected due to the 3D network character of the cross-linked matrix for which an increase in salt anion size determined an increase in the friction/entanglement with the polymeric host. On the contrary, for the PEO-based SPEs prepared using the polyanionic series (x  4), the apparent activation energy of the segmental motion decreased with anion size (less for n ⬇ 225), whereas the A values were practically the same, suggesting that the increase in conductivity was the result of an increase in polymer matrix plasticization. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17]

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Chapter 11

Electrolytes for lithium batteries Kiyoshi Kanamura Department of Applied Chemistry, Tokyo Metropolitan University, 1-1 MinamiOhsawa, Hachioji, Tokyo 192-0397, Japan 1. INTRODUCTION Most of electrolytes for lithium batteries are aprotic organic solvents containing lithium salts. Propylene carbonate, ethylene carbonate, and diethyl carbonate are popular aprotic solvents used in lithium batteries. Dielectric constants of these solvents are relatively high, but are smaller compared to water. In order to dissolve more electrolyte salts, a pair of large and small ions is preferable. Therefore, most of lithium salts that can easily dissolve into aprotic organic solvents involve various large anions containing F elements. In this chapter, lithium salts including F elements are introduced. 2. REACTIONS IN LITHIUM BATTERIES Electrochemical reactions occurring in lithium batteries have been explained by transfer of Li ions from anode to cathode for discharge and from cathode to anode for charge. For example, in the case of rechargeable lithium ion batteries that are already commercialized, the electrochemical reactions for discharge and charge of battery are written by the following equations [1,2]: At cathode, LiCoO2 y xLi  xe  Li1xCoO2

(1)

At anode, C  xLi  xe y LixC

(2)

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The overall reaction of the battery is LiCoO2  C y Li1xCoO2  LixC

(3)

Here, LiCoO2 is used as cathode and graphite is used as anode. When the reactions proceed from left to right, the battery is charged, and when they proceed from right to left the battery is discharged. These reaction processes are schematically illustrated in Fig. 1 for rechargeable and primary batteries. In these reactions, Li ions are transferred from anode to cathode in the course of discharge process and simultaneously the electrons move from anode to cathode through external circuit. Although Li ions are not consumed in the course of discharge and charge processes, a large number of Li ions is necessary for the electrochemical reactions in lithium batteries to keep a high ionic conductivity of electrolytes, so that electrolyte salts dissociate well into ions in nonaqueous organic solvents. This leads to a high ionic conductivity of liquid electrolytes, which is very important to realize a high performance of lithium batteries. The dissociation of lithium salts depends on a kind of anion. So far, various kinds of anions have been used for lithium batteries. The typical lithium salts used in lithium batteries are summarized in Table 1 [3–6]. LiClO4 is one of well known and frequently used electrolyte salts for aprotic organic solvents. However, this electrolyte salt is not suitable for practical lithium batteries because of its high chemical reactivity. Electrolytes containing LiClO4 are more dangerous than those containing F element. Therefore, in practical (a)

Charge

Charge

(b)

Load Discharge

Discharge Discharge

Charge

Charge Discharge

Cathode

Electrolyte

Carbon

Anion in electrolyte

Oxygen

Cation in electrolyte

Anode

Co ions

Fig. 1. Schematic illustration of reaction processes for primary and rechargeable lithium batteries.

Electrolytes for lithium batteries

255

Table 1 Typical lithium salts used in primary and rechargeable lithium batteries and some of the physical properties of these salts Lithium salt

Molecular weight

Radius of anion (nm)

Conductivity in PC

LiClO4

106.39

0.237

5.6

LiBF4

93.75

0.229

3.4

LiPF6

151.90

0.254

5.8

LiAsF6

195.86

0.260

5.7

CF3SO3Li

156.01

0.270

1.7

Note: Concentration of 1 M Li salt, in mS cm1.

batteries, electrolyte salts containing F element have been used. Electrolyte salts for primary lithium batteries are different from those for rechargeable lithium batteries. In the case of primary lithium batteries, either LiBF4 or LiCF3SO3 has been used as electrolyte salt, and the case of rechargeable lithium batteries, LiPF6 has been used. This difference is due to their chemical properties and cost. 3. IONIC CONDUCTIVITY OF APROTIC SOLVENT CONTAINING ELECTROLYTE SALT Ionic conductivity of aprotic solvent containing electrolyte salt increases with increase in concentration of electrolyte salt when the concentration is not so high. At a high concentration, the ionic conductivity decreases with increase in concentration of the electrolyte salt. In other words, the ionic conductivity of nonaqueous electrolyte has a maximum value at a certain concentration of electrolyte salt. This is due to a low dissociation of electrolyte salts in aprotic organic solvents. At a high concentration, some of electrolyte salts are not present in ionic form but in a molecular form. This behavior decreases the concentration of dissociated ions in the electrolyte. Therefore, the highest solubility of electrolyte salt in organic solvent is obtained at optimized concentration, leading to the highest ionic conductivity [5]. Fig. 2 shows the dependence of ionic conductivity on the concentration of electrolyte salt (LiPF6 in a mixed solvent of ethylene carbonate and diethyl carbonate) [7,8]. The maximum ionic conductivity is obtained at 1.0 mol dm3 LiPF6. When an electrolyte salt has higher solubility than LiPF6, the maximum conductivity would be obtained at higher concentration. Therefore, the development of a new electrolyte salt is still important for practical lithium batteries.

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7

Conductivity / mS cm-1

6 5 4 3 2 1 0

0

1

2

3

4

Concentration / mol dm-3

Fig. 2. Dependence of ionic conductivity on concentration of electrolyte salts, a mixed solvent of ethylene carbonate and diethyl carbonate containing LiPF6 electrolyte salt.

4. ELECTROLYTES CONTAINING NEW SALTS WITH F ELEMENT Various kinds of electrolyte salts have been utilized in lithium batteries. Most of the electrolyte salts are inorganic compounds. Recently, organic anions have been investigated and utilized to practical batteries [9–14]. One of the representative organic electrolyte salt is LiCF3SO3 that has the simplest organic anion. This anion includes –CF3 group, so that CF3SO3 works as a strong acid. This chemical nature provides the high dissociation of the salt to anion and cation even in aprotic organic solvent, leading to a high ionic conductivity. Such organic electrolyte salts involving F element have been extensively investigated across the world. Similarly, imide compounds are very interesting as fluorinated organic electrolyte salt. A simple imide is Li(CF3SO2)2N, which has been used in practical primary lithium batteries. A negative charge of (CF3SO2)2N anion is mainly present on the N atom, but a part of the negative charge is also present on –SO2 group due to an effect of the CF3 group. However, a negative charge on the imide anion is distributed widely on the anion region. This implies that (CF3SO2)2N is a soft anion; however, Li ion is a very hard cation. These properties of cation and anion provide the high dissociation of electrolyte salts. In fact, the ionic conductivity of electrolyte with Li(CF3SO2)2N exhibits a ionic conductivity similar to that containing LiPF6 [3]. In addition, owing to the higher chemical stability, Li(CF3SO2)2N is preferable as an electrolyte in primary lithium batteries, and also for safety reasons. Various organic fluorinated electrolyte salts can be prepared by changing organic groups bonded to N atom. Table 2 shows several new electrolyte salts which have been developed and utilized for lithium batteries. In the case of Li(C2F5SO2)2N, –CF3 group can be substituted by –C2F5 group. The length of fluorinated alkyl group is changed in these new electrolyte salts, and is strongly related to an electron-withdrawing

Electrolytes for lithium batteries

257

Table 2 Fluorinated electrolyte salts for lithium batteries Li salt

Specific conductivity (mS cm1)

Molecular weight

LiCF3CO2

0.4

119.96

LiN(CF3CO)2

0.8

214.98

LiCF3SO3

2.3

156.01

LiC4F9SO3

2.3

306.03

LiC6F5SO3

1.1

254.06

LiC8F17SO3

1.9

506.06

LiN(CF3SO2)2

4.0

287.09

LiN(C2F5SO2)2

3.8

387.10

LiN(C4F9SO2) (CF3SO2)

3.5

437.11

LiN(FSO2C6F4) (CF3SO2)

3.0

385.14

LiN(C8F17SO2) (CF3SO2)

3.2

637.14

LiN(CF3CH2OSO2)2

3.0

347.14

LiN(CF3CF2CH2OSO2)2

3.0

447.16

LiN(HCF2CF2CH2OSO2)2

2.9

411.18

LiN((CF3)2CHOSO2)2

3.1

483.14

LiC(CF3SO2)3

3.6

418.16

LiTFPBa

2.7

1086.04

LiPF6

4.4

151.90

LiB[C6F3(CF3)23,5]4.

a

ability. Longer chain therefore provides larger electron-withdrawing ability. This results in lower negative charge density per unit area on anions. In the case of Li(CF3SO2)(C4F7SO2)N, fluorinated alkyl chains are not identical. This asymmetric structure also influences the distribution of negative charge. Thus, the chemical structure of fluorinated alkyl group is very important to prepare a new electrolyte salt with high dissociation ability. On the other hand, molecular weight increases with increase in the length of fluorinated alkyl group, leading to larger weight of electrolyte salt for preparing nonaqueous electrolyte. In such a case, the cost of electrolyte becomes higher. Therefore, from the cost performance point of view, more compact anion is preferable for lithium battery application.

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5. STABILITY OF FLUORINATED ELECTROLYTE SALT LiPF6 and LiCF3SO3 have been practically used in rechargeable and primary lithium batteries, respectively. Owing to its chemical stability, the latter electrolyte salt is preferable. Primary lithium batteries have been used for a long time by consumers. On the other hand, LiPF6 is the best electrolyte salt for rechargeable lithium battery at the moment, and provides a high ionic conductivity. An electrolyte salt has to be selected depending on applications. To understand the use of LiCF3SO3 in rechargeable lithium batteries, let us discuss first, the main components in lithium batteries. The main components of lithium batteries are positive electrode, negative electrode, and separator containing electrolyte, which are stacked up to produce lithium batteries. Fig. 3 shows a schematic illustration of lithium batteries. The positive sheet comprises active material, binder, conducting material, and Al foil. Generally, the active material, binder, and conducting material are mixed in organic solvent to prepare ink. This ink is painted on Al foil with a uniform thickness. After the removal of organic solvent, the sheet is pressed under appropriate pressure. This is a standard electrode preparation process for rechargeable lithium batteries. If LiCF3SO3 is used as an electrolyte salt for rechargeable lithium batteries with Al foil, the corrosion of Al takes place in the course of charge process [15–18]. Fig. 4 shows the scanning electron micrograph of Al electrode after an anodic polarization in propylene carbonate containing 1.0 mol dm3 LiCF3SO3 [18]. There are many corrosion pits on Cathode Sheet Anode Sheet

Separator

Anode composite Cu sheet Cathode composite Al sheet

Fig. 3. Schematic illustration of lithium batteries.

Electrolytes for lithium batteries

259

Fig. 4. Scanning electron micrograph of Al electrode after an anodic polarization in propylene carbonate containing 1.0 mol dm3 LiCF3SO3.

Al electrode surface, implying that the Al electrode is corroded during the anodic polarization. Al metal is originally a strong reducing reagent. It possibly reduces organic solvents or electrolyte salts. In practice, the Al metal surface is covered with native oxide film that prevents further reactions. Therefore, the stability of Al foil in nonaqueous electrolytes depends on the stability of the native oxide film. When this native film becomes unstable at a high anodic potential, dissolution of Al takes place at specific parts, resulting in corrosion with a pit formation. The surface film on Al consists of Al oxides. These oxides are dissolved in nonaqueous electrolyte, depending on the kind of electrolyte salt. Fig. 5 shows the current – potential curves of Al electrodes in propylene carbonate with various kinds of electrolyte salts [15,16]. When LiPF6, LiBF4, or LiClO4 is used as the electrolyte salt, anodic current decreases with increase in the cycle number in cyclic voltammetry study. After the second cycle, the anodic current is partially observed, indicating that the oxidation of Al electrode or electrolyte is suppressed. On the other hand, when LiCF3SO3 and imide salts are used as an electrolyte salt, a relatively large current is observed during the anodic scan. Even in cathodic scan, the anodic current is still observed; moreover, the anodic current increases with increase in cycle number. After the anodic polarization treatment, the Al surface is analyzed by some surface analysis equipment such as XPS [18]. These analyses indicate that the Al surface after the anodic polarization is still covered with Al oxides or Al oxyfluorides when using LiPF6 or LiBF4. Fig. 6 shows the schematic illustration of surface film on Al electrode before and after anodic polarization. The change in the surface state may improve the surface stability of Al electrode. However, in the case of fluorinated imide salts or LiCF3SO3, the surface film on Al electrode is partly destroyed by the

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18

1

LiCF3SO3 2nd run

0.8

Current density / mA cm-2

Current density / mA cm-2

16

LiClO4 0.6

0.4

LiBF4 0.2

LiPF6 0

3

4

5

6

14 LiCF3SO3 1st run

12 10 8 6 4 2 0

7

LiPF6 LiPF6 1st run 2nd run

2

4

6

8

E / V vs. Li / Li+

E / V vs. Li/Li+

Fig. 5. Current – potential curves of Al electrodes in propylene carbonate with various kinds of electrolyte salts.

Al oxides or oxyfluorides

Al oxides LiPF6 Al

Oxidation

Ox id

LiCF3SO3

Al Al3+

Al3+

ati on

Al

Fig. 6. Schematic illustrations of surface film on Al electrode before and after anodic polarization, and corrosion of Al electrode.

anodic polarization. As a result, Al electrode surface is in direct contact with the electrolyte, leading to an active dissolution of Al with reduction of nonaqueous electrolyte. In Fig. 6, the schematic representation of the corrosion model of Al electrode is also shown. For the above reasons, the Al current collector cannot be used in rechargeable lithium batteries using fluorinated imide salts or LiCF3SO3. On the other hand, primary lithium batteries have 3 V operational voltage and are never been charged, therefore imides and LiCF3SO3 can be used. A primary battery with MnO2 as the cathode material and Li as the metal anode is

Electrolytes for lithium batteries

261

constructed, using LiCF3SO3 as the electrolyte salt due to its high stability. This battery is sometimes used for more than 10 years. Therefore, the stability of the electrolyte salt is a more important factor. Fig. 7 shows the thermal stability of

LiBF4 →LiF+BF3

TG

82.1 °C DTA

840.2 °C 310.3 °C

0

200

400

600

800

1000

Temperature/°C

(a) TG

LiPF6 →LiF+PF5 DTA

187.7 °C

226.8 °C 840.5 °C

0

200

400

600

800

1000

Temperature/°C

(b)

LiCF3SO3 TG

DTA 82.1 °C 165.3 °C

418.7 °C

0 (c)

100

200

300

Temperature/°C

Fig. 7. Thermal stability of several electrolyte salts.

400

500

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Kiyoshi Kanamura

several electrolyte salts [9,14]. LiPF6 and LiBF4 decomposed at 200–300°C. Most of imides and LiCF3SO3 have a higher decomposition temperature (400°C). This thermal stability is very important for the safety of lithium batteries, especially, it is more critical for rechargeable lithium batteries. From the viewpoint of stability, imide salts are preferable, but they cannot be used in rechargeable lithium batteries due to the corrosion of Al current collector. This is a high motivation for development of new electrolyte salts. 6. ROLE OF ELECTROLYTE SALTS The primary role of lithium salts is to provide a high ionic conductivity to nonaqueous organic solvent. Another important role is an effect on surface state of active materials. Carbon materials have been used in the case of rechargeable lithium ion batteries. In primary batteries, Li metal has been used as anode material. Both the anodes have a strong reducing ability and hence electrolyte solutions are easily reduced. From thermodynamic considerations, batteries cannot be constructed using these anodes. However, active materials are often covered with surface films which prevent a direct contact between the electrode and the electrolyte. This is very important in lithium battery system. A number of papers have been published related to surface chemistry of Li metal and carbon materials by using various kinds of surface analysis tools [19–28]. In the batteries, aprotic organic solvents mainly react with Li metal or lithiated carbon. In addition, electrolyte salt plays an important role in the formation of surface film on Li metal and carbon materials. Since these surface films have an ionic conductivity, they are called solid electrolyte interface (SEI) [29]. Fig. 8(a) shows the SEI on Li metal and carbon when LiClO4 is used as an electrolyte salt [19]. These surface films consist of various alkyl carbonates, alkoxides, and so on. In these cases, fluoride compounds are not detected. On the other hand, when fluorinated electrolyte salts such as LiPF6 and LiBF4 are used, LiF is formed on surfaces of Li metal and carbon [19]. This result indicates that anions also react on anodes. Another possible explanation for the formation of LiF is based on the stability of electrolyte salts. Most of electrolyte salts involve impurities that are sometimes produced by decomposition of electrolyte salts, for example by hydrolysis of anions. In fact, LiPF6 and LiBF4 decompose even under low humidity as a result of chemical reactions with water to produce HF and LiPOFn or LiBOFm molecules, according to the following equations [30]: LiPF6  H2O → POF3  LiF  2HF

(4)

LiBF4  H2O → BOF  LiF  2HF

(5)

These equations represent one of the decomposition processes of fluorinated electrolyte salts; there are other possible reactions for hydrolysis and decomposition

Electrolytes for lithium batteries

263

Li2CO3, Alkoxides, Alkyl carbonate

LiOH LiF Li

(a)

Li2O

Li

LiClO4

Li2CO3, Alkoxides, Alkyl carbonate

Carbon

LiF, Alkoxides, Alkyl carbonate

Carbon

(b)

Fig. 8. SEI on Li metal and carbon in ethylene carbonate base solvent containing LiClO4 and LiPF6 electrolyte slats.

processes. In either case, HF is produced by these chemical reactions. It is an acidic compound and easily reacts with alkyl carbonate, alkoxides, and other basic inorganic and organic compounds such as LiOH and Li2O. When these basic compounds are present on the surface of anode materials, HF reacts with these compounds to form LiF. As a result of these reactions, the compounds on surface films on anode materials are converted into compounds containing a large amount of LiF. Fig. 9(a) shows the scanning electron micrograph of Li metal surface deposited on Ni metal substrate in aprotic organic solvent containing LiPF6 [20]. Fig. 9(b) shows the scanning electron micrograph of Li metal deposited on Ni metal substrate in aprotic organic solvent containing LiClO4 [20]. The former Li metal exhibited a flat and hemispherical morphology and the latter showed a typical dendrite form. This difference in surface morphology is due to the presence of a kind of surface film. The surface film providing hemispherical shape of Li metal consists of LiF thin layer as shown in Fig. 8(a). This surface film is formed by the following reactions of native surface film consisting of LiOH, Li2O, and Li2CO3: LiOH  HF → LiF  H2O

(6)

Li2O  2HF → 2LiF  H2O

(7)

Li2CO3  2HF → 2LiF  H2CO3

(8)

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Fig. 9. Scanning electron micrographs of Li metal surface deposited on Ni metal substrate in aprotic organic solvent containing (a) LiPF6 and (b) LiClO4. O

O

O

O Li+

B O

O

O

O

Fig. 10. Chemical structure of LiBOB.

In these reactions, HF, which is produced by the decomposition of fluoride anions, plays an important role. Thus, electrolyte salts containing F element influence the surface state of Li metal. 7. DEVELOPMENT OF NEW ELECTROLYTE SALTS Many kinds of electrolyte salts with or without F element have been developed; the most recent and well known being LiBOB [31]. This electrolyte salt provides a good surface film on carbon materials. The chemical structure of the anion in LiBOB is shown in Fig.10. It can be seen from the figure that two oxalate groups are bonded to B element. This anion decreases the irreversible capacity of carbon materials [32]. At this moment, F element is not involved in this electrolyte salt. However, substitution of oxalate anion by other chemical groups having low molecular weights would be needed to decrease the total molecular weight of the anion; otherwise, the anion cannot be utilized in practical lithium batteries. Therefore, the fluorine chemistry on this anion is very important. 8. SOLVENT FOR ELECTROLYTE In practical batteries, most of the solvents are hydrocarbons such as ethylene carbonate, diethyl carbonate, propylene carbonate, and so on. These organic

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solvents can be easily oxidized or reduced at anode or cathode. As a result of such electrochemical reactions, surface films are formed on the electrode materials that work as protective film for oxidation or reduction of electrolytes. These protective films are very important for high operation voltage of lithium batteries. If no protective films are present on electrode surfaces, batteries cannot be constructed using carbon and Li metal anode. So far, much work has been done to improve chemical stability of organic solvents. One of the research vectors is a new synthesis of fluorinated organic solvents. Organic compounds involving F element is usually very stable. This topic is discussed in a later chapter. 9. SUMMARY Several kinds of electrolyte salts containing F element have been utilized in practical primary and rechargeable lithium batteries. These electrolyte salts provide high ionic conductivities of organic electrolytes. In addition, the surface state of electrodes is also influenced by the presence of F in the electrolytes. Demand for primary and rechargeable lithium batteries is increasing day by day. In order to satisfy the demand, new electrolyte systems and additives to electrolytes have to be developed. In further work on electrolytes, F element may be a key compound to obtain a high performance of lithium batteries. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18]

T. Nagaura and K. Tozawa, Prog. Batt. Solar Cells, 9 (1990) 209. G. Pistoia (Ed.), Lithium Batteries, New Materials, Developments and Perspectives, Elsevier, Amsterdam, 1994. M. Ue, J. Electrochem. Soc., 141 (1994) 3336. M. Ue and S. Mori, J. Electrochem. Soc., 142 (1995) 2577. Y. Matsuda, M. Morita, and F. Tachihara, Bull. Chem. Soc. Jpn., 59 (1986) 1967. S. Tobishima and T. Okada, Electrochim. Acta, 30 (1985) 1715. K. Kondo, M. Sano, A. Hiwara, T. Omi, M. Fujita, A. Kuwae, M. Iida, K. Mogi, and H. Yokoyama, J. Phys. Chem. B, 104 (2000) 5040–5044. A.M. Christie and C.A. Vincent, J. Phys. Chem., 100 (1996) 4618–4621. K. Momota, Batt. Technol., 8 (1996) 108. F. Kita, A. Kawakami, T. Sonoda, and H. Kobayashi, J. Power Sources, 68 (1997) 307. F. Kita, A. Kawakami, T. Sonoda, and H. Kobayashi, Batt. Technol., 6 (1994) 45. F. Kita, A. Kawakami, T. Sonoda, and D. Kagaku, 65 (1997) 909. T. Sonoda, J. Nie, H. Kobayashi, F. Kita, and A. Kawakami, Batt. Technol., 10 (1998) 106. Y. Sasaki and N. Nanbu, Materials Chemistry in Lithium Batteries, N. Kumagai, S. Komaba, and M. Wakihara (Eds.), Research Signpost, Trivandrum, India, 2002, p. 415. K. Kanamura, T. Okagawa, and Z. Takehara, J. Power Sources, 57 (1995) 119. K. Kanamura, Batt. Technol., 10 (1998) 85. L.J. Krause, W. Lamanna, J. Summerfield, M. Engle, G. Korba, R. Loch, and R. Atanasoski, J. Power Sources, 68 (1997) 320. K. Kanamura, T. Umegaki, S. Shiraishi, M. Ohashi, and Z. Takehara, J. Electrochem. Soc., 149 (2002) A185.

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[19] K. Kanamura, H. Tamura, and Z. Takehara, J. Electroanal Chem., 333 (1992) 127. [20] K. Kanamura, H. Tamura, S. Shiraishi, and Z. Takehara, J. Electroanal Chem., 394 (1995) 49. [21] K. Kanamura, H. Tamura, S. Shiraishi, and Z. Takehara, J. Electrochem. Soc., 142 (1995) 340. [22] K. Kanamura, H. Tamura, S. Shiraishi, and Z. Takehara, Electrochim. Acta, 40 (1995) 913. [23] K. Kanamura, S. Shiraishi, and Z. Takehara, Chem. Lett., 1995 (1995) 209. [24] K. Kanamura, S. Shiraishi, and Z. Takehara, J. Electrochem. Soc., 143 (1996) 2187. [25] D. Aurbach, Y. Ein-Eli, O. Chusid, Y. Carmeli, M. Babai, and H. Yamin, J. Electrochem. Soc., 141 (1994) 603. [26] D. Aurbach, B. Markovsky, A. Shecheter, Y. Ein-Eli, and H. Cohen, J. Electrochem. Soc., 143 (1996) 3809. [27] M. Winter, G.H. Wronigg, J.O. Besenhard, W. Biberacher, and P. Novak, J. Electrochem. Soc., 147 (2000) 2427. [28] M. Inaba, Y. Kawatate, A. Funabiki, S.K. Jeong, T. Abe, and Z. Ogumi, Electrochim. Acta, 45 (1999) 99. [29] E. Peled, J. Electrochem. Soc., 126 (1979) 2047. [30] D.W. Sharp, Advances in Fluorine Chemistry, Vol. 1, M. Stacey, J.C. Taltow, and A.G. Sharpe (Eds.), Butterworths Scientific Publication, London, 1960, p. 69. [31] W. Xu and C.A. Angell, Electrochem. Solid-State Lett., 4 (2001) E1. [32] K. Xu, S. Zhang, T.R. Jow, W. Xu, and C.A. Angell, Electrochem. Solid-State Lett., 5 (2002) A 26.

Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 12

Thermally stable fluoro-organic solvents for lithium ion battery Jun-ichi Yamaki Institute for Materials Chemistry and Engineering, Kyushu University 6-l Kasuga Koen, Kasuga 816-8580, Japan Tel: 8l-92-583-7790 E-mail:[email protected] 1. INTRODUCTION Recently, the study of the utilization of the high-performance Li-ion cells as power sources of electric vehicles (EV) and other large-sized equipment has been undertaken by many corporations and laboratories. However, large-sized lithium ion batteries are not yet commercially available, primarily because the effect of the scale-up of this kind of battery on its safety has not yet been fully estimated. Several exothermic reactions (such as decomposition of electrolyte itself and decomposition of electrolyte with lithiated carbon anode and charged Li0.5CoO2 cathode) occur inside a cell as its temperature increases. A “thermal runaway” is considered to have occurred if heat output exceeds thermal diffusion. This has prompted many researchers to carry out thermal stability studies on these batteries [l–l5]. The nonaqueous electrolytes used in lithium-ion cells consist of a lithium salt and a flammable organic solvent. The latter is considered to be one of the reasons for the failure of these batteries in terms of safety. Organic compounds containing fluorine species are nonflammable and have unique properties. Therefore, many kinds of fluorinated organic solvents have been studied as potential cosolvents of electrolytes in order to improve the flammability and low-temperature performance of graphite anode and lithium-ion cells [16–18]. The thermal stability of fluorinated esters in fluorinated organic solvents was investigated by Yamaki et al. [19–21]. Then it was found that 1 M LiPF6/methyl difluoroacetate (MFA) exhibits better stability in coexistence with a lithium metal anode than does the conventional electrolyte. For example, LiPF6/MFA, as compared with LiPF6/(EC) ethylene carbonate-dimethyl carbonate (DMC) (1:l, by vol), raised the onset temperature of the exothermal reaction in coexistence with lithium metal, Li0.5CoO2 or lithiated carbon anodes.

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2. THERMAL STABILITY OF FLUORINATED ESTERS Yamaki et al. [191 investigated the thermal stability of fluorinated esters, which are the same fluorinated esters used as additives to improve the cycling performance reported by Nakajima et al. [17,18] prior to our study. In this study, partially fluorinated carboxylic acid esters (Table 1) were used as the electrolyte solvent and LiPF6 as the salt. LiPF6 salt was dissolved in esters 1 MFA, and 2 ethyl difluoroacetate EFA) to a salt concentration of 1 M. In the other fluorinated esters, however, LiPF6 salt could not be dissolved to a salt concentration of 0.2 M. Therefore, the solutions of fluorinated esters (1 and 2) with 0.2 M of LiPF6 were used, and the other fluorinated esters were saturated with LiPF6. For comparison, the solutions of corresponding esters with 0.2 M LiPF6 were prepared. Similar measurements were performed employing the conventional electrolyte solution used in lithium batteries: 1 M LiPF6 /EC  DMC (1:1, by vol). The thermal stability of fluorinated esters (Table 2) was examined using a TG-DSC. Each sample (5 μl for liquid) for TG-DSC measurement was packed in a stainless-steel case, which was then crimp-sealed in a glove box filled with argon. In some cases, a piece (weighing several milligrams) of lithium metal or a charged LiCoO2 pellet was packed and sealed along with a sample in the stainless-steel case. No leak of the case was confirmed by TG data, which were simultaneously measured with DSC measurement. A LiCoO2 pellet was prepared by mixing LiCoO2, acetylene black, and a polytetrafluoroethylene binder. Table 1 Esters used as solvents [19] Nonfluorinated solvent Sample no.

Fluorinated solvent

Solvent

Sample no.

Solvent

1

CH3COOCH3 (MA)

1

CHF2COOCH3 (MFA)

2

CH3COOCH2CH3 (EA)

2

CHF2COOCH2CH3 (EFA)

3

CH3CH2COOCH3

3

CF3CF2COOCH3

4

CH3CH2COO CH2CH3

4

CF3CF2COO CH2CH3

5

H(CH3)2CCOOCH3

5

F(CF3)2CCOOCH3

6

H(CH2)3COOCH3

6

F(CF2)3COOCH3

7

H(CH2)3COOCH2CH3

7

F(CF2)3COOCH2CH3

8

H(CH2)4COOCH2CH3

8

H(CF2)4COOCH2CH3

9

H(CH2)7COOCH2CH3

9

F(CF2)7COOCH2CH3

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269

Table 2 Initial peak temperatures in DSC curves [19] Nonfluorinated solvent

Fluorinated solvent

Initial peak temperature (°C)

Initial peak temperature (°C)

Electrolyte

Electrolyte  Li

Electrolyte  Li0.5CoO2

Electrolyte Electrolyte  Electrolyte  Li Li0.5CoO2

1

280

220

230

1

290

290

310

2

210

110

240

2

210

180

210

3

260

90

200

3

280

330

210

4

210

90

200

4

250

180

240

5

250

70

250

5

270

290

180

6

260

90

170

6

260

300

230

7

210

90

180

7

260

170

220

8

240

120

170

8

250

300

230

9

250

120

170

9

230

160

220

The mixture was packed in a coin cell with lithium metal anode and 1 M LiPF6 /EC  DMC electrolyte, and then charged to Li0.5CoO2 in a constantcurrent mode. 2.1. Thermal stability of LiPF6 /fluorinated ester electrolyte

The thermal stability of the LiPF6 /fluorinated esters was similar to those of the LiPF6 /corresponding esters, except for esters 3, 4, 5, and 7 (Table 2). It should be noted that the fluorinated esters contained smaller amounts of LiPF6 than the corresponding esters, except for 1 and 2. In LiPF6 electrolytes, ionic dissociation of LiPF6 is not high, and LiPF6 is in equilibrium with LiF and PF5. PF5 is a strong Lewis acid, which reacts with a small amount of water in electrolytes following the reaction PF5  H2O → POF3  2HF [22]. Based on the analogy with this reaction, organic solvents may have reacted with PF5 at a high temperature. The thermal decomposition of LiPF6 electrolytes is probably caused by PF5 in the electrolytes. It has been reported that the direct reaction of PF5 with EC/EMC is similar to the thermal decomposition of LiPF6 in EC/EMC [23]. PF5 may attack the carbonyl oxygen of nonfluorinated solvents. The reaction mechanism and the stabilities of PF5-solvent complexes are not changed significantly by the fluorination of the esters, since the thermal stabilities of LiPF6 /fluorinated esters were similar to those of the LiPF6 /corresponding esters.

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The DSC curves of electrolyte (3 μ1) with 1 M LiPF6 and 1 M lithium-imide salts are shown in Fig. 1. The electrolytes with the lithium-imide salt were not stable compared with LiPF6 /MFA. However, the Li-imide electrolytes showed better stability than LiPF6 /EC:DMC (l:1). There was no exothermic peak at 300°C for LiPF6/MFA, although solid LiPF6 decomposes to LiF and PF5 at around 300°C. 2.2. Thermal stability of LiPF6 /fluorinated ester electrolyte with Li metal anode.

Li metal was used as a model compound of lithiated carbon anodes because its reactivity is similar to that of lithiated carbon anodes. Many of the fluorinated ester systems coexisting with Li metal provide exothermic peaks at temperatures higher than the melting point of lithium metal, whereas the corresponding ester systems generate an exothermic reduction below the melting point of lithium metal (Table 2). This indicates the reduction of LiPF6 /MA solution by lithium metal. However, there is no other exothermic reaction of the LiPF6 /MFA/lithium metal system at this temperature. Nakajima et al. [17] reported that “the fluoroesters are more easily reduced than EC/DEC by electrochemical reduction”. This means that the fluoroester electrolytes are more reactive with Li metal than EC/DEC electrolytes. As a rule, the reduction stability of partly fluorinated solvents is low because of the strong electron-withdrawing effect of the fluorine atoms. Therefore, it is necessary to consider the contribution of fluoroesters to the solid electrolyte interphase (SEI) [24]. The details on the SEI are introduced in the next section. It is also necessary to consider the reactivity of the solvent itself with lithium metal. Many esters can also exist in enol form in addition to the keto form. The enol form reacts quickly with lithium metal [25]. The enol form of MFA may be unstable compared with that of MF because of the strong electron-withdrawing effect of fluorine atoms. However, the content of enol MFA is very small. This is another possible reason for the low reaction temperature of LiPF6 /MA/lithium metal system.

Fig. 1. DSC curves of electrolyte (3 μl) with 1 M LiPF6 and 1 M lithium-imide salts.

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The cycling efficiency of lithium metal electrode was estimated by a cycle test using a coin cell. The cycling efftciency of the lithium metal electrode with 1 M LiPF6 /MA, MFA, EA, or EFA is 30, 84, 0, or 50%, respectively. The DSC curves of electrolyte (3 μ1) with 0.56 mg of Li metal are shown in Fig. 2 [26]. The Li-imide salt electrolytes were stable; however, their SEI was found to be slightly different from that of LiPF6 /MFA by FT–IR measurement. A main component of SEI formed between LiPF6 /MFA and Li metal is CHF2COOLi [21], which is a reaction product of Li and MFA. However, there was another component formed by a reaction with Li salts. 2.3. SEI on lithium metal anode [21]

SEI covers the surface of lithium metal and prevents further reduction of electrolytes. In the case of heating the LiPF6/MA/lithium metal system, the SEI layer may be effective up to 200°C and the reduction rate may be so accelerated that the SEI layer cannot protect the reduction at higher temperatures. The effectiveness of the SEI layer can be estimated by the position of the exothermic peak; that is, the higher temperature of the exothermic peak may indicate that, in this case, the SEI layer is thick enough to prevent any substantial reduction. When the LiPF6/MFA/lithium metal system was heated, no exothermic peak could be observed below 250°C and a broad peak appeared around 290°C. This indicates that the LiPF6 /MFA/lithium metal system may provide a thicker and more effective SEI layer on the lithium surface than the corresponding ester system. CHF2COOCH3 (MFA, Daikin Fine Chemical Co. Ltd) was used after dehydration by a molecular sieve 3A l/8. Lithium hexafluorophosphate (LiPF6) was dissolved in MFA to 1 M. The water content of the solution as estimated by a Karl Fischer aquameter (Aquacounter AQ-7; Hiranuma Co.) was 20 ppm or less.

Fig. 2. DSC curves of 1 M LiPF6 and 1 M lithium-imide salt electrolytes (3 μl) with 0.56 mg of Li metal.

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The solution of 1 M LiPF6 /EC-DMC (1:1, by vol) (Tomiyama Chemical Co.) as received was used for comparison. The lithium surface for the FT–IR and XPS analyses was prepared as follows: a coin-type cell was fabricated using a working electrode of Li metal foil and a counterelectrode of stainless steel. Lithium was removed from the Li surface at a constant current of 1.0 mA. Subsequently, the cell was stored for 1 day at 25°C in order to grow an SEI on the surface of the Li metal electrode. The coin cell was disassembled in a glove box filled with argon, and the deposited lithium was washed with DMC or MFA solvent. When FT–IR was performed, the lithium deposited on the stainless steel was mixed with KBr powder. The mixture was pressed into a tablet outside the box. For XPS measurement, lithium metal electrode was put into the sample holder in situ after vacuum drying. For the measurements using both apparatuses, commercial CHF2COOLi (SynQuest Laboratories, Inc.) was used as a reference. The MFA solvent can be more easily reduced than EC-DEC as reported by Nakajima et al. [17]. Generally, a partly fluorinated organic compound is rather unstable against reduction because of the strong electron-withdrawing effect of the fluorine atoms. An electrolyte solution containing such a fluorinated solvent is considered to form an effective SEI on an anode surface. The extensive work by Aurbach et al. in analyzing the basic chemistry of the SEI layer that is formed on lithium metal and carbon surfaces revealed the reduction mechanisms of the electrolyte on the anode surface as shown below [22]. According to them, lithium metal and lithiated carbon provide SEI on their surfaces with a similar composition, and for the most commonly used alkyl carbonate solvents, EC and DMC, the main reduction products generally accepted are lithium alkyl carbonates (R-OCO2Li), which are produced through the following reactions [22]: 2(CH2O)2CO (EC)  2e  2Li → (CH2OCO2Li)2 ↓ C2H4 ↑

(1)

CH3OCO2CH3 (DMC)  e  Li → CH3OCO2Li ↓  CH3.

(2)

Therefore, when MFA-based electrolyte is used, the expected main reduction product of lithium is lithium difluoroacetate (CHF2COOLi), which is formed through the following reaction: CHF2COOCH3  e  Li → CHF2COOLi ↓  CH3.

(3)

CHF2COOLi was used as a reference for XPS and FT–IR spectroscopy. Fig. 3 shows the FT–IR spectra obtained from the reference CHF2COOLi, lithium electrodes in EC–DMC and MFA-based electrolyte. When MFA-based electrolyte is used, the FT–IR spectra are different from those of EC–DMC-based electrolyte. From these results, it is clear that SEI in MFA-based electrolyte has a different

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composition than that in EC–DMC-based electrolyte. Moreover, the FT–IR spectra obtained from the reference CHF2COOLi and lithium electrodes in MFAbased electrolyte exhibited two pronounced peaks around 1500 cm1. These peaks may be attributed to the presence of CO(–COOLi) bonding. However, from the FT–IR measurements of the SEI formed at the MFA-based electrolyte/lithium interphase, sufficient information was not obtained to determine that the main component of this SEI is CHF2COOLi. Fig. 4 illustrates the energy levels of Lils, Cls, Ols, and Fls in the XPS spectra of SEI formed between lithium metal and various electrolytes. The Li peak corresponding to CHF2COOLi was obtained around 58 eV, and the energy levels of Lils show that the SEI formed on Li metal in MFA-based electrolyte has a

Absorbance [Arb. Unit]

(a)

3000

(b)

(c)

2500

2000

1500

1000

500

Wavenumber [cm−1]

Fig. 3. FT–IR spectra of various deposited lithium: (a) CHF2COOLi as reference, (b) deposited in 1 M LiPF6/MFA, and (c) deposited in 1 M LiPF6/EC-DMC (1:1, by vol).

Fig. 4. XPS spectra of various lithium electrodes: (a) CHF2COOLi as a reference, (b) in 1 M LiPF6/MFA, and (c) in 1 M LiPF6/EC-DMC (1:1, by vol). The spectra were obtained for the energy levels of Lils, Cls, Ols, and Fls for all systems.

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chemical composition similar to that of CHF2COOLi in terms of Li species. The Cls spectrum for CHF2COOLi shows a broad peak at 282.5–290 eV. It seems that the peak position is similar to that in the SEI formed on Li metal in MFA-based electrolyte in terms of the C species. The O1s spectrum for CHF2COOLi also contains a broad peak at 530–539.5 eV, with the major peak obtained at 536.1 eV, and this spectrum is similar to the SEI layer of the MFA system. Furthermore, a strong peak shown in the Fls spectrum in (a) at 689.2 eV is similar to that shown for SEI in the spectrum in (b). From these results, it is clear that the main component of SEI formed between the MFA-based electrolyte/anode surfaces is CHF2COOLi, which is expected to enhance the thermal stability of the electrolyte – anode system. In contrast, all the peaks for the four materials in the SEI of the EC–DMC system (c) differ from those in the MFA system (b). It was clear that the SEI of the composition in the MFA system differs from that in EC–DMC system. The thermal stability of CHF2COOLi was also assessed using DSC measurement. Fig. 5 shows the DSC profile of CHF2COOLi. This profile shows an endothermic peak at 150°C and a broad exothermic peak around 220°C. The endothermic peak at 150°C was not observed after the CHF2COOLi was preheated up to 180°C. CHF2COOLi remained in a similar white powder form after heating to 180°C. The chemical composition of CHF2COOLi before and after the heat-treatment at 180°C and 300°C was estimated using FT–IR. Fig. 6 shows the FT–IR spectra of CHF2COOLi before and after the heat treatment. The correspondence of the peaks in spectra (a) and (b) clearly indicates that the composition of CHF2COOLi did not change up to 180°C. The endothermic reaction at 150°C is considered to be water evaporation. In contrast, spectrum (c) is different from spectra (a) and (b). The thermal decomposition of CHF2COOLi should occur

Fig. 5. DSC profiles of CHF2COOLi as a reference: (a) as received, (b) heated up to 180°C and cooled up to 50°C and (c) heated up to 180°C and cooled up to 50°C after it has been preheated up to 180°C.

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275

up to 300°C, preferably around 220°C as shown in the DSC profile. It was found that the thermal stability of CHF2COOLi itself is not sufficient for this material to contribute to the thermal stability of the MFA-based electrolyte, because CHF2COOLi in the SEI does not exist stably at 220°C. It is expected that the decomposition product or the reaction product of CHF2COOLi with lithium under high temperatures contributes to the improved thermal stability of its SEI. Fig. 7 shows the DSC profiles of CHF2COOLi with or without lithium metal. This result indicates that no additional exothermic peaks appeared with the inclusion of lithium metal. The heat generation at 200°C markedly increased

Fig. 6. FT–IR spectra of CHF2COOLi as a reference: (a) initial, (b) heated up to 180°C and (c) heated up to 300°C.

Fig. 7. The DSC profiles of CHF2COOLi coexisted with and without lithium metal (a) the profile of CHF2COOLi and (b) the profile of CHF2COOLi coexisting with lithium metal.

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from 598 to 1681 J/g in the presence of lithium metal while the endothermic heat at around 120°C did not change. The increase in the exothermic heat at 200–300°C indicated that CHF2COOLi or its decomposition product reacted with lithium metal in this temperature range. As already mentioned, MFA-based electrolyte clearly provides the thermally stable SEI onto the surface of lithium or lithiated carbon. However, CHF2OOLi, which is expected to be the main component of SEI provided by MFA-based electrolyte, is not thermally stable enough to account for its thermal stability when coexisting with lithium or lithiated carbon. The component of SEI in the high-temperature region on the lithium in MFA-based electrolyte is a decomposition product of CHF2COOLi or a decomposition product formed by the reaction with lithium metal at approximately 200–300°C. The peak indicated that the decomposition of CHF2COOLi did not appear in the DSC profiles of MFA-based-electrolyte/lithium or lithiated carbon surface. This material was to be decomposed; however, because the SEI layer was considered to be very thin, it was so small in amount that the heat generation was not observable. Even though there is no information about the actual “effective” material in the SEI layer in the high-temperature region, it is at least clear that this material also prevents the reaction of MFA-based electrolyte with lithium metal or lithiated carbon. The most likely candidate for the product is LiF. 2.4. Thermal stability of LiPF6 /fluorinated ester electrolyte with charged Li0.5CoO2 cathode

The positions of the exothermic peaks of the LiPF6/various fluorinated esters and corresponding esters/Li0.5CoO2 systems are also summarized in Table 2. With the exception of esters 2 and 5, all fluorinated esters tended to inhibit the reaction with Li0.5CoO2. The cycling performance of Li/LiCoO2 cells with 1 M LiPF6/MA, MFA, EA, or EFA was investigated. The MFA and EFA electrolytes showed very good performances. However, the cell with the MA or EA electrolyte could not be cycled because of the low oxidation potentials of the solvents. It is well known that charged LixCoO2 (x l) is metastable, and that oxygen evolution was observed at temperatures above 200°C [1]. LixCoO2 , delithiated by a chemical method using H2SO4 [27,28], was investigated by means of DSC with or without an electrolyte (1 M LiPF6/EC  DMC) [13]. The lithium content x in the delithiated LixCoO2 was determined by atomic absorption spectroscopy. The DSC measurements of Li0.49CoO2 with the electrolyte at various mixing ratios showed two exothermic peaks, one beginning at 190°C and the other at 230°C. The exothermic heat of each peak was proportional to the amount of Li0.49CoO2. The peak starting at 190°C probably resulted from the decomposition of solvent due to an active cathode surface, and the peak starting at 230°C was electrolyte oxidation caused by released oxygen from Li0.49CoO2. However, no exothermic peak caused

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Fig. 8. DSC curves of 1 M LiPF6 and 1 M lithium-imide salt electrolytes (3 μl) with 3 mg of Li 0.48CoO2.

by cathode surface reaction was observed when 1 M LiPF6 /MFA electrolyte was used (Fig. 8) because of the difficulty of the oxidation by fluorination. 3. THERMAL STABILITY OF 1 M LiPF6/EC/DMC  FLUORINATED ESTER ELECTROLYTE [20] Unfortunately, many fluorinated carboxylic acid esters did not dissolve in 0.2 M LiPF6, and therefore saturated solutions were used for the experiments. However, ca. 1 M LiPF6//EC/DMC/fluorinated carboxylic acid esters (1:1:2, by vol) can be prepared for all the fluorinated carboxylic acid esters listed in Table 1. Using the mixed solvent electrolyte, the thermal stability of the electrolyte system coexisting with Li metal was investigated [20]. The results are shown in Table 3. The onset temperature increased by the addition of CHF2COOCH3. The addition of CHF2COOCH3, CF3CF2COOCH2CH3, or F(CF2)3COOCH3 decreased the exothermic energy. These results indicate that CHF2COOCH3 was the most effective additive in this study. A precise thermal study was undertaken [20] for 1 M LiPF6/CHF2COOCH3 and 1 M LiPF6 /EC  DMC or propylene carborate (PC) (ca. 1 M LiPF6/ CHF2COOCH3 /EC/DMC [CHF2COOCH3 : EC : DMC  x : (100  x)/2 : (100  x)/2] and ca. 1 M LiPF6//CHF2COOCH3 /PC [CHF2COOCH3 : PC  x : (100  x)] changing the mixing ratio with the coexistence of lithium metal (Figs. 9 and 10). The remaining lithium metal content was estimated using DSC by the endothermic vs. exothermic heat ratio at lithium melting and freezing from 300°C. The electrolyte volume for DSC is 5 μl, the lithium weight is 1.3 mg, the heating rate is 5°C/min., and the temperature range is from room temperature to 300°C. As the

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Table 3 Onset temperatures of ca. 1 M LiPF6//EC/DMC/fluorinated carboxylic acid esters (1:1:2, by vol) with the coexistence of lithium metal, and the exothermic energy of electrolytes from 180 to 220°C [20] Fluorinated solvent

Onset temperature (°C)

Exothermic energy (J/g)

No addition (1 M EC  DMC)

180

2900

CHF2COOCH3

210

700

CHF2COOCH2CH3

175

 3500

CF3CF2COOCH3

168

2100

CF3CF2COOCH2CH3

172

 3500

F(CF3)2CCOOCH3

173

3400

F(CF2)3COOCH3

167

1900

F(CF2)3COOCH2CH3

169

3000

F(CF2)4COOCH2CH3

168

3000

Fig. 9. DSC profiles of the mixture of 1 M LiPF6/EC  DMC and 1 M LiPF6/MFA with Li metal. 1 M LiPF6/EC  DMC : 1 M LiPF6/MFA (by vol) are (a)1:0, (b)4:1, (c)l:1, (d)3:7 and (e)0:1.

volume ratio of CHF2COOCH3 increased from 0 to 100%, the amount of lithium remaining increased from 0 to 95% (Figs. 11 and 12). Using FTIR and XPS, the main component of SEI in the MFA electrolyte was found to be CH3COOLi [21]. From these results, it is clear that CHF2COOLi

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279

Fig. 10. DSC profiles of the mixture of 1 M LiPF6/PC and 1 M LiPF6/MFA with Li metal. 1 M LiPF6/PC : 1 M LiPF6/MFA (by vol) are (a)1:0, (b)4:1, (c)1:1, (d)3:7 and (e)0:1.

Fig. 11. Dependence of MFA content on thermal stability of ca. 1 M LiPF6/EC  DMC  MFA with lithium metal. (䊊) the total exothermic energy from 140 to 300°C; (䊉) the amount of lithium metal remaining after heating.

is expected to enhance the thermal stability of the electrolyte – anode system. However, CH3COOLi does not exist in the SEI of the EFA electrolyte. The MFA electrolyte appears to react with Li metal to form SEI through the following reaction [21]: CHF2COOCH3  e  Li → CHF2COOLi  CH3*

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Fig. 12. Dependence of MFA content on the thermal stability of ca. 1 M LiPF6/PC  MFA with lithium metal; (䊊) the total exothermic energy from 140 to 300°C; (䊉): the amount of lithium metal remaining after heating.

It is also probable that the EFA electrolyte reacts with Li metal to form CHF2COOLi through the following reaction: CHF2COOC2H5  e  Li → CHF2COOLi  C2H5* The question is why CH3COOLi does not exist in the SEI of the EFA electrolyte. The answer is that CHF2COOLi does not dissolve in MFA but it dissolves in EFA, as shown in Fig. 13 [29]. Unfortunately, CHF2COOLi was found to dissolve in EC, PC, DMC, and DEC. Therefore, it is expected that the addition of MFA in usual carbonate electrolytes is not effective in improving the thermal stability with Li metal. The Li cycling efficiencies of 1 M LiPF6//EC/DMC, 1 M LiPF6// CHF2COOCH3-mixed EC  DMC (50 wt%), 1 M LiPF6/PC, 1 M LiPF6/ CHF2COOCH3-mixed PC (50 wt%), and 1 M LiPF6/CHF2COOCH3 were 70, 75, 71, 71, and 85%, respectively. These results indicate that the single solvent electrolyte of CHF2COOCH3 showed the highest cycling efficiency and that the addition of CHF2COOCH3 to EC  DMC can improve cycling efficiency. However, the addition to PC cannot improve cycling efficiency. The conductivities of 1 M LiPF6//EC/DMC, 1 M LiPF6/CHF2COOCH3mixed EC  DMC, 1 M LiPF6/PC, 1 M LiPF6/CHF2COOCH3-mixed PC, and 1 M LiPF6/CHF2COOCH3 were 12, 14, 7, 13, and 12 mS/cm, respectively. The conductivity increased by the addition of CHF2COOCH3 to EC  DMC (50 wt%)

Thermally stable fluoro-organic solvents for lithium ion battery

281

Fig. 13. CHF2COOLi in EFA (a) and MFA (b).

or PC (50 wt%). These results are attributed to the lower viscosity of CHF2COOCH3 as compared with the viscosity of EC  DMC or PC. 4. THERMAL STABILITY OF LiPF6/FLUORINATED ESTER ELECTROLYTE WITH LITHIATED CARBON ANODE [21] In this section, in order to assess the availability of LiPF6 /MFA to a lithium ion battery, thermal stability with lithiated carbon anodes was studied using a DSC. The electrochemical characteristics of the carbon anode were obtained in this electrolyte [21]. The lithiated graphite was prepared using a coin cell (2032 type, can size: 2.0 cm in diameter and 0.32 cm in height) for DSC measurement. Natural graphite (LF-18D from Chuestsu Graphite) and MCMB-6–10 (Osaka Gas Chemical) were used as the carbon anodes. MCMB-6–10 was treated as a representative of carbon materials that have a less graphitic structure. The composite electrode used in this study was prepared by mixing 95 wt% of carbon anode material with 5 wt% of polyvinylidenefluoride (PVdF) binder (KF#9100 from Kureha Chemical) dissolved in 1-methyl-2pyrrolidinone (NMP). The slurry was coated onto a copper current collector. The thermal stability of those anodes with 1 M LiPF6 /EC-DMC (1:1, by vol) or 1 M LiPF6/MFA was monitored by a DSC apparatus (Rigaku Thermo plus DSC 8230L, Rigaku). LiCoO2 electrodes used in this study were prepared by mixing 90 wt% of LiCoO2, 5 wt% of acetylene black (Denki Kagaku), and 5 wt% of PVdF-binder (KF#1300 from Kureha Chemical) dissolved in NMP. The charge/discharge cycle tests of the coin cells were applied at 25°C. Fig. 14 shows the respective DSC profiles of MFA electrolyte in coexistence with lithiated natural graphite (LF-18D) and with MCMB-6–10. For the sake of comparison, profiles of 1 M LiPF6 /EC-DMC (1:1, by vol) electrolyte in conventional lithium-ion batteries were also included. In our previous study on

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Fig. 14. DSC profiles of electrolytes with various lithiated carbon anodes: (a) natural graphite (LF-18D) in 1 M LiPF6/EC-DMC (1:1, by vol); (b) natural graphite (LF-18D) in 1 M LiPF6/MFA; (c) less graphitized MCMB (MCMB-6–10) in 1 M LiPF6/EC-DMC (1:1, by vol); (d) less graphitized MCMB (MCMB-6–10) in 1 M LiPF6/MFA.

the thermal stability of the graphite anode with 1 M LiPF6 /EC-DMC (1:1, by vol) [121, it was concluded that the small exothermic peak from 140 to 280°C originated in the reaction (SEI formation) of the electrolyte and the lithiated graphite through SEI, and that the sharp exothermic peak at 280°C originated from a direct reaction of the lithiated carbon with the electrolyte. When MFA-based electrolyte is used, a direct reaction of the lithiated carbon with the electrolyte is clearly observed at 400°C on both natural graphite (C6 Li 0.8) and MCMB-6–10 (C6 Li 0.44). The thermal stability of 1 M LiPF6 /MFA was improved by about 100°C in comparison with to that of 1 M LiPF6 /EC-DMC (1:1, by vol), though there are several small exothermic reactions from 110°C to the temperature of the main reaction, even in the case of the MFA-based electrolyte. It is well known that a graphite electrode cannot be cycled stably with propylene carbonate (PC)-based electrolyte [30,31]. It is not obvious whether MFA, as the component of the electrolyte, behaves as EC or as PC during the lithium intercalation into the graphite. Therefore in this work, two kinds of carbon materials were used for the electrode: natural graphite (LF-18D), and less graphitized MCMB (MCMB-6–10) that can be cycled with the PC-based electrolyte. Here the cycling behavior using 1 M LiPF6 /MFA was compared with that using 1 M LiPF6 /EC-DMC, where both of these carbon anodes could be cycled well. The cycling behavior of the LF-18D/LiCoO2 and MCMB-6–l0/LiCoO2 cells with 1 M LiPF6 /MFA was also evaluated, since the cycle stabilities of both carbon anodes in MFA-based electrolytes were confirmed. The cycle performances of the discharge capacities are plotted in Fig. 15 [21]. Also, in this set of electrodes the discharge capacity of the cell with both electrolytes retained the

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283

Fig. 15. The cycle number dependences of the discharge of LiCoO2/the natural graphite (LF-18D) ion cell and LiCoO2/the less graphitized MCMB (MCMB-6–10 ion cell). (The cells were cycled at 25°C between 3.0 and 4.2 V at 0.2 mA/cm2 with a constant voltage charge at 4.2 V for 3 h.) (䊐) 1 M LiPF6/EC-DMC (1:1, by vol); () 1 M LiPF6 /MFA.

value as high as 180 mAh/g after 50 cycles. These results indicate that the MFAbased electrolyte has adequate properties for the use with various sets of electrode materials in lithium-ion batteries. ACKNOWLEDGEMENTS This work was supported by CREST of JST (Japan Science and Technology Corporation). The authors thank Daikin Industries, the Society of Advanced Battery Technologies, Osaka Science & Technology Center, NEC, Japan Storage Battery, Central Glass and Mitsubishi Heavy Industry for financial support. REFERENCES [l] J.R. Dahn, E.W. Fuller, M. Obravae, and U. von Sacken, Solid State Ionics, 69 (1994) 265. [2] D. Wainwright, J. Power Sources, 54 (1995) 192. [3] U. von Sachen, E. Nodwell, A. Sundher, and J.R. Dahn, J. Power Sources, 54 (1995) 240. [4] H. Arai, S. Okada, Y. Sakurai, and J. Yamaki, J. Electrochem. Soc., 144 (1997) 3117. [5] Z. Zhang, D. Fouchard, and J.R. Rea, J. Power Sources, 70 (1998) 16. [6] H. Arai, S. Okada, Y. Sakurai, and J. Yamaki, Solid State Ionics, 109 (1998) 295. [7] A.M. Anderssson, K. Edstrom, and J.O. Thomas, J. Power Sources, 81–82 (1999) 8. [8] A. Okamato, T. Sasaki, S. Komatsu, K. Nakamitsu, H. Tsukamoto, and M. Mizutani, GS News Tech. Rep., 56 (l) (1999) 18. [9] G.G. Bottle, R.E. White, and Z. Zhang, J. Power Sources, 97–98 (2001) 570. [10] J.P. Cho and B.Park, J. Power Sources, 92 (2001) 35.

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[11] K. Edstrom, A.M. Andersson, A. Bishop, L. Fransson, J. Lindgren, and A. Hussenius, J. Power Sources, 97–98 (2001) 87. [12] J. Yamaki, H. Takatsuji, T. Kawamura, and M. Egashira, Solid State Ionics, 148 (2002) 241. [13] Y. Baba, S. Okada, and J. Yamaki, Solid State lonics, 148 (2002) 311. [14] T. Kawamura, A. Kimura, M. Egashira, S. Okada, and J. Yamaki, J. Power Sources, 104 (2002) 260. [15] N. Katayama, T. Kawamura, Y. Baba, and J. Yamaki, J. Power Sources, 109 (2002) 321. [16] J.O. Besenhard, W.K. Appel, L.H. Lie, G.H. Wrodnigg, K.-C. Moeller, and M. Winter, Abstracts of the Second Hawaii Battery Conference, Organized by A.N. Dey, Big Island of Hawaii, Jan. 4–7, 1999, p. 181. [17] T. Nakajima, K. Dan, and M. Koh, J. Fluorine Chem., 87 (1998) 221. [18] T. Nakajima, K. Dan, M. Koh, T. Ino, and T. Shimizu, J. Fluorine Chem., 111 (2001) 167. [19] J. Yamaki, I. Yamazaki, M. Egashira, and S. Okada, J. Power Sources, 102 (2001) 288. [20] K. Sato, I. Yamazaki, S. Okada, and J. Yamaki, Solid State lonics, 148 (2002) 463. [21] M. Ihara, B.T. Hang, K. Sato, M. Egashira, S. Okada, and J. Yamaki, J. Electrochem. Soc., 150(11) (2003) A1476. [22] D. Aurbach, A. Zaban, Y. Ein-li, I. Weissman, O. Chusid, B. Markovsky, M. Levi, E. Levi, A. Schechechter, and E. Granot, J. Power Sources, 68 (1997) 91. [23] S.E. Sloop, J.K. Pugh, S. Wang, J.B. Kerr, and K. Kinoshita, Electrochem. Solid State Lett., 4 (2001) A42. [24] E. Peled, J. Electrochem. Soc., 126 (1979) 2047. [25] R. Herr, Electrochim. Acta, 35 (1990) 1257. [26] J. Yamaki, T. Tanaka, I. Watanabe, M. Egashira, and S. Okada, Abstracts of 204th Meeting of the Electrochem. Soc., ECS, Orland, FL, 2003, Abstract No. 290. [27] R. Gupta and A. Manthiram, J. Solid State Chem., 121 (1996) 483. [28] E. Zhecheva and R. Stoyanova, J. Solid State Chem., 109 (1994) 47. [29] J. Yamaki, T. Tanaka, M. Ihara, K. Sato, and S. Okada, Extended Abstracts of LiBD 2003 – Electrode Materials, Arcachon, France, 2003, Abstract No 34. [30] A.N. Dey and B.P. Sullivan, J. Electrochem. Soc., 117 (1970) 222. [31] M. Arakawa and J. Yamaki, J. Electroanal. Chem., 219 (1987) 273.

Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter13

Physical and electrochemical properties and application to lithium batteries of fluorinated organic solvents Yukio Sasaki Department of Nanochemistry, Faculty of Engineering, Tokyo Polytechnic University, 1583 Iiyama, Atsugi, Kanagawa 243-0297, Japan 1. INTRODUCTION Efforts to develop improved solvents for rechargeable lithium batteries with high-energy density, oxidation durability, high liquidus range and non-flammability for power sources of mobile equipments and electric vehicles have been made by many researchers [1–3]. Fluorinated organic solvents show very different physical properties compared with those of common organic solvents because of very high electronegativity, high ionic potential and low polarizability of the fluorine atom. In general, partially fluorinated organic solvents among the fluorinated organic solvents show fairly high polarity in comparison with that of perfluoro organic solvents. Therefore, one of the appropriate methods to find a solvent with good cell performance is the introduction of fluorine atoms into the solvent molecules. There are three methods for partial fluorination of organic compounds: chemical fluorination using fluorinating reagents, electrolytic fluorination and direct fluorination using elemental fluorine (F2 gas). The direct fluorination is the simplest method to prepare partially fluorinated organic solvents; in addition, it makes possible to obtain many interesting fluorinated organic solvents. In this chapter, physical and electrochemical properties of the partially fluorinated solvents by the use of F2 gas to develop several important organic solvents for practical lithium batteries are mainly described. 2. AN APPARATUS FOR DIRECT FLUORINATION Fig. 1 shows an schematic representation of the apparatus for direct fluorination. F2 gas was generated by an electrolysis of KF·2HF (Fluorodec TM30; Toyo Tanso Co. Ltd.) at 80–90°C. Fluorination of solvents was carried out using diluted F2 gas by N2.

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Yukio Sasaki

N2

Bypass

F2

NaF Pellet

Al2O3Scrubbers N2

Trap

Electrolytic bath

HF

Gas dilution system −

Anode : 2F → F2 ↑ + 2e−

Reactor

Cathode : 2H+ + 2e− → H 2 ↑ Fig. 1. Schematic representation of the apparatus for direct fluorination.

3. DIRECT FLUORINATION OF γ -BUTYROLACTONE (γ -BL) Direct fluorination of γ-BL was carried out using 20% F2/N2 gas at 30°C [4]. The gas was poured onto the γ-BL sample surface by stirring in PFA (perfluoroalkoxy) vessel at 30°C. Fig. 2 shows a typical gas chromatogram (GC) after fluorination for 16 h in the presence of NaF as a HF scavenger. In Fig. 2, three monofluorinated γ-BL (α-, β-, and γ-F–γ-BL) and six difluorinated γ-BL (F2-γBL) derivatives were formed according to Scheme 1. An important characteristic of direct fluorination is the fact that many γ-BL derivatives are usually produced as shown in Scheme 1. In contrast, γ-F–γ-BL is selectively produced by electrochemical fluorination [5]. Fig. 3 shows the yield of the fluorinated γ-BL derivatives with time in the presence (a) and absence (b) of NaF. The γ-F–γ-BL formation proceeded preferentially in the presence of NaF as a HF scavenger. However, γ-F–γ-BL was not stable under some conditions (particularly hydrolysis). It seemed that γ-F–γ-BL was decomposed by the HF produced. On the other hand, β-F–γ-BL formation became dominant after NaF was consumed at around 10 h. A similar fluorination was carried out in the absence of NaF as shown in Fig. 3(b). Furthermore, from the total yield in both in Fig. 3 (a) and (b), it is found that the fluorination of γ-BL occurs quantitatively. 3. 1. Physical properties

Variations of dielectric constants, densities and viscosities [6] of a mixed solvent (F-γ-BLmix) of α-F–γ-BL and β-F–γ-BL (molar ratio 3:7), that have

Physical and electrochemical properties and application to lithium batteries of fluorinated organic solvents

287

Fig. 2. A typical GC of the fluorinated γ-BL derivatives. F O 20% F2/N2 30°C

O F

F O

+

O

O

O

γ -BL

α-F-γ-BL

β-F-γ-BL

O +

+ O

γ-F-γ-BL

F2 O O F2-γ-BL

Scheme 1.

Fig. 3. Time – yield curves of fluorinated γ-BL derivatives in the presence (a) and absence (b) of NaF.

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Yukio Sasaki

almost the same boiling points in Fig. 2, obtained by fractional distillation of the fluorinated γ-BL, are shown in Figs. 4, 5 and 6, respectively. The dielectric constant of F-γ-BLmix was 80.3 at 25°C, which is more than twice that of γ-BL in a whole temperature range. Similarly, viscosity and density of F-γ-BLmix increased to much higher values than those of γ-BL. This shows that polarity and molecular interaction increase in F-γ-BLmix with the introduction of a fluorine atom with high electron withdrawing.

100

Dielectric constant

80 F- γ-BLmix 60 γ-BL 40

20 0

20

40 60 Temperature /°C

80

100

Fig. 4. Temperature dependence of γ-BL and F-γ-BLmix on dielectric constants. 1.5

Density/103kgm-3

1.4 F- γ -BLmix 1.3

1.2 γ-BL 1.1

1

0

20

40 60 Temperature/°C

80

Fig. 5. Temperature dependence of γ-BL and F-γ-BLmix on densities.

100

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289

Viscosity/cP

6

4 F-γ-BLmix γ-BL 2

0

0

20

40 60 Temperature/°C

80

100

Fig. 6. Temperature dependence of γ-BL and F-γ-BLmix on viscosities.

3. 2. Electrochemical properties

Fig. 7 shows the plot of specific conductivities against temperature in γ-BL (a) and F-γ-BLmix (b) solutions containing 1.0 mol dm3 electrolytes [6]. The conductivities of F-γ-BLmix solutions are lower than those of γ-BL solutions. By comparing dielectric constant with viscosity in Figs. 4 and 6, it is found that the conductivities are affected much by viscosities rather than dielectric constants of γ-BL and F-γ-BLmix. In general, the introduction of a fluorine atom to organic solvents tends to increase their oxidation durability and decease their reduction durability. Fig. 8 shows the i–E curves in PC, γ-BL and F-γ-BLmix solutions containing 1.0 mol dm3 LiPF6. Two oxidation peaks in F-γ-BLmix solution were detected at 6.4 and 5.9 V, which correspond to α-F- and β-F-γ-BL, respectively. The oxidation potentials of these peaks were lower than that of PC. On the other hand, two reduction peaks in F-γ-BLmix solution were also observed at 1.1 and 1.3 V, which were higher than that of γ-BL. It is well known that the reductive decomposition of electrolytes contributes to form a passivation film on the electrode. It is called as solid electrolyte interface (SEI) and plays a very important role in rechargeable lithium batteries. Fig. 9 shows lithium electrode cycling efficiencies in F-γ-BLmix, γ-BL and PC solutions containing 1.0 mol dm3 LiPF6, and SEM images on a Ni electrode after the discharge of the 10th cycle. The cycling efficiency in F-γ-BLmix solution is maintained at about 75% in a higher range of cycles. From the SEM photographs, it is found that the electrode surface (film) in F-γ-BLmix solution is very homogeneous, and consists of a uniform and small grain size, compared with those in γ-BL and PC solutions. Accordingly, the cycling efficiency is highly

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Specific conductivity / mScm-1

LiPF6

20

LiClO4

10

(a)

0 0

LiBF4

20

40 60 Temperature /°C

80

100

Specific conductivity / mScm-1

30

30

20 LiPF6 LiClO4

10 LiBF4

0

0

20

(b)

40 60 Temperature /°C

80

100

Fig. 7. Temperature dependence of γ-BL(a) and F-γ-BLmix(b) solutions on specific conductivities. 1.5 PC

i / mA cm-2

1.0

F-γ-BLmix

0.5

γ-BL

PC

0 -0.5 -1.0 γ-BL -1.5

0

1.0

F-γ-BLmix 2.0

3.0

4.0

5.0

6.0

7.0

E / V vs. Li / Li+

Fig. 8. i–E curves in PC, γ-BL and F-γ-BLmix solutions containing 1.0 mol dm3 LiPF6 using a Pt working electrode, Li wire counter and reference electrodes at a scan rate of 5 m Vs1 at 25°C.

dependent on the morphology of the films. F-γ-BLmix is a good electrolyte for forming a uniform film on the electrode (anode). Fig. 10 shows cycling efficiencies of Li/LiCoO2 coin cell using 1.0 mol dm3 LiPF6 solutions at 25°C. F-γ-BLmix solution had much higher efficiency (which is close to that of PC) than γ-BL solution. Accordingly, it is found that Fγ-BLmix has enough oxidation durability against the cathode (LiCoO2) and the film does not interfere with the dissolution of Li ion. 4. DIRECT FLUORINATION OF ETHYLENE CARBONATE (EC) Direct fluorination of ethylene carbonate (EC) [7] was successfully carried out to provide 4-fluoro-ethylene carbonate (FEC) as in Scheme 2.

Physical and electrochemical properties and application to lithium batteries of fluorinated organic solvents

291

100 F- γ -BLmix

Efficiency/%

80 60 40

γPC

20 γ-BL 0 0

20

40

60

80

100

Cycle number

γ-BL

F-γ-BLmix

PC

⋅ 7.5 × 7.5

Fig. 9. Upper panel: Variation of lithium electrode cycling efficiencies (ip  is  1.0 mA cm2, Qp  0.3 cm2) in F-γ-BLmix, γ-BL and PC solutions containing 1.0 mol dm3 LiPF6 at 25°C and Lower panel: SEM images on a Ni electrode after the discharge of the 10th cycle. 100

Efficiency/%

95

90

85

80

75 0

10

20 30 Cycle numer

40

50

Fig. 10. Cycling efficiencies of Li/LiCoO2 cells in constant current – constant voltage (CCCV) mode in 1.0 mol dm3 LiPF6 electrolytes at 25°C (䊐) F-γ-BLmix , (䊊)γ-BL, (䉭)PC.

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FEC was also further fluorinated to give three difluoro derivatives as shown in Scheme 3. Fluorination of EC was strongly dependent on a choice of solvents as shown in Table 1. At 0°C, the conversion of EC was low and the reaction rate was slow. On increasing the temperature, the reaction proceeded faster but there was no increase in selectivity of FEC. F O

30% F2N2

O

O

O

50˚C, no solvent O

O EC

FEC (70%)

Scheme 2.

F

F

F O

O

30% F2/N2

F

F

F

F O

O

+

50, no solvent O

O

FEC

cis-4.5-difluoroEC(c-DFEC)(11%)

O

O

+

O

O trans-4.5-difluoroEC(t-DFEC) (59%)

O O

4.4-difluoroEC(DFEC) (5%)

Scheme 3.

Table 1 Direct fluorination of EC Solvent

Ratio (w/w) Reaction F2a Passing time conversion Selectivityb EC/solvent temperature (eq) (ml min1) (%) of FEC (%) (°C)

CHCl3

5:95

0

0.6

50

0

0

(CF3)2CFCF2CF2CF3

2:98

0

0.4

50

12

95

(CF3)2CFCF2CF2CF3

2:98

50

1.0

50

33

79

CClF2CCl2F

10:90

50

1.0

50

52

68

HF

17:83

10

1.0

50

71

93

No solvent

100:0

50

1.7

350

85

86 (70)c

a Fluorine, 30% mixture in nitrogen, was passed through the reaction mixture (50–350/ml min1. bGC analysis. c Isolated yield (in parentheses) based on EC.

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293

4. 1. Physical properties

Table 2 [7] shows physical properties of fluorinated EC. An introduction of fluorine atom tends to decrease the boiling point (b.p.) and melting point (m.p.). Table 2 Physical properties of fluorinated EC Compounds

O

B.p. (°C)

M.p. (°C)

Density (103 kg m3)

Viscosity (cP)

Dielectric constant

238

37

1.321

1.9a

90a

210

17.3

1.497

4.1b

78.4b

F

187

56.6

1.592

2.3c

F

129

7.8

1.508

2.5a

37.1a

134

3.2

1.567

2.1b

34b

O

O F

O

O

O F

O

O

O F

O

O

O F F O

O

O a

At 40°C. bAt 23°C. cAt 60°C.

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In addition, the boiling point, melting point and the density of c-DFEC among three difluorinated ECs, which have two fluorine atoms, became higher. This means that c-DFEC is thermally stable and has a high molecular interaction. However, the viscosity of c-DFEC is lower than that of t-DFEC. Variation of densities (a) and viscosities (b) with temperature of FEC and EC is shown in Fig. 11. The viscosity of FEC increases rapidly with a decrease in the temperature. 4. 2. Electrochemical properties

Fig. 12 shows cyclic voltammograms of FEC synthesized chemically and PC solutions containing 1 mol dm3 LiPF6 [8]. A necessary requirement for

1.6 4 Viscosity/cP

Density/103kgm-3

FEC 1.5

1.4

FEC

2

EC EC 1.3

1 0

(a)

3

20 40 Temperature/°C

60

(b)

0

20 40 Temperature/°C

60

Fig. 11. Temperature dependence of FEC and EC on dielectric constants (a) and viscosities (b).

Fig. 12. Cyclic voltammograms in FEC and PC solutions containing 1.0 mol dm3 LiPF6 at a scan rate of 0.4 Vs1. A lithium wire was used as a reference electrode, a Pt wire as a working electrode and Pt mesh as a counterelectrode.

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295

electrolytes of lithium-ion battery is stability to oxidation at potentials up to 4.3 V vs. Li/Li. The cyclic voltammetry was conducted in the range of 3.0–5.0 V on FEC electrolyte and for comparison, on EC/DMC and EC/DEC electrolyte. FEC electrolyte is at least as stable as EC/DMC electrolyte to oxidation. Fig. 13 shows the potential curve plotted as a function of the cell capacity per weight of graphite for a C/LiCoO2 cell using 1mol dm3 LiPF6 in FEC/PC/EC (1:3.5:3.5, by vol) [8]. In Fig.13, a fairly good reversible capacity is obtained. The cycling efficiency of the cell has become higher than 99.5% as shown in Fig.14. Fig.15 shows a life cycle of

Fig. 13. Potential profile of the first cycle of a C/LiCoO2 cell using 1 mol dm3 LiPF6 FEC:PC:EC (1:3.5:3.5, by vol) electrolyte. The cell is cycled at a 20 h rate.

Fig. 14. Cycling efficiency of the C/LiCoO2 cell. The cell was charged and discharged at a 20 h rate for the first three cycles between 3.2 and 4.0 V.

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Fig. 15. A life cycle plot of discharge capacity vs. cycle number for the C/LiCoO2 cell.

the cell. The decrease in the capacity of the cell is around 37% at 200th cycle. The decrease is probably due to incomplete intercalation of the graphite anode. 5. DIRECT FLUORINATION OF DIMETHYL CARBONATE (DMC) Fluorination of DMC was carried out using 15% F2 gas diluted using N2 [9–11]. The gas was poured onto the DMC sample in the same way as γ-BL at 25°C. Fig.16 shows GC–MS analysis of the fluorinated DMC derivatives after fluorination for 24 h. RT is the retention time. It is found that four fluorinated DMC derivatives are formed, which are monofluorinated dimethyl carbonate (MFDMC), two difluorinated dimethyl carbonates (DFDMC and gem-DFDMC), and trifluorinated dimethyl carbonate (TFDMC) as shown in Fig.16. Fig. 17 shows the variation in yields of the fluorinated DMC derivatives with time. The initial fluorinated DMC derivative is MFDMC. The yield of MFDMC increases with decreasing DMC concentration. The yield decreases gradually after the maximum value at about 15 h. On the contrary, formation of DFDMC and gemDFDMC begins slowly after about 10 h, and then TFDMC formation begins successively. DFDMC formation continues to increase after TFDMC formation, compared with that of gem-DFDMC. Accordingly, it is supposed that TFDMC formation takes place mainly through gem-DFDMC. From the total yield of the fluorinated DMC derivatives, fluorination of DMC occurs quantitatively. By considering the variation in the yields with time, the formation process of the fluorinated DMC derivatives can be presumed as shown in Fig. 18. 5. 1. Physical properties

In Fig. 16, the boiling point of DFDMC is very different from that of gemDFDMC though both DFDMC and gem-DFDMC have two fluorine atoms. This

Physical and electrochemical properties and application to lithium batteries of fluorinated organic solvents

H3C

O

C

O

O

F

H3C

O

O

C

F F O O H 2C CH C CH3 F O

O

F

DMC

C

F

O

F F O O H2C C CH2

CH2

O

O

TFDMC

CH

O

H3C

297

DFDMC

MFDMC

gem-DFDMC 0.0

2.0

4.0

6.0

Time/ min

127

RT: 2.1min

(gem-DFDMC) 91 (DMC)

RT: 2.3 min

145 (TFDMC)

RT: 3.1 min 109 (MFDMC)

RT: 3.5 min

127 (DFDMC)

RT: 4.9 min 50

60

70

80

90

100

110

120

130

140

150

Mass to charge (m z-1)

Fig. 16. GC–MS analysis of the fluorinated DMC derivatives after fluorination for 24 h. 100 DMC

90

Total

80 MFDMC

Yield/%

70 60 50

DFDMC

40

TFDMC

30 20

gem -DFDMC

10 0 0

5

10

15

20

Time/h

Fig. 17. Variation in yields with time of the fluorinated DMC derivatives.

suggests that physical properties of the fluorinated DMC with the same molecular weight are highly dependent on the positions of fluorine atoms introduced into the DMC molecule. In addition, the boiling point of gem-DFDMC is slightly

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lower than that of DMC. Therefore, it was actually difficult to obtain the pure gem-DFDMC by distillation. Fig. 19 shows dielectric constants of the fluorinated DMC derivatives with temperature. The dielectric constant of DFDMC is very high compared with those of TFDMC and MFDMC. The higher dielectric constant of DFDMC is dependent on the higher polarity of DFDMC due to the introduction of fluorine atoms with high electron withdrawing into methoxy groups (CH3O). Furthermore, every dielectric constant of the fluorinated DMC derivative decreases linearly with increase in temperature. This means that the orientation of the dipoles in the DMC derivatives reduces with an increase in the temperature. On the other hand, the dielectric constant of DMC with lower polarity remains almost constant in the whole temperature range. Variation of viscosities and densities of the fluorinated F O O F CH2 C H2C O

DFDMC O H3C

C

O

CH3

O H3C

F O O F C CH H2C F O

O F C CH2

O

O

DMC

MFDMC

TFDMC O O F C CH H3C F O

gem-DFDMC

Fig. 18. Formation process of the fluorinated DMC derivatives.

15 Dielectric constant

DFDMC TFDMC

10

MFDMC 5 DMC

0

20

40 Temperature/°C

60

Fig. 19. Temperature dependence on dielectric constants of the fluorinated DMC derivatives measured with 1 MHz.

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299

DMC derivatives with temperature are shown in Figs. 20 and 21, respectively. The viscosities decrease in the order DFDMC  TFDMC  MFDMC  DMC, similar to those of dielectric constants in Fig. 19. However, this order is different from that of densities in Fig. 21. The highest density of TFDMC seems to be due to the higher molecular weight of TFDMC with three fluorine atoms. 5. 2. Electrochemical properties

Table 3 shows the HOMO and LUMO energies for fluorinated DMC derivatives calculated by B3LYP. In HOMO energies, oxidation durability decreases in

2

Viscosity/cP

DFDMC TFDMC

1

MFDMC

DMC

0

20

40 Temperature / °C

60

Fig. 20. Temperature dependence of the fluorinated DMC derivatives on viscosity.

Density/103kgm-3

1.6 TFDMC 1.4 DFDMC 1.2 MFDMC DMC 1 0

20

40 Temperature/°C

60

Fig. 21. Temperature dependence of the fluorinated DMC derivatives on densities.

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the order TFDMC  DFDMC  gem-DFDMC  MFDMC  DMC. On the other hand, reduction durability increases in the same order as shown for LUMO energies. Therefore, it is found that the potential windows of the fluorinated DMC derivatives are shifted to the oxidation sides. In addition, the results described above was clearly evidensed by comparing the HOMO and LUMO energies for gem-DFDMC and DFDMC, which have two fluorine atoms together. Fig. 22 shows i–E curves for the fluorinated DMC derivatives except TFDMC, which has poor solubility of 0.1 mol dm3 LiPF6 at 25°C. In Fig. 22, the order of oxidative decomposition voltage of the derivatives is DFDMC  MFDMC  DMC, as expected from HOMO energies in Table 3. In general, it seems that the oxidation durability of the fluorinated organic solvents increases with an increase in the number of fluorine atoms introduced into the solvent molecule. Table 3 Correlation of HOMO and LUMO energies for fluorinated DMC derivatives calculated by B3LYP HOMO energy (kJ mol1)

LUMO energy (kJ mol1)

DMC

790.8

6.9

MFDMC

835.8

1.0

gem-DFDMC

873.3

16.7

DFDMC

877.1

37.2

TFDMC

881.1

55.6

5 DMC

i/mA cm-2

MFDMC

DFDMC

0

-5 0

2

4

6

8

E / V vs. Li/ Li+

Fig. 22. i–E curves in 1.0 mol dm3 LiPF6 solutions using a Pt working electrode, Li wire counter and reference electrodes at a scan rate of 5 m Vs1 at 25°C.

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301

Table 4 shows specific conductivities at various LiPF6 concentrations in DFDMC, MFDMC and DMC solutions at 25°C. The specific conductivities in DFDMC solutions (except 0.25 mol dm3) with higher dielectric constants, in Fig. 19, are considerably lower than those in MFDMC and DMC solutions. Accordingly, it is considered that the conductivities are affected much by the viscosity of these three solvents as shown in Fig. 20. Specific conductivities at various LiPF6 concentrations in EC-DFDMC, EC-MFDMC and EC-DMC equimolar binary solutions at 25°C are shown in Fig. 23. The conductivities in Table 4 Specific conductivities at various LiPF6 concentrations in DFDMC, MFDMC and DMC solutions at 25°C Specific conductivities (mS cm1) Concentration (mol dm3)

DFDMC

MFDMC

DMC

0.25

0.70

0.76

0.57

0.50

1.36

2.00

2.33

1.00

2.39

4.21

5.99

1.25

2.80

4.93

7.33

1.50

3.10

5.30

7.90

1.75

–

5.24

8.15

2.00

–

4.99

7.71

Specific conductivity / mScm-1

12 EC-DMC

10 8

EC-MFDMC

6 EC-DFDMC

4 2 0

0

0.5

1

1.5

2

LiPF6concentration / mol dm-3

Fig. 23. Variation of specific conductivities with LiPF6 concentrations in EC-MFDMC and EC-DMC equimolar binary solutions at 25°C.

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these binary solutions show the maximum values at around 1 mol dm3 LiPF6 concentrations. For example, the conductivities in the EC-DFDMC and ECMFDMC solutions are 6.91 and 8.47 m S cm1, respectively. These conductivities are available for practical lithium batteries. Fig. 24 shows lithium electrode cycling efficiencies (charge – discharge coulombic efficiencies of cycling for lithium electrode) in DFDMC, MFDMC, EC-DFDMC, EC-MFDMC and EC-DMC equimolar binary solutions containing 1 mol dm3 LiPF6 at 25°C. Both DFDMC and EC-DFDMC solutions show higher efficiencies than those of EC-MFDMC, MFDMC and EC-DMC solutions at a higher range of cycle numbers. In particular, EC-DFDMC solution shows the highest efficiency of more than 80% at around the 100th cycle. This is a good electrolyte for rechargeable lithium batteries. Fig. 25 shows SEM photographs on Ni electrode (working) in these solutions after a discharge of the 40th cycle. In Fig. 25, a dendrite formation is observed on the electrode in EC-DMC solution showing a rapid decrease with an increase in the cycle number in Fig. 24. However, the surface films formed on the electrode in MFDMC, DFDMC and EC-MFDMC solutions are very homogeneous, and consist of uniform and very small grain sizes. In addition, it seems that the films are thin compared with that of EC-DFDMC solution, which shows the highest efficiency. In analogy with Fγ-BLmix in Fig. 9, the lithium electrode cycling efficiency is related to the morphology and thickness of the films (SEI) formed on the electrode. Fig. 26 shows discharge capacities of Li/LiCoO2 coin cells in EC-DMC, EC-MFDMC and EC-DFDMC solutions. The EC-MFDMC and EC-DMC solutions show higher capacities,  100 mAh g1, at the 50th cycle. On the other 100 DFDMC

EC-DFDMC

Efficiency / %

80

60 EC-MFDMC MFDMC EC-DMC

40

20

0

0

20

40 60 Cycle number

80

100

Fig. 24. Variation of lithium electrode cycling efficiencies in DFDMC, MFDMC, ECDFDMC, EC-MFDMC and EC-DMC equimolar binary solutions containing 1.0 mol dm3 LiPF6 at 25°C.

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Fig. 25. SEM photographs on Ni electrode in MFDMC, DFDMC, EC-DMC, EC-MFDMC and EC-DFDMC equimolar binary solutions containing 1.0 mol dm3 LiPF6 after the discharge of the 40th cycle. 150

Capacity/mAh g-1

EC-MFDMC

100 EC-DMC EC-DFDMC

50

0

0

10

20 30 Cycle number

40

50

Fig. 26. Discharge capacities of Li/LiCoO2 cells in EC-DMC, EC-MFDMC and EC-DFDMC equimolar binary solutions containing 1.0 mol dm3 LiPF6 at 25°C.

hand, the capacity in EC-DFDMC solution, which showed the highest lithium electrode cycling efficiency in Fig. 24, is very low. This suggests that the intercalation of Li ion to cathode (LiCoO2) does not proceed smoothly in EC-DFDMC solution. In general, an intercalation to cathode is influenced by many factors, such as characteristics (wetting) of the electrolyte and the solvent (selective ion solvation), conductivity and viscosity in electrolytes, etc. It is thought that the lower capacity in EC-DFDMC solution is dependent on the lower conductivity

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EC-DMC

Efficiency / %

100

EC-MFDMC

80 60

EC-DFDMC

40 20 0

0

10

20 30 Cycle number

40

50

Fig. 27. Variation of cycling efficiencies of Li/LiCoO2 cells in EC-DMC, EC-MFDMC and EC-DFDMC equimolar binary solutions containing 1.0 mol dm3 LiPF6 at 25°C

(Fig. 23) due to high viscosity of DFDMC in Fig. 20. The cycling efficiencies of Li/LiCoO2 cells are shown in Fig. 27. EC-DMC and EC-MFDMC solutions show almost 100% efficiencies. These solutions have good recycleability for the Li/LiCoO2 cell. On the contrary, EC-DFDMC solution showed lower efficiency, as presumed by the discharge capacity in Fig. 26. ACKNOWLEDGEMENT The author thanks Dr. M. Takehara and Mr. R. Ebara, Mr. M. Hagiyama and Mr. S. Watanabe, graduate students of Tokyo Polytechnic University, for their help in preparing this chapter. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11]

K. Xu, S. Zhang, J.L. Allen, and T.R. Jow, J. Electrochem. Soc., 149 (2002) A1079. M.S. Ding, K. Xu, and T.R. Jow, J. Electrochem. Soc., 149 (2002) A1489. N. Nanbu, T. Shibazaki, and Y. Sasaki, Electrochemistry, 71 (2003) 1205. Y. Sasaki, R. Ebara, N. Nanbu, M. Takehara, and M. Ue, J. Fluorine Chem., 108 (2001) 117. M. Hasegawa, H. Ishii, and T. Fuchigami, Tetrahedron Lett., 43 (2002) 1503. M. Takehara, R. Ebara, N. Nanbu, M. Ue, and Y. Sasaki, Electrochemistry, 71 (2003) 1172. M. Kobayashi, T. Inoguchi, T. Iida, T. Tanioka, H. Kumase, and Y. Fukai, J. Fluorine chem., 120 (2003) 105. R. McMillan, H. Slegr, Z.X. Shu, and W. Wang, J. Power Sources, 81–82 (1999) 20. Y. Sasaki, M. Takehara, S. Watanabe, M. Oshima, N. Nanbu, and M. Ue, Solid Sate Ionics, in press. Y. Sasaki, M. Takehara, S. Watanabe, N. Nanbu, and M. Ue, J. Fluorine chem., 125 (2004) 1205. M. Takehara, S. Watanabe, N. Nanbu, M. Ue, and Y. Sasaki, Synth. Commun., 34 (2004) 1367.

Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 14

PVdF-based polymers for lithium batteries Jean-Yves Sancheza, Fannie Alloina, and Johanna Saunierb a

LEPMI, UMR 5631, CNRS, INPG, UJF, BP 75, F.38402 Saint-Martin-d’Hères, France b

EA 401 Faculté de pharmacie 5, rue J.B.Clément Chatenay, 92296 Malabry, France 1. INTRODUCTION The world market of electrochemical energy sources (EES) has been growing continuously at a rapid rate for about a decade. In the consumer market, in particular the 4C market, i.e. cellular phones, camcorders, computers, cordless tools, non-rechargeable batteries still have a prevalent position but rechargeable batteries should improve from 2010 (Fig. 1). Now, even in the 4C market, fuel cells, PEMFC (proton-exchange membrane fuel cells) and/or DMFC (direct methanol fuel cell), may, according to their technological improvements, also have a significant share of the market around 2010. In addition to the growing rate of the 4C market [1,2], for about the last 14 years, voluntary policies originating from California have led to the implementation of federal programmes with the aim of developing electrical vehicles. Further to these policies, national consortium USABC (U.S. Advanced Battery Consortium) has funded several programmes on rechargeable batteries. Mid- and long-term performances, i.e. specific and volumetric energy, cyclability at 80% DOD (depth of discharge), peak power and price per kWh were required and certain batteries as lithium-polymer batteries [3] have met most of the USABC requirements. Later, the U.S. PNGV programme (Partnership for New Generation Vehicles) aimed at decreasing the energy consumption of thermal vehicles and to providing alternatives to traditional vehicles, in particular hybrid vehicles which associate electric and thermal engines. As severe limitations still remain with regard to the rate of charge of the best batteries, PEMFC appear as a promising alternative for electric vehicles. Not only environmental issues but also the depletion of oil resources, foreseen for one half of the present century, will influence battery and car manufacturers to develop efficient, clean and safe EES.

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Metal sprayed contact

0

VOx/collector

Li

Metal sprayed contact

Fig. 1. Laminated lithium-polymer electrolyte cell assembly.

Although the importance of the polymer component in batteries is often underestimated, it is prevalent in fuel cells and in large rechargeable batteries as it conditions both the electrochemical performances and the safety of the EES. In this chapter, attention will be focused on polyvinylidene fluoride (PVdF)-based polymer electrolytes, in the frame of lithium batteries. We will, however, first provide an overview of the different options involving polymers as electrolyte components of lithium batteries. 2. ELECTROLYTE COMPONENTS OF LITHIUM BATTERIES 2.1. Polymer electrolytes

The word polymer electrolyte refers a priori to salts or acids dissolved in a host polymer to give salt – polymer complexes. The salt dissolution, the ion-pair dissociation, the ion solvation and mobility are ensured by the host polymer. The complexes between a lithium salt such as LiClO4 or LiI and polyoxiranes as poly(oxyethylene) POE ([CH2–CH2–O]n), whose trivial name is polyethylene oxide, or poly(oxypropylene) POP ([CH2–CH(CH3)–O]n) belong to this category. Such complexes have been known for about 40 years [4,5], and their ionic conductivity is known since mid-1970s [6,7], and their use as electrolyte for lithium batteries was proposed in 1979 by Armand et al. [8]. In addition, host polymers with an appropriate electrochemical stability window must satisfy the following requirements: ●

Provide high concentration in charge carriers. This means that they must be able to dissolve appropriate lithium salts, up to high concentrations, to ensure

PVdF-based polymers for lithium batteries

●

●

307

the ion’s solvation. Solvating ability vs. cations is related to the Lewis basicity of the solvent while solvating ability vs. anions is related to the Lewis acidity of the solvents. Donor Number (DN) and acceptor number (AN) tables are very convenient scales to select the solvent, i.e. the host polymer. Due to the use of highly reductive negative electrodes, Li metal or LiC6, solvents with high AN (hydrogen bonding solvents) are excluded. Provide high ionic mobility. This implies amorphous polymer electrolytes with glass transition temperatures, Tg, as low as possible, and restricts the use of polar monomer repeat units that might favour the ion-pair dissociation but would increase Tg. Provide thermomechanical properties allowing the polymer electrolyte film to be shaped to the thinnest form as possible, while maintaining a high safety level. The gain in film thickness allows compensating, at least partially, for the ohmic drop in the electrolyte. As the polymer electrolyte must be amorphous under the battery operating conditions, the resulting polymer electrolytes, whose Tg values must be as low as possible, have poor mechanical properties. Their reinforcement through their cross-linking is therefore required, although an alternative was recently proposed to reinforce linear and cross-linked polymer electrolytes with natural nanofibres [9–11].

2.2. Set liquid electrolyte  separator vs. gelled polymer electrolytes 2.2.1. Liquid electrolytes

Due to the highly reductive negative electrode, i.e. lithium metal and, to a lesser extent, lithium graphite, lithium battery electrolytes must fulfil severe constraints. Among the negative electrodes, lithium metal is obviously the most restricting one, but lithium graphite induces other constraints in relation to the solvent co-intercalation and to the formation of a stable passivation layer. As a result the liquid solvents must be aprotic and, therefore, have a low AN, i.e. they have a poor solvating ability vs. anions. Therefore, they must be selected from among solvents with good solvating ability vs. cations, i.e. with a fairly high DN. Lithium-ion battery electrolytes mainly consist of a molar solution of a superacid lithium salt, such as LiPF6 or LiBF4, in a mixture of two or three solvents. Thus ethylene carbonate (EC) provides high dielectric constant, which favours the ion-pair dissociation and is also essential in the formation of the passivation layer on graphite. EC is a solid at ambient temperature; however, and even if the salt decreases the melting temperature, EC cannot be used alone at sub-ambient temperatures. In addition, salt solutions in EC are highly viscous and restricts the ion mobility. Therefore the selection of different solvents becomes difficult, e.g. some solvents, such as EC, provide high permittivity but enhance the viscosity while others such as dimethyl carbonate (DMC) or diethyl carbonate (DEC) have both low permittivity and low viscosity. In addition to these electrochemical criteria, attention must also be paid to other

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physicochemical properties such as toxicity, melting point, boiling point, flash point, and flammability. 2.2.2. Separators

Contrary to polymer electrolytes, liquid electrolytes allow a high conductivity level to be obtained at ambient and sub-ambient temperatures and are, therefore, essential for electronic applications. However, even with highly conductive liquid electrolytes, a significant part of the internal resistance of the batteries originates from the electrolyte. A natural solution consists, therefore, in decreasing the electrolyte thickness, a porous separator being used to separate the electrodes. In this case, the conductivity is ensured by the liquid electrolyte which fills the separator porosity. In lithium batteries, the separator must be selected among polymer matrixes that are both chemically and electrochemically stable vs. Li metal or LiC6. Most of the commercial macroporous separators dedicated to lithium batteries are manufactured from apolar polyolefins such polyethylene, polypropylene or polyethylene  polypropylene. Thus, Celgard® is widely used but the apolar nature of the polymer does not favour the pore wetting by the highly polar electrolytes. Abraham and Alamgir [12] reported a resistivity increase by a factor of 5 while we found a factor close to 10 with usual electrolytes [13]. Ooms et al. [14] reported the performances of Solupor® from DSM Solutech, which is made from oriented Ultra high-molecular weight (HMW) polyethylene. They claimed a fast wetting, a high rate capability and a low tortuosity. However, they did not provide comparative impedance data between the liquid electrolytes alone and the Solupor®  liquid electrolyte sets. The apolar nature of the separator is detrimental to the pore wetting and, therefore, to the internal resistance of the battery. However, it prevents the inter-pore matrix from swelling by the electrolyte, thus preserving its mechanical properties. In order to keep the high tensile strength provided by the semi-crystalline inter-pore matrix, while enabling a good pore wetting, Gineste et al. [15] proposed a surface treatment followed by UV irradiation in order to initiate the free-radical photopolymerization of acrylic monomers grafted at the surface of the pores. This attractive technology allows the pore coating by a polar layer that improves the pore wetting, while preserving the mechanical properties of the inter-pore matrix [15,16]. Recently Augustin et al. proposed new separators Quallion®, based on ceramics deposited on a tissue consisting of polymer fibres [17]. In addition to the tensile strength of the separator, battery manufacturers pay attention to the safety aspect provided, through the shut-down effect, by some separators. This effect originates from the melting of the semi-crystalline polymer. Thereafter, the porous volume clogs and the resistivity increases dramatically. According to the melting point of the crystalline phase of the polymer, the shut-down may occur between 110 and 150°C.

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2.3. Gelled polymer electrolyte

An alternative to the above discussed electrolyte consists in using a polymer swollen by the liquid electrolyte. In this case a porous separator is not required. Thirty years ago, Feuillade et al. [18] proposed the use of gelled polymer electrolytes instead of the classical set i.e. liquid electrolyte  porous separator in order to obtain thin batteries as claimed by the title of their French patent “Eléments électrochimiques en couche mince.” Later they reported [19] the use of several gelled polymer electrolytes, based on commercial polymeric matrixes, e.g. polyacrylonitrile (PAN), polyvinyl acetate (PVAc), polyvinylbutyral (PVB), polyvinylpyrrolidone (PVP), polyacrylonitrile and poly(vinylidene fluoride)-hexafluoropropene (VdF–HFP) copolymers as electrolyte of Li/CuS batteries. In order to correlate the solvent/polymer interactions with the formation of gels, they investigated (i) the behaviour of several thermoplastics in diluted propylene carbonate solutions by membrane osmometry and (ii) the mechanical properties of the gels. This technique allows not only the average molecular weight (Mn) to be calculated but also the second virial coefficient A2, to be reached. Unfortunately, it was impossible to correlate A2, obtained from diluted polymer solutions, to the gel formation that occurs at high polymer concentration. Among the polymer matrixes, PAN is unstable vs. metallic lithium and this was shown in particular by following the impedance of symmetric cells lithium/PAN gelled electrolyte/lithium [20–22]. Poly(VdF-co-HFP) also appears unstable [23] vs. metallic lithium. The chemical instability of fluorinated polymers vs. metallic lithium is probably due to halogen–lithium exchange reaction while that of PAN is partially due to the acidic character of the hydrogen located in the α position of the nitrile [22,24]. Among the polymeric matrixes another industrial thermoplastic, polymethylmethacrylate (PMMA), was proposed by Bohnke et al. [25–27], in particular to prepare gelled polymer electrolytes for electrochromic windows. However, the PMMA gel is closer to a very viscous liquid than to a self-supported film. Obviously, very early times, it was proposed to use POE-gelled polymer electrolyte. Surprisingly, despite the high conductivities of POE–lithium salt complexes the use of gelled POE electrolytes does not lead to higher conductivities than those obtained with the previously gelled thermoplastic electrolytes. In addition, due to the strong interaction between poly(oxyethylene) and lithium cation, a decrease in cationic transport may be anticipated [28]. More recently, it was reported that replacement of the fairly acidic hydrogen of PAN by a methyl group, leading to polymethacrylonitrile, results in a clear improvement of the electrochemical stability in reduction [24]. Lastly, a third type of electrolyte intermediary between a gelled electrolyte and a macroporous separator  liquid electrolyte set was more recently described. It consists of a macroporous PVdF separator, the polymeric skeleton of which is swollen by the liquid electrolyte while the porous volume is filled with the liquid electrolyte [29]. The swollen part of the membrane might possibly behave as an electrolyte reservoir.

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PVdF has high thermomechanical properties and fairly high permittivity. Due to an expected low DN, it has no solvating ability vs. lithium salts and cannot be used in polymer–salt complexes. On the other hand, copolymers of PVdF have been used for three decades as gelled polymer electrolytes as well as binders for the electrodes. Before introducing the properties of electrolytes based on the PVdF family, the main characteristics of PVdF must be taken into account. 3. ABOUT THE MORPHOLOGY OF VdF POLYMERS 3.1. VdF homopolymers

PVdF homopolymer (CH2–CF2)n is a semi-crystalline linear thermoplastic, whose crystallinity varies between 35 and 70%, according to the method of synthesis and to the thermal history of the sample. This crystallinity makes it possible to have a good mechanical resistance, and to use this polymer in applications requiring creep strength, resistance to tiredness and abrasion. In addition, these properties are also conditioned by the average molecular weights, the molecular weight distribution, and the head-to-head content. The PVdF Kynar 301F from Totalfina is synthesized by emulsion. For instance, PVdF samples with high molecular weight, i.e. Mn  100,000 g/mol, and low head-to-head defects ( 7%) make it possible to limit the swelling ratio by the liquid electrolyte, thus preserving the mechanical integrity of the swollen membrane. However, the membrane elaboration by dissolution in hot acetone remains possible and does not pose the problems of implementation noticed with some PVdF grades, which have both a higher molecular weight and weaker defect content. The thermal degradation of PVdF starts from 390°C, producing HF as the by-product [30]. In addition, it is resistant to most of the inorganic acids, halogens, and oxidants, even at high temperature, which justifies its use in the chemical processing equipment. However, it is sensitive to bases which induce a dehydrofluoration and generate double bonds. As compared to its counterpart polytetrafluorethylene Teflon® (CF2–CF2)n, due to its methylene group, PVdF presents a satisfactory compromise between chemical resistance and simplicity of implementation. PVdF dissolution can thus be performed in hot solvents, after fusion of its crystallites. Among its best solvents are polar solvents, such as acetone, dimethylsulfoxide (DMSO) and dimethylformamide (DMF) i.e. which have CO or SO groups. On the other hand, benzene, hexane or alcohols are non-solvents of PVdF. Several types of crystalline phases exist for PVdF, which is rare for a synthetic polymer. Most popular is the α phase (form II) of the type trans-gauche TGTG, hydrogen and fluorine alternating in a regular way on both sides of the chain [31–33]; conformation is of helix type. The α form is thermodynamically the most stable form. Both polymeric chains are packed in such a way that the molecular dipoles are in antiparallel position and cancel each other out. PVdF

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can in addition, comprise phases with a dipole moment and have ferro, piezo and pyro electric properties. The β phase (form I) and the γ phase (form III) are the polar phases most commonly met. For all these phases, the chains are packed in the unit cell so that the dipoles associated with the individual molecules are parallel resulting in a dipole moment for the crystal; the dipoles are perpendicular to the chain axis. The β phase leads to the best ferro, piezo and pyro electric properties. The β phase is made up of trans–trans TTTT sequences, the fluorine atoms are always on the same side of the chain. This phase has a zigzag planar conformation [34]. The γ phase is made up of TTTGTTTG sequences, while there is a controversy about the unit cell: monoclinic according to Takahashi and Tokodoro [35] or orthorhombic according to Weinhold et al. [36]. There is, in addition, a δ phase with a dipole moment that results from the α phase [37]. This δ phase (also named IV or IIp form) that results from the alignment of the parallel dipoles is only obtained by a corona discharge under a strong polarizing field (1 MV/cm). Such a field is by far much higher than that existing in a lithium-ion battery and, therefore, the δ phase should not be present in PVdF-based polymer electrolytes. Obtaining a crystalline phase depends on several factors. First, it depends on the structure of the polymer, i.e. molecular weight [38] and on the head-tohead defect content [39]. It depends also on the elaboration process. Thus, cooling at T 150°C, under high pressure, a PVdF melt leads to γ phase [40], while a quenching under high pressure leads to β phase [41]. The nature of the PVdF phase obtained by film casting depends on the solvent used and on its evaporation rate [42,43]. Film casting may thus result in γ, β or α phases according to whether dimethylacetamide (DMAc) hexamethylphosphoramide (HMPA) and acetone are, respectively, used as solvents. The density of the crystalline part of PVdF slightly depends to some extent on the nature of the crystalline phase (1.92–1.97) while the density of the amorphous phase (1.68) is clearly lower. 3.2. VdF copolymers

Vinylidene fluoride (VdF) can be polymerized with a variety of comonomers leading to statistical or random copolymers [44]:

H2C CF2 x CF2

F C CF3

p

(P(HFP-co-VdF))

However, using radical telomerization or cotelomerization from iodine-terminated perfluoropolyether Gelin and Ameduri [45] succeeded in preparing block copolymers consisting of a perfluoropolyether block and either of a PVdF

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block or of a poly(VdF-co-HFP) block: F F3C

CF2

CF2 O C

O CF2

CF2

CF3

CF2

I +

n

C=C

+

C=C F

F

F

F

F

F

F

CF3

(VdF)

(HFP)

(PPFE)

O O F F3C

CF2

CF2 O C CF3

F CF2

O CF2 n

CF2

H2C CF2

CF2 x

C CF3

p

These new polymers are interesting but they remain at the laboratory scale. 3.2.1. Crystallinity

Crystallinity and glass transition temperatures were found to depend both on the block length and on the HFP content. Among the various comonomers of VdF, hexafluoropropene does not give homopolymers but leads to a variety of industrial VdF copolymers manufactured, in particular by Totalfina® and Solvay®. The incorporation of HFP decreases the crystallinity ratio and the melting temperature with regard to PVdF and increases the glass transition temperature. From wide-angle diffraction measurements performed on a Kynarflex 2801® sample, which is a poly(VdF-co-HFP) with 12 wt% of HFP, Abbrent et al. [46] show that the PVdF units are capable of inducing a partial crystallization and estimate the crystallinity ratio at 31%. The nature and the positions of the three strong crystalline peaks allow them to be associated to an orthorhombic unit cell typical of the α phase (form II) of the homopolymer PVdF, namely a  4.96, b  9.64 and c  4.62 Å. 4. NON-POROUS-GELLED POLYMER ELECTROLYTES BASED ON PVdF COPOLYMERS VdF copolymers have been extensively used as a component of electrolyte and electrodes of lithium-ion batteries. 4.1. Gelled polymer electrolyte based on VdF homo- and copolymers

If PVdF-HFP copolymers are unstable vs. lithium they may be used as gelled polymer electrolytes in contact with lithium graphite negative electrode. The main advantages of the fluorinated copolymer concern its (i) availability, (ii) stability in oxidation and (iii) presence as binder, at least in the positive electrode. In order to

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improve the mechanical properties of gelled VdF–HFP copolymers, Feuillade and Perche [19] performed, their crosslinking by a nucleophilic attack, using a diamine, namely dicinnamylidene hexanediamine in ethyl and methyl ketone. However, PVdF homopolymer does not lead, by film casting, to dense membranes free of porosity. 4.1.1. Additives used with VdF copolymer

Prior to their extrusion or injection or to other processing, industrial polymers are formulated and incorporate a variety of additives. These additives must be used temporarily as the aid-process plasticizers or must remain in the final material. Thus the polymer may consist of aid-use plasticizers, antioxidants, fire retardants and pigments. Some additives, such as starch or flour, must also be used to partially replace the polymer to reduce the cost. Polymer reinforcements require the use of fillers, e.g. carbon black or fibres, such as e.g. glass or carbon fibres. In battery applications, battery-grade silica are often used. One of the processes to obtain gelled PVdF copolymer electrolytes is to use an external aid-process plasticizer, namely dibutyl phthalate (DBP). Owing to the external plasticizer separator electrodes can be bonded together. Then the DBP plasticizer is extracted by using an appropriate solvent, which simultaneously generates some porosity [47]. Irrespective of its advantages the process requires an additional step to extract the aid-use plasticizer. To improve their mechanical properties and to some extent their conductivity, fillers such as silica fume, TiO2 or Al2O3 , have often been incorporated into the gelled polymer electrolyte. However, few papers report data on mechanical analyses of these very complex polymer electrolytes that consist of polymer, fillers, salt and solvent mixtures. Thus, Kim et al. [48] reported that addition of rutile TiO2 nanoparticles up to 70 wt% vs. polymer of particles led to self-supported gelled polymer electrolyte film, but this qualitative description is not sustained by mechanical characterizations. They pointed out that an addition of TiO2 increases the conductivity simultaneously with the electrolyte uptake. As the liquid electrolyte uptake greatly decreases the mechanical strength of PVdF–HFP gelled electrolytes, Wang et al. [49] reported the use of a microporous polyolefin to improve the mechanical properties. It is a unfortunate because these gelled polymer electrolytes are supposed to replace the microporous separator used in commercial batteries. 4.1.2. Blends based on VdF copolymers

The scientific electrochemistry community frequently uses mixtures of (i) inorganic materials in electrodes, (ii) solvents and, sometimes, (iii) salts in the electrolyte. Obviously, with great imagination it has extended this approach to polymers and various blends have been proposed. But, in fact, due to thermodynamic reasons pertaining to the entropy of mixing, compatibility between

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polymers is very rare. Incompatible blends can be used, depending on the size of the phase-separation domains but the situation may evolve with time, in particular, if the system is swollen by solvents and submitted to electrical fields. Fluorinated copolymers have often been blended with other polymers, but few polymers are compatible with fluorinated homopolymers and copolymers. An interesting paper has recently been published by Wang and Tang [50], who blended VdF–HFP copolymers with PVP. They established, by DSC measurements, the compatibility between PVP and VdF–HFP copolymers. Although PVdF has previously been reported as compatible with PVP [51], the compatibility between PVP and the copolymer VdF–HFP has not been established. Oh and Kim [52] reported blends obtained by mixing, in the presence of silica fillers, VdF–HFP copolymer with a methyl methacrylate-vinyl acetate random copolymer. If PVdF and PMMA homopolymers are undisputedly compatible, there is no evidence of any compatibility between the previous copolymers. This general lack of precisions about the blends contrasts with the careful approach of polymer scientists who, despite the fact that the compatibility between PVdF and PMMA was unambiguously established, investigated in depth the interactions between PVdF and isotactic, syndiotactic and atactic PMMA before concluding that (i) the three forms provide compatibility and (ii) PVdF interaction with isotactic PMMA is stronger than that with syndiotactic PMMA [53]. Another approach reported by Cheng et al. [54] dealt with a semi-IPN (Inter Penetrated Network) obtained by dissolving VdF–HFP copolymer with polyethylene glycol (PEG) and a macromonomer polyethyleneglycol dimethacrylate (PEGDMA), in acetone/ethanol mixture the functionality of which (f  4) allows a three-dimensional network to be obtained by free-radical polymerization. CH3 O

CH2 = C

O

C=O O H2C

+

H2C O

H2C CF2

x

CF2

F C CF3

n

C=O CH2 = C CH3

O

O

O

p O

O

O

(P(HFP-co-VdF)) (Semi -IPN)

(PEGDMA)

The solvent evaporation generates some porosity. The resulting membrane is swollen by 1M LiPF6 solution in EC–DMC and the liquid uptake decreases with the content of polyether network. The authors reported an improvement in the mechanical properties, the tensile modulus reaching 78 MPa for a semi-IPN free of PEG. However, these mechanical characterizations were performed on non-swollen membranes. The conductivities of swollen membranes laboriously reach 1.5 mS at 298 K.

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After the first announcement of macroporous PVdF prepared by phase inversion most of the papers related to gelled fluorine polymer electrolytes dealt with this process. They often selected PVdF copolymers rather than homopolymers as the decrease in crystallinity induced by the copolymerization both widened the range of solvents and decreased the temperature of dissolution. 5. MACROPOROUS PVdF MEMBRANES The word microporous has been used for a very long time to refer to porous polymeric membranes, even when the average porous diameter is close to 0.5 μm. According to IUPAC rules, it is recommended to refer to them as macroporous membranes. 5.1. Macroporous membrane elaboration

Although film casting from a solution of PVdF generates some porosity, obtaining macroporous PVdF membranes is mainly carried out according to three methods. The first method deals with a phase-separation process from a dramatically cooled PVdF solution. The other processes can be gathered as inversion processes and use ternary mixtures consisting of PVdF/solvent/nonsolvent. 5.1.1. Description of phase-inversion processes

Two phase-inversion processes lead to macroporous membranes. In the first one, the solvent used is selected from among volatile solvents while the non-solvent is selected from among the solvents having high boiling points. After casting the ternary mixture, the progressive evaporation of the solvent increases the concentration in the non-solvent and generates the porosity in the membrane (Fig. 2). Thus, by evaporation of a mixture of acetone/ethanol i.e. solvent/non-solvent polymer

gel

miscibility lacuna

solution

solvent

Non-solvent

Fig. 2. Principle of phase inversion: ternary polymer/solvent/non-solvent diagram. The route to obtain membranes is shown by the dotted lines.

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from a swollen semi-IPN, Cheng et al. [54] reported obtaining microporous membranes. According to their SEM image they may be ranked among the macroporous membranes. In the second one, a viscous solution of PVdF is cast onto a support and is then immersed in a coagulation bath consisting of a non-solvent of PVdF. The solution becomes unstable and generates the porosity. This method has been used, for instance [55], to shape PVdF into a macroporous membrane for filtration purpose. They immersed an acetone solution of PVdF in an 80:20 acetone/water mixture behaving as a non-solvent. It is obviously possible to vary solvent, the composition of the coagulation bath and the temperature [56]. In fact, the principle of this process is to generate a phase separation leading to a polymer- and solvent-rich phase. 5.1.2. Phase-inversion mechanism

When, by immersion or selective evaporation, the solution becomes unstable, phase separation occurs with the appearance of liquid microdroplets, polymer-poor, dispersed within a polymer-rich phase. In these droplets, the polymer is gradually rejected in the periphery. When a certain viscosity threshold is reached within the phase richest in polymer, the structure coagulates into a gel. The polymeric chains get organized, which then results in the formation of balls of polymeric chains around liquid globules. These balls coalesce and interpenetrate; this organization can also result, in the case of PVdF, in a partial crystallization. The solvents and non-solvents then evaporate and create small diameter pores. 5.1.3. Alternative processes

Among the alternatives to the previous elaboration associated with a solvent and a non-solvent of the polymer, most of the film-casting methods by solvent evaporation generate porosity, but often this porosity was not reported and its extent as its reproducibility are questionable. However, Michot et al. [57] reported that phase separation occurs during the casting process. This was performed from solutions – in N-methylpyrrolidone (NMP), methyl, ethyl, ketone (MEK), THF, acetone – of VdF polymers  liquid electrolyte (LiBF4/PC) mixtures. They found that phase separation increases with the crystallinity and is more significant with PVdF homopolymers than with VdF–HFP copolymers. A very attractive alternative may consist of electrospinning a PVdF solution, which results in the formation of nanofibres that can be deposited at a rate of several metres per second. The process has been known for more than 70 years. Flat ribbons and various other shapes can be obtained from solutions of different polymers such as polyhydroxyethyl methacrylate, polystyrene, polyetherimide or PVdF [58]. In addition to the characteristics of the electrospinning jet, special attention must be paid to those of the solution, i.e. its viscosity and its

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surface tension coefficient, higher viscosity favouring the formation of fibres free of beads [59]. Recently, Choi et al. [60] have reported that electrospinning, on a stainless-steel plate, PVdF Kynar Flex 761® (Mn  5.5.105 g/mol) from an acetone/N,N-dimethylacetamide mixture leads to a fibrous membrane, with a thickness of 30 μm. Solvent evaporation is very rapid, some milliseconds, because the phase separation occurs in the polymer jet. The resulting fibres exhibit a smooth surface, their average diameter is 250 nm and the average pore size is about 0.65 μm. Wide-angle X-ray scattering shows that the nanofibres have typical characteristics of the α type (form II). The conductivity of a membrane soaked in 1 M LiPF6 in ethylene carbonate/dimethyl carbonate mixture reaches 1.7 mS/cm at 0°C. 5.2. Macroporous membrane specifications

The conditions necessary to obtain an appropriate macroporous membrane are: ●

●

●

●

Use, of non-toxic solvents at low temperature (T 100°C), with regard to industrial requirements and environmental friendly processes Obtaining a sufficiently viscous initial solution allowing the film to keep a mechanical integrity during the immersion step in the coagulation bath A porosity, uniform and opened, exceeding at least 40%, as it governs the conductivity performances Neither skin on the surface of the membrane (blocking of ionic transport), nor on the macroscopic channels

The macroporous PVdF microstructure is strongly influenced by the operating conditions, namely, solvent, composition, and coagulation temperature. Boudin [61] determined the appropriate parameters to obtain, from Kynar® 301F, the required microstructure: ●

●

●

Use of a 17% PVdF solution in hot acetone. This solution is close to the demixtion conditions and must be miscible with the non-solvent, in order to balance the flows between solvent and non-solvent and to obtain a uniform microstructure. Coagulation in an ethanol bath, a sufficiently bad solvent, to generate a high porosity, but not enough to lead to a coarse porosity with macrochannels. Wiping and soft drying in a drying oven at 60°C.

After membrane processing, complete evaporation of the coagulation bath showed that partial dissolution of the shortest macromolecular chains does not occur during the phase-inversion process. PVdF used is therefore fully recovered [62,63].

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5.3. Characterizations and swelling behaviour of the macroporous membrane 5.3.1. Macroporous PVdF free of solvents 5.3.1.1. Porosity The comparison between macroporous membranes based on fluorinated polymers is not easy. Indeed often the porosity characterization is more qualitative than quantitative and when the characterizations are quantitative they do not use the same polymer PVdF or VdF–HFP copolymers. Lastly, many papers deal with very complex mixtures involving inorganic additives such as silica or other fillers or organic ones such as high-molecular weight plasticizers, etc. Scanning electron microscopy is often used and provides useful, but mainly qualitative, information. Most of the papers only report the electrolyte uptake. This information is necessary as the electrolyte uptake governs the conductivity, but does not exactly reflect the porous volume as the inter-pore polymeric matrix, swollen by the electrolyte, also contributes to this uptake. Another approach consists of soaking the macroporous membrane in a nonsolvent of the polymer [64], i.e. n-butanol. This method is attractive but the incompatibility between the polymer and n-butanol may induce a poor wetting of the pores, thereby lowering their filling by this non-solvent. Boudin et al. [65] estimated the average pore size in PVdF Kynar® 301F at about 0.5 μm, using a Coulter porosimeter. The mean pore diameter and pore size distribution were also measured [60] using a capillary flow porometer. Lastly, characterization of membrane porosity was also performed, using BET [64]. Michot et al. [57] compared the porosity obtained by film casting using different solvents and different polymer grades. It defined the porosity with regard to the volume occupation of the polymer as to the ratio of the apparent density of the porous membrane to the density of the related polymer. From these values it clearly appears that the extent of porous volume depends on the casting solvent and on the fluorinated polymer. Thus, film casting from NMP solution of PVdF–HFP Kynar Flex 2801® results in 2% porosity while using MEK instead of NMP enhances up to 53%. The porosity also depends on whether a homopolymer or a copolymer is used to cast the membrane. Thus Kynar Flex 2801® and Kynar 301F®, both cast in MEK, lead, respectively, to 9 and 47% porosity. This clearly shows that the higher the crystallinity, the higher will be the porosity. Although very convenient, the previous method based on the volume occupation of the polymer neither allows the pore size distribution nor the discrimination between open and close porosity to be known. Another method consists of using mercury porosimetry. Saunier et al. [66] found porosities close to 72% with a very homogeneous pore diameter close to 0.64 μm, using PVdF Kynar 301F®. As the return curve occurs with a negligible hysteresis, in agreement with a complete shrinkage of the mercury, they assume that this technique does not affect the structure of the membrane.

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5.3.1.2. Structure of macroporous PVdF Comparative study performed by X-ray scattering and infrared on hot-pressed PVdF Kynar 301F® powder and macroporous PVdF shows that both samples exhibit a high α phase ratio [67], the amount of β and γ phases being slightly higher in the hot-pressed sample. The crystallinity ratio (by DSC) is close to 42% for the hot-pressed sample and reaches up to 60% in the macroporous PVdF [68]. The melting point of the macroporous PVdF (141°C), lower than that of the hot-pressed PVdF (150°C), indicates a thinner microstructure. Michot et al. [57] reported for the same PVdF grade a fusion occurring between 145 and 170°C. Safety, thermomechanical and electrochemical performances depend on the liquid electrolyte uptake by the membrane. Special attention must, therefore, be paid to (i) the thermodynamic and kinetic aspects of the swelling, (ii) the modifications (Tg, crystallinity, storage modulus) and degradations (cracking, macrovoids, etc.) and (iii) the safety aspects in relation with the shut-down effect. 5.3.2. Swelling behaviour: Thermodynamical aspects

The electrolyte uptake by a dense or macroporous polymer matrix has a direct influence both on the thermomechanical properties and on the conductivity of the electrolyte  polymer set. Swelling of a polymer by a multicomponent mixture, as the liquid electrolyte, is very complex and requires sound knowledge in thermodynamics of polymer solutions, especially because few (or no) thermodynamical approaches were reported about interactions between polymers and such exotic solvents. Saunier et al. [68] selected a step-by-step approach to study the interactions between PVdF Kynar 301F® and the liquid electrolyte and they first investigated binary mixtures between PVdF and a single solvent. Thus, swelling ratios were measured and interaction parameters were calculated for the aprotic solvents usual in lithium-ion batteries, namely diethyl carbonate, dimethyl carbonate, propylene and ethylene cyclic carbonates. The calculation of the polymer/solvent interaction, χsp, defined in the Flory–Huggins theory [69,70] allows the estimation of the swelling ability of a polymer by a given solvent. The lower the χsp the better the solvent is. The lowest interaction parameter was found for DMC. This clearly indicates that DMC is the best solvent for PVdF while DEC is the worst (Fig. 3). A high affinity PVdF/DMC may lead to a possible dissolution of PVdF around 80°C. The size of the solvent molecule is not the only relevant factor and, in fact, its shape, i.e. cyclic or acyclic, permittivity, viscosity, may also affect the swelling behaviour. A surprising overswelling was reported after soaking PVdF Kynar 301F® dense membranes in binary or ternary solvent mixtures. This means that, at equilibrium, the swelling ratio is higher than the sum of the mass fraction swelling ratios. This overswelling can be explained by

5.3.2.1. Swelling ratios in solvent mixture

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0.2 DMC

Molar fraction in gel

DEC EC

0.15

PC 0.1

0.05

0

10

20

30

40 50 60 Temperature in °C

70

80

90

Fig. 3. Solvent molar fraction in gels at different temperatures for pure solvents.

a bad affinity between the solvents [71,72]. As a result both solvents interact preferentially with PVdF and the swelling ratio increases. In a similar way as for polymer/solvent interaction, solvent/solvent interaction can be quantified through the solvent/solvent interaction parameter χsisj, where i and j represent respectively, to solvent i and solvent j. It may be emphasized that positive values of χsisj correspond to a poor affinity between i and j solvents. Few papers deal with complex systems containing more than two components [73,74]. However, modelling, based on the Flory–Huggins theory, of the swelling equilibrium of PVdF by cyclic and acyclic carbonate mixtures has been successful. This is surprising as this theory is above all well-adapted to apolar solvents and elastomers. Thus, starting from the experimental interaction parameters of PVdF with each solvent, a good agreement between the experimental and calculated values of polymer/solvent and solvent/solvent interaction parameters was observed, except for χEC/DEC, as the calculated value provides too high interaction parameters. As PVdF is soaked in a multicomponent liquid electrolyte and not in a single solvent, selectivity of PVdF dense membranes vs. the electrolyte components is predictable. Fig. 4 plots the swelling volume ratio of PVdF Kynar 301F® from binary equivolume mixtures, i.e. DEC/DMC, EC/DMC vs. the swelling temperature. From this plot it is clear that the swelling selectivity in binary mixtures decreases when the temperature of swelling mixtures increases. The solvent for which the interaction parameter, χsp, is the lowest is in the highest proportion in the PVdF membrane. This selectivity depends on the absolute value of the difference between the χsp values, relative to PVdF and each of these solvents. Thus, the selectivity is higher for the DEC/DMC mixture. 5.3.2.2. Swelling selectivity

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0.9 0.85

Volume ratio

0.8 0.75 0.7 0.65 0.6

VDEC/VDMC VEC/VDMC

0.55 0.5 10

20

30

40

50 T (°C)

60

70

80

90

Fig. 4. Polymer selectivity/Mixture composition (volume ratios) inside the PVdF gel for swelling in equivolume binary solvent mixtures at different temperatures.

These data were obtained from equivolume mixtures but swelling selectivity also depends on the swelling solution composition. The complexity still increases by the addition of a third solvent. However, and despite a lesser reproducibility, for the ternary mixture EC/DMC/DEC, an increase in DMC and EC concentrations compared to DEC and a slight decrease in the EC concentration compared to the DMC have been shown. Despite the information about the swelling ability of the usual solvents of lithium batteries, most of the processes incorporate the salt and the solvents at the same time and the salt may affect the swelling equilibrium, i.e. the swelling ratios and the swelling selectivity. Infrared and Raman spectroscopic studies, on the one hand [75,76], and NMR experiments based on 19F, on the other, show a bad affinity between LiPF6 and PVdF [77] which results in a much lower salt concentration in the electrolyte trapped in the membrane than in the swelling solution with a decrease by a factor ranging between 4 and 10. The addition of salt in a single solvent or in solvent mixtures modifies the swelling ratios and the swelling selectivity of PVdF membranes with respect to unsalted solvents [67]. Thus, swelling ratios in LiPF6/EC solutions decrease linearly with the salt concentration in EC. Not only are the swelling ratios but also the swelling selectivity greatly affected on addition of LiPF6 in solvent mixtures. The salt effect is a very complex problem. Indeed at least three interaction parameters are required to explain the swelling behaviour in solvent mixtures free of salt. Addition of salt leads to additional interactions, i.e. polymer/salt and

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salt/solvents and, due to its ionic character, it might also modify the solvent/solvent and solvent/polymer interaction parameters. Not only the salt but also silica fillers may modify the liquid electrolyte uptake as reported by Caillon-Caravanier et al. [78]. They reported that addition of SiO2 in VdF–HFP copolymer gelled by electrolytes, comprising LiTFSI and valerolactone, increases the membrane porosity, the electrolyte uptake and, therefore, the conductivity 5.3.3. Swelling behaviour: Kinetic aspects

If the knowledge on the swelling equilibrium allows optimization of the electrolyte uptake, and therefore of the conductivity, both for dense and for macroporous membranes, the kinetics is also essential for the battery manufacturing. Indeed, in order to overcome the problems related to the moisture sensitivity of lithium salts, the Bellcore group cleverly proposed [79] to delay the electrolyte incorporation using a PVdF–HFP copolymer temporarily plasticized. The aid-process plasticizer is then removed to be replaced by the liquid electrolyte. Nowadays it is a common practice to perform a late incorporation of the electrolyte on an electrolyte-free battery, but the swelling has to be fast. As both dense and macroporous PVdF membranes consist of dense polymeric parts the swelling kinetics of dense PVdF membranes should provide useful information to the battery manufacturers. The nature of the polymer matrix, i.e. glassy, rubbery or semicrystalline on the one hand and the penetrating molecules on the other make the swelling kinetics often very complex. Instead of being Fickian, an anomalous behaviour may be observed with glassy [80,81] cross-linked rubbery [71,72] or semicrystalline polymers [82]: the sorption curves exhibit a sigmoidal S shape. This behaviour is illustrated in Fig. 5, which plots several sigmoidal sorption curves of DMC in dense PVdF Kynar 301F® at different swelling temperatures. As can be seen, the first part of the sorption curve is linear, i.e. Fickian behaviour and then becomes anomalous (Fig. 5). The Fickian part of the sorption curves allows the diffusion coefficient, D, of the solvent to be calculated. It must be emphasized that D can be correlated to the polymer/solvent interaction parameter. At 40°C, DMC, the smallest molecule has the highest diffusion coefficient i.e. D  2  109 cm2/s while DEC, the largest molecule has the lowest coefficient, i.e. D  4  1011 cm2/s. As for the cyclic carbonates, such as EC or PC, the D values are close and intermediate between those of DEC and DMC. The diffusion coefficients of binary and ternary carbonate mixtures are intermediate between those calculated from the single solvents. The salt modifies the shape of the sorption curves in agreement with a twostage sorption [83,84]. A two-stage sorption was not only observed in glassy polymers but also in the rubbery polymers [85]. The first stage is Fickian while the second one is related to polymer chain relaxations. The diffusion coefficients

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Fig. 5. Sigmoidal sorption curves of DMC in dense PVdF slabs (mass uptake is the weight of solvent in gel divided by dry polymer weight and is plotted vs. square root of time) at different swelling temperatures. Curves are normalized to the sample thickness e by dividing the square root of time by the sample thickness.

can be calculated from the Fickian part. At 50°C, the D value is close to 7.1  108 cm2/s for EC/LiPF6, i.e. two orders of magnitude higher than those obtained by swelling the sample from solutions free of salt. In macroporous PVdF membranes, the situation is still more complex as both porous volume and PVdF support are swollen by the electrolyte. The kinetic study allowed the equilibrium times to be calculated for a ternary solvent mixture EC/DMC/DEC in a macroporous membrane PVdF Kynar 301F® made of filaments, the diameter being 1 μm. This equilibrium time ranges from 2.4 ms at 40°C to 0.8 ms at 60°C, that means that, irrespective of the swelling temperature, the filling of a macroporous PVdF membrane by a liquid electrolyte is very rapid. 5.3.4. Swelling behaviour: Structural modifications, degradations and safety aspects

In addition, swelling can also lead to a macroscopic or microscopic degradation, and cracking has already been reported [86] for other kinds of polymer after their swelling. 5.3.4.1. Structural modifications

Thermodynamic studies of the thermoreversible and thermal gelation were reported earlier by Guenet and Mc Kenna [87]. More recently, Mal and Nandi [88] reported a thermodynamic study on the thermoreversible PVdF gels in several solvents, e.g. acetophenone and found (i) a regular decrease in the gel fusion temperature with the content of solvent and (ii) a positive 5.3.4.1.1 Crystalline phase

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deviation from linearity in the plots of enthalpy of gel fusion vs. weight fraction of PVdF. They assumed that the enthalpy of gel melting consists of the sum of four contributions, i.e. solvent, PVdF crystal, random mixing PVdF/acetophenone and a forth term related to the formation of oriented structure. They attributed the positive deviation to the last two contributions. Despite this sound study being performed using single solvent, obviously free of salt, and not electrolyte based on solvent mixtures, it provides relevant information. Jiang et al. [89] prepared gelled polymer electrolytes from both PVdF Kynar 741® and PVdF–HFP Kynar 2822® using (CF3SO2)2NLi and LiPF6 salts in a 1:1 solvent mixture EC/PC. From X-ray diffractions they conclude that plasticizing both fluorinated polymers by the previous solution induces disorder into the polymer structure and results in a clear decrease in crystallinity. The decrease in the melting point of swollen PVdF depends, according to Tazaki et al. [90], not only on the solvent uptake but also on the nature of the solvent. Thus, γ-butyrolactone PVdF gels melt roughly 55°C lower than octanone ones. Saunier et al. [91] reported PVdF Kynar 301F® crystalline-phase modifications that affect (i) the melting point, (ii) the melting enthalpy and (iii) the crystalline phase. The melting point depression may be related to a decrease in the crystalline lamella thickness and/or to the presence of low-molecular-weight diluents. Its amplitude strongly depends on the polymer/solvent affinity, the best solvents, DMC and EC, leading to the more pronounced depression. The decrease in melting enthalpy may be attributed both to a decrease in the crystallinity ratio and to an easier melting in the presence of a solvent. From wideangle X-ray scattering, Abbrent et al. [46] showed that the incorporation of solvent cannot alone significantly decrease the crystallinity ratio. However, the membrane processing used by Saunier and Abbrent was different, as the film casting performed by Abbrent, involving solvent evaporation, generates porosity. The composition of the crystalline phase of PVdF Kynar 301F® is slightly modified by the swelling. Indeed, X-ray spectra shows that the β phase increases with the solvent uptake. This content in β phase is maintained after the solvent evaporation. In Fig. 6, the melting temperature is plotted vs. the volume fraction. 5.3.4.1.2. Amorphous phase Reversible modifications are related to a lowering of the chain–chain cohesion that results in a decrease in the glass transition temperature Tg. Using both modulated differential scanning calorimetry (DSC) and dynamic mechanical analysis (DMA), Saunier et al. [91] showed, on PVdF Kynar 301F® an almost linear decrease of Tg with the volumic solvent fraction. The plasticizing effect decreases according to DEC  DMC  EC. EC even seems to behave as an antiplasticizer but this behaviour is probably related to its crystallization in the gel. Irreversible modifications as chain disentanglements were demonstrated by rheometric measurements.

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Fig. 6. Variation of the melting temperature with the volume fraction of solvent for different solvents.

Optical microscopy allowed parallel cracks to be characterized on PVdF Kynar 301F® samples swollen at a high solvent ratio. This results in sample delamination. Open porosity is not modified in non-macroscopically degraded samples. Scanning electron microscopy (SEM) did not allow detecting microcracks while macroscopic cracking suggests their existence.

5.3.4.2. Macroscopic and microscopic degradations

The safety aspects deal with the thermomechanical stability of the gelled polymer electrolyte and the possible shut-down effect. Gelling dense PVdF Kynar 301F® membrane results in a loss of storage modulus by about a factor 10 at ambient temperature for a 20 vol% solvent uptake [91]. As the uptake of dense membrane is obviously limited, the mechanical property loss is also limited. Despite the polymer electrolyte being self-supportive, the situation is more critical for macroporous PVdF electrolytes, but an accurate mechanical characterization is not easy. From a study undertaken jointly by DSC, impedance, and SEM, shut-down effects [68] have been characterized for PVdF Kynar 301F® samples at around 135°C when the porosity is close to 70%. Decreasing the porosity to 47% results in an enhancement of the shut-down temperature up to 150°C while, under the same conditions, for Celgard® it is close to 145°C (Fig. 7). More recently, Liu et al. [92] reported a shut-down effect starting at 90°C for a composite gel electrolyte consisting of polyethylene and a VdF–HFP copolymer. However in this case, polyethylene provides a shut-down effect but it 5.3.4.3. Safety aspects

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300

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PVDF porosity 47% Celgard PVDF porosity 70%

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Fig. 7. Comparison of the shut-down temperature of Celgard and macroporous PVdF.

also decreases significantly the ionic conductivity which laboriously reaches 0.2 mS/cm at ambient temperature. 6. ELECTROCHEMICAL BEHAVIOUR OF VdF-BASED SEPARATORS 6.1. Conductivity of macroporous PVdF filled by liquid electrolyte

The technology of lithium-ion battery manufacturing allows electrolytefree storage. This avoids the detrimental self-discharge during the battery storage. The liquid electrolyte is incorporated before the first utilization of the battery. The macroporous membrane may be considered to some extent as a composite material, consisting of the liquid electrolyte solution and of the interpore semi-crystalline polymeric matrix. However, according to whether there are more or fewer affinities between the whole electrolyte or some of its components, and the polymer matrix, two cases must be considered. Case 1: When the electrolyte has a poor affinity with the polymer matrix, the pore wetting is non-optimal and the polymer matrix is not swollen by the electrolyte. A good example is the commercially available and well-known Celgard®, which consists of semi-crystalline apolar polyolefins, i.e. polyethylene/polypropylene. As a result, the conductivity exclusively depends on the liquid electrolyte filling the pores. Case 2: At least one of the electrolyte components has a good affinity for the polymer matrix. Despite its crystallinity the filling of the porous volume by the liquid electrolyte induces a swelling of the polymer matrix, weakening its mechanical properties but improving the filling of the pores. In this case, the

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conductivity results a priori from two contributions: that of the liquid electrolyte filling the porosity and that of the gel phase. 6.1.1. Porosity impact on the conductivity

The conductivity due to the contribution of the liquid electrolyte filling the pores depends on two main factors, i.e. the separator geometry and its affinity with the liquid electrolyte. The membrane geometry, i.e. the pore size distribution, the porosity shape and its percentage is thus an important factor. The key questions are the following: Is the porosity fully opened? What is its shape: finger or sponge? Does the membrane present some macrovoids? Does a non-porous skin exist? All these questions will have a great impact on the ionic conductivity and on the electrochemical efficiency. The insertion of a porous membrane will indeed first introduce a decrease in the electrode surface in contact with the liquid electrolyte and thus will introduce an increase in the system resistivity. For example, a 70% porosity results in a decrease in the electrode effective surface that contributes roughly by a factor 1.4 to the resistivity increase. In addition to this geometrical factor, the tortuosity greatly affects the conductivity of the set liquid electrolyte  macroporous separator. The tortuosity depends on (i) the pore size distribution, (ii) the porous volume and (iii) the porosity shape. An effective way for the conduction to be defined is only by the inter-connected network of pores, this can be increased by a ratio, named tortuosity, compared to the geometrical inter-electrodes distance. The more straight its way can be, the lower the conductivity decrease is, tortuosity being thus close to one. It corresponds roughly to the additional distance that ions must travel to move from the negative to the positive electrode. High porous volume and large pores will reduce this tortuosity, but too large macrovoids are also to be avoided because they may allow dendritic growth to occur more easily, therefore shortening the battery lifespan by creating short circuits. This obviously may occur when metallic lithium is used as a negative electrode. In addition, graphite grains may migrate through too large pores. To obtain good conductivity performances, the porosity of the separator must be higher than 40%, it must be uniform and open, and free of non-porous skin. 6.1.2. Inter-porous gel impact on the conductivity

The contribution of the gel phase to conduction mechanisms in macroporous membranes is less obvious than that of the liquid electrolyte filling the porosity. In gel phase, ionic transport depends on the ionic interactions among ions, solvents and polymer: the more the polymer is swelled, the more enhanced is the ionic conductivity in the gel phase. Kataoka et al. [93], from PFGNMR, assume that the conduction mechanism in macroporous PVdF membrane, obtained by phase inversion, is controlled not by the liquid electrolyte filling the porosity, but by the polymer-rich region swollen by the electrolyte. Data on dense PVdF Kynar 301F® membranes, prepared by hot pressing to avoid the formation of porosity,

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show unambiguously that their conductivities after swelling by usual liquid electrolyte are very low. Thus, the best conductivity obtained, 6  105 S/cm at 60°C [91], is about two orders of magnitude lower than that measured using macroporous PVdF membranes. These low values might be related to (i) the poor solvent uptake and (ii) the poor salt incorporation into the gel phase of PVdF macroporous membrane, as the thermodynamical study of PVdF swelling (see Section 5.3.2) have shown that the gel phase contains only 0.2 M salt in the electrolyte. To improve the affinity and thus to increase swelling and ionic conductivity in the gel fraction, we have to aim for fewer crystalline polymers, different solvents or modified separator (copolymer blends with polymer which has a higher affinity with the liquid phase). However, in addition to enhancing conductivity in the gel phase, the interaction between polymer and the electrolyte has certainly a non-negligible impact on the conductivity of the liquid phase within the porous volume. First, this affinity will condition wettability and accessibility of the whole porosity. From this point of view a good interaction is required. Second, in some case, it has been noticed that the better the affinity is, the deeper the ion mobility falls. 6.2. Ionic mobilities

Using pulse field gradient spin-echo NMR, it is possible to access the diffusion coefficients of several nuclei such as 19F, 31P, 1H, and 7Li present in the anion, solvents, polymer and cation, and to deduce transference number of anions and cations. Capiglia et al. [94], comparing the cationic transference number, t, of LiPF6 in pure liquid electrolyte and in PVdF gelled by the same liquid electrolyte, found an increase in cationic transference number in the gel. Kataoka et al. [93] performed investigations on macroporous PVdF by the same method and concluded that the conduction mechanism is controlled not by the liquid electrolyte filling the porosity but by the polymer-rich region swollen by the electrolyte. Saunier et al. [77], using the same method on similar macroporous PVdF (PVdF Kynar 301F®), disagree with this conclusion as the conductivity contribution of the gelled part of that membrane is quite negligible. They found that both solvent and ions are slowed down in the porosity. However, the slowing down of the cation is higher than that of the anion, leading to a decrease in t value with regard to the liquid electrolyte. 7. BATTERY PERFORMANCES The comparison of gelled polymer electrolytes through the battery performances is not easy. On the one hand, battery performances are not only tributary to the intrinsic performances and to the purity of the materials, i.e. electrodes, electrolyte, separators, and current collectors, but also to technological aspects (knowhow), often confidential, in relation with the (i) possible use of additives, (ii) composite electrode processing and (iii) battery assembly. On the other hand,

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companies have specialized technicians or engineers to perform the battery tests while this activity, clearly technological, is generally far from the concerns of scientific laboratories. Lastly, battery tests often deal with different positive or negative electrodes, making the comparison of the electrolyte performances difficult. Tarascon et al. [79], from the Bellcore group, manufactured “plastic rechargeable Li-ion batteries,” and claimed, for a plastic LiMn2O4/graphite battery, specific and volumetric energies of, respectively, 110 Wh/kg and 280 Wh/L and a lifespan  2000 cycles at 25°C (rate  1C) using an innovative technology [95] based on lamination of fluorinated copolymers. After the first announcement by Boudin et al. [29], from SAFT-France Company, of the performances obtained using a porous PVdF separator, scientists and companies moved from dense fluorinated separators towards porous ones. Thus, Du Pasquier et al. [96], from Telecordia (formerly Bellcore), reported the performances of LiMn2O4/graphite batteries of 35 and 115 mAh, based on PVdF–HFP macroporous separator. The capacity retention of the 35 mAh at a C/2 rate and at room temperature ranged between 80 and 70% after 500 cycles. Prosini et al. [97], from ENEA, reported performances of LiMn2O4 batteries using macroporous VdF–HFP copolymer filled by inorganic oxides such as MgO. Unfortunately, this preliminary study did not lead to a battery exceeding 70 cycles. More recently, Saunier et al. [68] reported battery tests performed at SAFT Recherche using a macroporous PVdF membrane. Lithiated cobalt oxide and a mixture of MCMB  graphite were, respectively, used as positive and negative electrodes. The liquid electrolyte consisted of a molar solution of LiPF6 in a ternary solvent mixture EC/DMC/DEC (2:2:1, in by vol.). The specific energy is 95 Wh/kg for one monostack and increases up to 130 Wh/kg when 6–7 monostacks are assembled in parallel. The battery charge is shown in Fig. 8. When it is carried out at 4.1 V and C rate, 82% of the capacity is

Fig. 8. GSM discharge curves at different temperatures.

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recovered after 1 h. The discharge was performed at different rates. At slow discharge rate of C/5, the capacity ranges between 93 and 115 mAh, while at 1C and 2C, respectively, 6 and 40% capacity loss are observed. As the discharge rate of a cellular phone is close to 0.4C, limitations due to lithium diffusion should not occur. GSM discharges are pulsed discharges, i.e. a sequence of peak current and a baseline current. The GSM discharge curves at different temperatures, corresponding to a 2C peak current for 0.5 ms, followed by a C/2 baseline current for 4.5 mS are also presented in Fig. 8. The initial autonomy is 1.78 h and decreases to 1.35 h after 400 cycles while the specifications of GSM application are respectively, 1.5 and 1.2 h. 8. CONCLUSION Thirty years ago Feuillade and Perche et al. [19], pioneers in the use of polymer electrolytes in thin lithium batteries wrote: The development of batteries using such ionic gels therefore combines non-classical macromolecular techniques with the fundamental electrochemical behaviour of batteries. Because of limited contact between electrochemical technology and macromolecular chemistry, the combination is rarely seen in the literature. Applications in common aqueous batteries are generally empirical solutions of particular needs, using macromolecular materials which may not be optimum to the electrochemical requirements Has the situation evolved over the last 30 years? Indeed, much attention is paid to material science in electrochemistry but this expertise has been mainly concentrated on inorganic materials used both in negative and positive electrodes. Yet all kinds of electrolytes are essential to improve the battery performances and the conduction mechanisms are far from being understood, even in liquid electrolytes. However, it is very important to develop in situ, i.e. in laboratories, a good expertise in polymer science and electrochemistry. The use of PVdF-based polymer electrolytes is very old, and it is time to move towards functional VdF polymers, either by chemical modifications of available polymers or by designing new monomers and polymers. Are the partially fluorinated polymers, such as PVdF, threatened by the necessary protection of the environment? In that case, VdF copolymers incorporating significant amount of non-fluorinated comonomers would probably be less threatened or more easily tolerated. However, is the market wide enough to justify a significant industrial research? The answer is probably negative if these polymers are designed exclusively for the 4C batteries, the prices of which are very tense. But the situation is different for large lithium-ion batteries intended for electric or

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hybrid vehicles where safety is a deciding criterion. In addition, polymer electrolytes should be designed to be used both in lithium batteries and in fuel cells, thus expanding the market and lowering the production costs. ACKNOWLEDGMENTS We acknowledge Cristina Iojoiu for her collaboration to the figures in this paper. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21] [22] [23] [24]

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[66] J. Saunier, F. Alloin, and J.-Y. Sanchez, Colloque Gaston Planté, Paris, 30–31 October 2000, Proc., pp. 7–12. [67] J. Saunier, F. Alloin, J.Y. Sanchez, and B. Barrière, J. Polym. Sci., B: Polym. Phys., 42 (2004) 532. [68] J. Saunier, F. Alloin, J.Y. Sanchez, and G. Caillon, J. Power Sources, 119–121 (2003) 454. [69] P.J. Flory, J. Chem. Phys., 10 (1942) 51. [70] M.L. Huggins, Annals of the New-York Academy of Science, XLIII (1942) 1. [71] B. Barrière, Ph.D. Thesis, Université Paris VI, 1997. [72] B. Barrière and L. Leibler, J. Polym. Sci. B: Polym. Phys. 41 (2003) 166. [73] W.Y. Chuang, T.H. Young, D.M. Wanh, R.L. Luo, and Y.M. Sun, Polymer, 41 (2000) 8339. [74] E. Favre, Q.T. Nguyen, R. Clément, and J. Neel, Eur. Polym. J, 32 (1996) 303. [75] S. Abbrent, J. Lindgren, J. Tegenfeldt, J. Furneaux, and A. Wendsjö, J. Electrochem. Soc, 146 (1999) 3145. [76] M.M.E. Jacob and A.K. Arof, Electrochim. Acta, 44 (1999) 2909. [77] J. Saunier, W. Gorecki, F. Alloin, and J.Y. Sanchez, J. Phys. Chem., Part. B, 109 (7), (2005) 2487. [78] M. Caillon-Caravanier, B. Claude-Montigny, D. Lemordant, and G. Bosser, J. Power Sources, 107 (2002) 125. [79] J.M. Tarascon, T. Gozdz, C. Schmutz, F. Shokoohi, and P.C. Warren, Solid State Ionics, 86–98 (1996) 49. [80] T. Alfred, E.F. Gurne, and W.G. Loyd, J. Polym. Sci. C: Polym. Symp., 12 (1966) 249. [81] J. Cranck and G.S. Park, Diffusion in Polymers Academic Press, London, 1968. [82] M.S. Hendeqvist and W.W. Gedde, Polymer, 40 (1999) 2381. [83] E. Bagley and F.A. Long, J. Amer. Chem. Soc., 77 (1955) 2172. [84] G.S. Park, Diffusion in Polymers, J. Cranck, and G.S. Park (Eds.), Academic Press, London, 1968 p. 141. [85] M.S. Hendeqvist, M. Kroock, and U.W. Gedde, Polymer, 43 (2002) 3061. [86] G.C. Sarti, Polymer, 20 (1979) 827. [87] J.M. Guenet and G.B. Mc Kenna, Macromolecules, 21 (1988) 1752. [88] S. Mal and A.K. Nandi, Langmuir, 14 (1998) 2238. [89] Z. Jiang, B. Caroll, and K.M. Abraham, Electrochim. Acta, 42 (1997) 2667. [90] M. Tazaki, R. Wada, M. Okabe, and T. Homma, J. Appl. Polym. Sci., 65 (1997) 1517. [91] J. Saunier, F. Alloin, J.Y. Sanchez, and L. Maniguet, J. Polym. Sci., B: Polym. Phys., 42 (2004) 2308. [92] X. Liu, H. Kusawake, and S. Kuwajima, J. Power Sources, 97–98 (2001) 661. [93] H. Kataoka, Y. Saito, T. Sakai, E. Quarterone, and P.J. Mustarelli, Solid State Ionics, 122 (1999) 285. [94] C. Capiglia, Y. Saito, H. Yamamoto, H. Kageyama, and P. Mustarelli, Electrochim. Acta, 45 (2000) 1341. [95] T. Gozdz, C. Schmutz, J.M. Tarascon, and P.C. Warren, U.S. Patent 5,456,904 July 30, 1996. [96] A. Du Pasquier, P.C. Warren, D. Culver, A.S. Gozdz, G.G. Amatucci, and J.M. Tarascon, Solid State Ionics, 135 (2000) 249. [97] P.P. Prosini, P. Villano, and M. Carewska, Electrochimi. Acta, 48 (2002) 227.

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335

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Lithium-ion-conductive polymer electrolytes exhibit a high lithium-ion transference number with the incorporation of fluorine atoms Takeshi Abe and Zempachi Ogumi Graduate School of Engineering, Kyoto University, Nishikyo-ku, Kyoto 615-8510, Japan 1. INTRODUCTION Since Sony commercialized lithium-ion batteries in 1991, most portable electric devices now use these batteries due to their light weight and high-energy density. Recently, lithium-ion batteries have received considerable attention for the use as power sources in hybrid electric vehicles (HEV) and electric vehicles (EV). However, to use lithium-ion batteries for HEVs, serious safety issues must be solved. There have been many approaches to enhance the safety of lithium-ion batteries. Among these, the use of lithium-ion-conductive polymer electrolytes or inorganic electrolytes has been shown to drastically improve the safety of these batteries. The latter inorganic electrolytes have attracted much attention since very high ion-conductive sulfide-based electrolytes have been reported [1,2]. However, the compatibility of battery-active materials with inorganic electrolytes is a major problem for the practical use of inorganic electrolytes in lithium-ion batteries. All-solid batteries using inorganic electrolytes should be quite safe in this regard. The compatibility at the electrode/electrolyte interface is greatly enhanced when polymer electrolytes are used, and therefore, many researchers have used polymer electrolytes to fabricate all-solid batteries. Polymer electrolytes were first reported by Wright and co-workers [3]. Since then, much work has been done on ion-conductive polymer electrolytes for the use in lithium-ion batteries, sensors, electrochromism, etc. Regarding the polymer electrolytes, there is no leakage of liquid electrolytes and devices with various shapes can be fabricated due to their flexibility. However, several problems must be solved before the practical use of polymer electrolytes. For example, low ionic conductivities of polymer electrolytes will impair the

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performance of batteries and low-lithium-ion transference numbers are serious drawbacks for high-power use of these batteries. Many approaches have been considered to increase the ion conductivities and lithium-ion transference numbers [4–15]. Croce et al. [14] reported that polymer electrolytes containing nanometer-sized ceramic powders of TiO2 and Al2O3 showed high conductivity at moderate temperatures as well as high lithium-ion transference numbers. Sun et al. [15] showed that the addition of ferroelectric materials as a ceramic filler to polyethylene oxide mixed with various lithium salts enhanced the ionic conductivity at low temperatures and that interfacial resistance between the lithium metal and polymer electrolyte decreased. The above results can be explained by Lewis acid–base reactions between the lithium salt and the surface of the ceramic filler, and a high ionic conduction path may be created near the ceramic surface [16]. Fujinami et al. [7] reported siloxyaluminate polymers with high lithium-ion conductivities. Anion in the polymer electrolyte is bonded to the polymer backbone, and therefore, a high lithium-ion transference number can be achieved. Recently, Sun et al. reported a lithium-ion battery electrolyte based on LiF and tris(pentafluorophenyl)borane (TPFB) in 1,2-dimethoxyethane [17–20]. They used TPFB as an anion receptor. Abe et al. reported lithium-ion-conductive polymer electrolytes containing TPFB and found an increase in the lithium-ion transference number [21]. Fujinami’s group showed that aluminate and borate complex polymers containing fluoroalkane dicarboxylate exhibited high lithiumion transference numbers [22,23]. As mentioned above, fluorine atoms play an important role in the enhancement of the lithium-ion transference number. This chapter describes polymer electrolytes that exhibit a high lithium-ion transference number with the incorporation of fluorine atoms. 2. VARIOUS ANION RECEPTORS A new family of boron-based anion receptors was first reported by Lee et al. [24]. The strategy using anion receptors in non-aqueous electrolytes involves the reduction of ion pairing. Since the dielectric constants of organic solvents used in non-aqueous electrolytes are usually lower than those that of water, ion pairs remain in the electrolyte, resulting in a decrease in ionic conductivity and in the lithium-ion transference number. Fig. 1 shows the chemical structure of boron-based anion receptors containing fluorine atoms reported by Lee et al. [24]. Due to the high electronegativity of fluorine atom, boron atom probably possesses a partial positive charge, which will trap an anion via Lewis acid–base interaction. Using the anion receptors shown in Fig. 1, Lee et al. examined the ionic conductivities of lithium salts of various concentrations of 1,2-dimethoxyethane (DME). Typical examples are given in Table 1.

Lithium-ion-conductive polymer electrolytes exhibit a high lithium-ion transference number with the incorporation of fluorine atoms

337

Fig. 1. Chemical structures of fluorinated boron-based anion receptors reported by Lee et al. [24].

Table 1 Ionic conductivities of electrolytes consisting of lithium salts, anion receptors, and 1,2-dimethoxyethane (DME) Conductivity (103; S/cm) Concentration (mol dm3)

CF3COOLi

C2F5COOLi

LiF

Anion receptor : (C6F5O)3B 0.2

3.3

3.2

4.1

0.5

5.9

6.1

6.0

1.0

5.2

5.8

6.8

Anion receptor : (C6F5)3B 0.2

3.2

3.0

1.7

0.5

6.8

5.4

6.4

1.0

6.0

5.4

6.2

The dielectric constant of DME is very low, therefore, lithium salts hardly dissolve in DME. As shown in Table 1, lithium salts of CF3COOLi, C2F5COOLi, and LiF give high ionic conductivities. LiF is insoluble even in solvents with high dielectric constants. However, even LiF in DME shows high ionic conductivities due to the boron-based anion receptors. Therefore, anion receptors are promising compounds for use as electrolytes in lithium-ion batteries. 3. LITHIUM-ION-CONDUCTIVE ELECTROLYTES THAT USE ANION RECEPTORS We previously reported polyethylene oxide (PEO)-based polymer electrolytes containing an anion receptor, tris(pentafluorophenyl)borane (TPFB) [21]. Polyethylene oxide (MW  400,000), and lithium salts of LiCF3SO3 and LiF

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were used without further purification. The anion receptor TPFB was used as received. To decrease the resistance of polymer electrolytes, 10 wt% of lowmolecular-weight plasticizer (polyethylene glycol dimethylether, MW  250) was added to the electrolytes. The ratio of ethylene oxide unit to lithium ion was maintained at EO/Li  25, and that of TFPB/Li was varied from 0 to 1. The ionic conductivities of the resultant polymer electrolytes were evaluated in a cell composed of stainless-steel (SUS)/polymer electrolyte/SUS using a Solartron 1260 frequency response analyzer over a frequency range 1–100 kHz. Various methods for evaluating the lithium-ion transference number have been reported [25–27]. We used the method reported by Evans et al. [28] due to its ease of use. Fig. 2(a) shows the temperature dependence of ionic conductivities for polymer electrolytes prepared from PEO, LiCF3SO3, PEGDME, and various amounts of TFPB. Solid circles show ionic conductivities of polymer electrolytes without TFPB. In Fig. 2(b), ionic conductivities at 30°C are plotted against the ratios of TPFB/TFA, where TFA denotes LiCF3SO3. The following main points should be noted. (1) With an increase in the ratio of TPFB/LiCF3SO3 from 0 to 0.1, the ionic conductivity increases slightly, as clearly shown in Fig. 2(b), except at TPFB/TFA  0.033. This is because TPFB promotes the dissociation of LiCF3SO3. (2) Ionic conductivities decrease with an increase in the ratio of TPFB/LiCF3SO3 from 0.1 to 1 at temperatures below 50°C. This decrease in ionic conductivity can be ascribed to the interaction between CF3SO3 anion and TPFB in the polymer electrolytes. The formation of a CF3SO 3 anion and TPFB complex should suppress anion mobility in the electrolytes, and therefore, the total conductivity should decrease. The decrease in anion conductivity reflects an increase in the lithium-ion transference number, which will be discussed later. At higher temperatures, the ionic conductivities are almost the same regardless of the amount of TPFB. The difference in the interaction between the anion and TPFB becomes less important at high temperatures. (3) Finally, the ionic conductivities of polymer electrolytes using PEO as a host matrix usually increase at temperatures between 60 and 70°C. However, a rapid increase in the ionic conductivity of a polymer electrolyte with a TPFB/LiCF3SO3 ratio of 1 is observed between 50 and 60°C. Fig. 3 shows DSC data for polymer electrolytes containing various amounts of TPFB. Peaks in Fig. 3 are due to the transition of crystalline PEO to amorphous PEO. The transition temperatures clearly decrease with increasing amounts of TPFB. Further, the transition temperature of a polymer electrolyte with TPFB/LiCF3SO3  1 is 58.2°C. These DSC data are consistent with the ionic conductivities in Fig. 1, and also indicate that TPFB acts as a plasticizer to some extent.

Lithium-ion-conductive polymer electrolytes exhibit a high lithium-ion transference number with the incorporation of fluorine atoms

339

t (°C) 90

80

70

60

50

40

30

3.2

3.3

−3.5 −4.0

log (/ Scm−1)

−4.5 −5.0 −5.5 −6.0

TPFB/TFA =1 TPFB/TFA = 0.5 TPFB/TFA = 0.2 TPFB/TFA = 0.1 TPFB/TFA = 0

−6.5 −7.0 −7.5 2.8

2.9

(a)

3.0

3.1 1000 /T (K−1)

−5.0 30°C

log (/ Scm−1)

−5.5

−6.0

−6.5

−7.0

−7.5 0.0 (b)

0.2

0.4

0.6

0.8

1.0

TPFB/TFA

Fig. 2(a). Temperature dependency of ionic conductivities for polymer electrolytes prepared from PEO, LiCF3SO3, PEGDME, and various amounts of TFPB. TFA denotes LiCF3SO3. (b). Dependency of ionic conductivities for polymer electrolytes prepared from PEO, LiCF3SO3, PEGDME and TPFB/TFA at 303 K.

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Takeshi Abe and Zempachi Ogumi

0

Heat Flow(μW)

−2000

−4000

TPFB/TFA= 1 TPFB/TFA= 0.5 TPFB/TFA= 0.33 TPFB/TFA= 0

58.2 60.8

−6000

63.7 61.8

−8000 40

60

80

100

120

Temperature (°C)

Fig. 3. DSC curves of polymer electrolytes prepared from PEO, LiCF3SO3, PEGDME, and various amounts of TFPB.

Fig. 4 shows the ionic conductivities of polymer electrolytes containing LiF and TPFB. Since it is very difficult for LiF to dissociate PEO with a low dielectric constant, ionic conductivity could not be measured without the addition of TPFB. As shown in Fig. 4, the ionic conductivities of polymer electrolytes increased with an increase in the amount of TPFB. This is simply due to an increase in the carrier concentration by the dissociation of LiF promoted by the anion receptor TPFB. Fig. 5 shows the lithium-ion transference number against the ratio of TPFB/Li salt. In a qualitative comparison, the lithium-ion transference number increased with an increase in the TFPB/Li ratio. As discussed above, an increase in the transference number can be ascribed to the formation of a complex between CF3SO3 anion or F anion and TPFB, leading to the suppression of anion mobility. The lithium-ion transference numbers of polymer electrolytes using LiF are larger than those using LiCF3SO3. This difference can be ascribed to the sizes of the anions. Since F anion is smaller than CF3SO3 anion, there is greater interaction between F and TPFB than between CF3SO3 and TPFB. Therefore, some CF3SO3 anion is not trapped by TPFB, which leads to a decrease in the lithium-ion transference number. As shown in Fig. 5, a lithium-ion transference number of 0.5 was obtained with the addition of TPFB. Therefore, the addition of TFPB drastically enhances

Lithium-ion-conductive polymer electrolytes exhibit a high lithium-ion transference number with the incorporation of fluorine atoms

341

t (°C) 110 100 90

80

70

60

50

40

30

−3

log ( / Scm−1)

−4

−5

−6

TPFB/LiF = 1 TPFB/LiF = 0.5 TPFB/LiF = 0.2 TPFB/LiF = 0.1 TPFB/LiF = 0.05

−7 2.6

2.8

3.0 1000 / T

3.2

3.4

(K−1)

Fig. 4. Temperature dependency of ionic conductivities for polymer electrolytes prepared from PEO, LiF, PEGDME, and various amounts of TFPB. 0.7 0.6 Li+ transference number

LiF

0.5 0.4 0.3

LiCF3SO3

0.2 0.1 0.0 0.0

0.2

0.4

0.6

0.8

1.0

TPFB/Li

Fig. 5. Lithium-ion transference number of polymer electrolytes prepared from PEO, LiCF3SO3 or LiF, PEGDME and various amounts of TFPB.

342

Takeshi Abe and Zempachi Ogumi

the lithium ion-transference number. In contrast, total ionic conductivities decrease with the addition of TPFB, as shown in Fig. 4. Other lithium salts of LiN(C2F5SO2)2 (LiBETI) and LiClO4 were used to determine the effect of TPFB. Fig. 6 shows the ionic conductivities of PEO mixed with LiBETI containing PFB in various molar ratios. The polymer electrolytes were prepared as described above. In contrast to the results shown in Fig. 4, the ionic conductivities decreased with an increase in TPFB. This decrease is simply due to a decrease in anion conduction. LiClO4 gave similar results. The results shown in Fig. 6 indicate that even the large anion, N(C2F5SO2)2, is trapped by TPFB. Fig. 7 shows the correlation between the molar ratio of TPFB and the lithiumion transference number. The results for LiCF3SO3 are shown for comparison. The transference number increased with an increase in anion receptor, which is in good agreement with the results given in Fig. 5. For LiF, LiCF3SO3, LiBETI, and LiClO4, the lithium-ion transference numbers were in the order LiF  LiCF3SO3  LiClO4, LiBETI. Again, a large anion is associated with a decrease in the transference number due to a decrease in interaction between the anion and anion receptor. t (°C) 110 100 90

80

70

60

50

40

30

20

−3

log ( /Scm−1)

−4

−5

−6

BETI/LiF = 1 BETI/LiF = 0.5 BETI/LiF = 0.2 BETI/LiF = 0.1 BETI/LiF = 0.05

−7

2.6

2.8

3.0 1000 /T

3.2

3.4

(K−1)

Fig. 6. Temperature dependency of ionic conductivities for polymer electrolytes prepared from PEO, LiBETI, PEGDME, and various amounts of TFPB.

Lithium-ion-conductive polymer electrolytes exhibit a high lithium-ion transference number with the incorporation of fluorine atoms

343

Li+ ion transference number (t+)

0.6 LiTFA LiBETI LiClO4 LiF

0.5

0.4

0.3

0.2

0.1 0.0

0.2

0.4

0.6

0.8

1.0

PFB/Li+

Fig. 7. Lithium-ion transference number of polymer electrolytes prepared from PEO, LiCF3SO3, LiBETI, LiClO4, LiF, PEGDME, and various amounts of TFPB.

Based on the results shown in Figs. 4 and 6, the ionic conductivities seem to be too low. However, the use of a comb-like polymer instead of PEO enhanced the ionic conductivities, and we obtained ionic conductivities of 104 S/cm with a transference number of 0.33. Therefore, polymer electrolytes containing anion receptor should be promising for the use in lithium-ion batteries. 4. LITHIUM-ION-CONDUCTIVE ALUMINATE AND BORATE COMPLEX POLYMERS [22] Single-ion-conducting polymer electrolytes containing a fluoroalkane dicarboxylate-substituted aluminate or borate backbone and two methoxy [oligo (ethyleneoxide)] side chains directly bonded to the ate complex centers (aluminum and borate) were reported by Fujinami et al. [22]. The incorporation of a Lewis acid or electron-withdrawing groups into the inorganic backbone facilitates the delocalization of negative charge to reduce ion pairing, which results in a drastic enhancement of conductivity. Aluminate or borate polymers were prepared as shown in Scheme 1. The ionic conductivities of polymers A and B are shown in Fig. 8. For both aluminate and borate polymers, ionic conductivity was enhanced with the incorporation of longer ether chains from n  3 to 11.8. This is commonly observed in polymer

344

Takeshi Abe and Zempachi Ogumi − 80 oC

LiMH4 + 2 ROH

LiMH2(OR)2

THF R = CH3(OCH2CH2)n

n = 3, 7.2, 11.8

OR

LiMH2(OR)2 + HOOC(CF2)3COOH

− 80 oC

Li+ [–M −–OC(CF2)3CO–]m

THF

OR O

M=Al : Polymer A (n = 3, 7.2, 11.8) M=B : Polymer B (n = 3, 7.2, 11.8)

LiMH2(OR)2 + HOOC(CH2)3COOH

− 80 oC

OR

Li+ [–M-–OC(CH2)3CO–]m

THF

OR O

M=B : Polymer C (n = 11.8)

Scheme 1. Synthesis of fluorinated ate complex polymers.

−4.0 B (n = 11.8) B (n = 7.2)

−4.5 B (n = 3) −5.0

log (Scm−1)

A (n = 11.8) −5.5 −6.0 A (n = 7.2) −6.5 −7.0

A (n = 3)

−7.5 −8.0 2.8

2.9

3.0

3.1 1000 / T

3.2

3.3

3.4

(K−1)

Fig. 8. Temperature dependency of ionic conductivity for polymers A and B.

Lithium-ion-conductive polymer electrolytes exhibit a high lithium-ion transference number with the incorporation of fluorine atoms

345

electrolytes of this type and can be ascribed to the increasing organic component of the polymer, since lithium-ion motion is promoted by the segment motion of the oligoether chains. Borate polymers showed higher ionic conductivity than aluminate polymers. The maximum conductivity was 1 106 S/cm (30°C) with n  11.8 for aluminate polymers and 1105 S/cm (30°C) with n  11.8 for borate polymers. In Fig. 9, the formation of ion pairing between the lithium ion and oxygen atoms around the ate complex center is shown based on optimization calculations by MOPAC (PM5). While the ate complex center has a formal negative charge, a positive charge on the ate complex center and a dispersed negative charge on the oxygen atoms around the ate complex center were shown by the optimization calculations by MOPAC. The negative charge of the oxygen atoms for borate was smaller than that for aluminate. This can be ascribed to the large positive charge on the aluminum atom due to the smaller electronegativity of the aluminum atom compared to boron. Therefore, weaker ion pairing between the lithium ion and oxygen atoms was estimated from the smaller negative charge on the oxygen atoms for the borate polymer, and resulted in an increase in the number of charge carriers and accounted for the enhancement of ionic conductivity. The enhancement of ionic conductivity with the incorporation of electronwithdrawing groups was investigated by comparing fluoroalkane- and alkanesubstituted polymer. The temperature dependence of the ionic conductivities of polymer B (n  11.8) and polymer C (n  11.8) is shown in Fig. 10. Polymer B (n  11.8) exhibited higher ionic conductivity than polymer C (n  11.8) by two orders of magnitude. This can be ascribed to the difference in the electronic effect between the trifluoromethyl group and the methyl group. Therefore, a weaker interaction between the lithium ion and oxygen atoms was estimated from a smaller negative charge on oxygen atoms for the electron-withdrawing fluoroalkane-substituted polymer, and this resulted in the enhancement of ionic conductivity. The lithium-ion transference number of polymer B (n  11.8) was determined to be 0.95 by the Evans method as modified by Abraham et al. [9] and polymer B was confirmed to be an approximately single-ion conductor.

Fig. 9. Formation of ion pairing between lithium ion and oxygen atoms.

346

Takeshi Abe and Zempachi Ogumi −4.0 B (n = 11.8)

log  (Scm-1)

−5.0

−6.0 C (n = 11.8)

−7.0 2.8

2.9

3.0

3.1 3.2 1000 /T (K−1)

3.3

3.4

Fig. 10. Temperature dependence of ionic conductivity for polymers B and C.

5. SUMMARY Based on the results in sections 3 and 4, polymer electrolytes with high lithiumion transference numbers are likely to be obtained with the incorporation of fluorine atoms. Although the ion conductivities of the resultant polymer electrolytes are not sufficiently high for practical use in lithium-ion batteries at room temperature, the values can be enhanced by the incorporation of plasticizers, ionic liquids, etc. Further studies are required. REFERENCES [1] [2] [3] [4] [5] [6] [7]

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Chapter 16

Room-temperature molten salts as new electrolytes Rika Hagiwara and Kazuhiko Matsumoto Graduate School of Energy Science, Kyoto University, Sakyo-ku, Kyoto 606-8501, Japan 1. ROOM-TEMPERATURE MOLTEN SALTS OF FLUOROANIONS Salts such as sodium chloride need high temperatures to melt, although mixing these considerably lowers the melting points. However, some salts, mostly organic, are known to melt below room temperature. These salts are called room-temperature molten salts (RTMSs) or room-temperature ionic liquids. When RTMSs are used as electrolytes in electrochemical devices, the following advantages are expected: nonvolatility makes their handling easier and prevents the electrolyte from drying up; nonflammability improves the safety of devices; and a wide electrochemical window raises power and energy densities. A wide liquid-phase temperature range enables the operation of the devices in various environments. However, in the case of classical chloroaluminate salts ((cation)(AlCl3)nCl), difficulty in their handling that arose from their instabilities against moisture tended to give an impression to researchers that RTMSs were not easy to handle, in spite of their excellent functionalities as described above. In 1992, novel moisture-stable RTMSs, 1-ethyl-3-methylimidazolium tetrafluoroborate (EMImBF4) and 1-ethyl-3-methylimidazolium triflate (EMImCF3SO3), were reported for the first time [1,2]. After these reports, a wide variety of moisture-stable RTMSs have been reported, most of them having been applied as reaction solvents of organic syntheses and contain organic or inorganic fluoroanions as counteranions that are combined with some onium cations [3–6]. In this chapter, syntheses and properties of RTMSs containing fluoroanions are described. Also, a brief introduction is given on their applications to electrochemical devices such as electric double-layer capacitors (EDLCs), fuel cells, lithium batteries and dye-sensitized solar cells.

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2. SYNTHESES OF ROOM-TEMPERATURE MOLTEN SALTS 2.1. Molecular design of room-temperature molten salts

Fig. 1 shows typical cations used for RTMSs. Generally (1) salts of aromatic cations such as alkylimidazolium and alkylpyridinium ions exhibit low melting points compared with the salts of nonaromatic cations such as tetraalkylammonium and dialkylpyrrolidinium ions; (2) the chemical and electrochemical stabilities of nonaromatic cation-based salts are higher than those of aromatic cation-based salts; and (3) the melting point and conductivity of the salt decreases with the increase in carbon number of alkyl side-chain on the cation. In addition, substituted groups on the cations are also important factors to determine the properties of the salt. Although RTMSs of the cations with perfluoroalkyl chains usually do not exhibit low melting points, some of them combined with bis(trifluoromethylsulfonyl)amide ion (CF3SO2)2N form liquid at room temperature [7,8]. From the viewpoint of environmental safety, halogen-free anions are being developed today as counteranions for RTMSs. However, as mentioned above, fluoroanions still attract much attention because of their superior characteristics. Organic anions containing perfluoroalkyl chains as well as some typical inorganic fluorocomplex anions such as BF4 and PF6 are widely studied. 2.2. Synthetic methods of room-temperature molten salts

For the preparation of RTMSs, halide salts, (CXs), containing cations for RTMSs, are often used as starting materials. They are prepared by the reaction of the corresponding halogenoalkane and amine. For example, one of the most

Fig. 1. Cations used for the syntheses of RTMSs.

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351

popular salts, 1-ethyl-3-methylimidazolium chloride, EMImCl, is synthesized by the reaction of 1-methylimidazole (MeIm) and chloroethane in acetonitrile: MeIm  CH3CH2Cl → EMImCl

(1)

Except for the reactions of volatile chloroalkane like chloroethane performed in an autoclave, other reactions are usually carried out under refluxing. All the reagents should be well dried before use. The next procedure is the metathesis of CX and MA (M, alkali metal, silver or ammonium cation; A, anion for RTMS) in a solvent. One of the preparative methods of EMImBF4 is the reaction of EMImI and AgBF4 in a water/methanol solvent [1]: CX  MA → CA  MX

(2)

Since RTMSs are involatile and cannot be distilled, the by-product MX is removed only by filtration or extraction. However, the halide is not completely eliminated in this procedure because the solubility of MX in RTMS is not absolutely zero. In the case of a water-insoluble salt such as 1-ethyl-3-methylimidazolium bis(trifluoromethanesulfone)amide (EMIm(CF3SO2)2N), washing with water is an effective method to remove the halide impurities. A protic acid HA or its solution is occasionally used instead of MA, where the by-product HX is completely removed from the RTMS obtained in vacuo at elevated temperatures: CX  HA → CA  HX↑

(3)

Fluorohydrogenate salts are synthesized in the same manner by the reactions of starting chlorides and large excess of anhydrous hydrogen fluoride (aHF), and the vacuum stable liquid salts of various cations such as alkylimidazolium and alkylpyrrolidinium possess the same HF composition formulated by (cation)(HF)2.3F at room temperature [9–12]. They are composed of the cation and two fluorohydrogenate anions, (HF)2F and (HF)3F, shown in Fig. 2, in the ratio to give the composition shown above. The reactions of amine and ester, giving alkylammonium salt directly, are also used for preparations of RTMSs. For example, the reaction of 1-ethylimidazole (1-EIm) and methyltriflate (MeOSO2CF3) in 1,1,1-trichloroethane gives EMImCF3SO3 [7]. This reaction does not produce any by-product and is regarded as a clean method, although rigorous anhydrous condition is required to avoid hydrolysis. EIm  MeOSO2CF3 → EMImCF3SO3

(4)

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Fig. 2. Fluorohydrogenate anions: (a) μ-fluoro-bis(fluorohydrogenate) ion ((HF)2F); (b) μ3-fluoro-tris(fluorohydrogenate) ion ((HF)3F).

The salts of tertiary ammonium cations are prepared by neutralization of starting tertiary amine with acid. For example, 1-methylimidazolium tetrafluoroborate (MeImBF4) is synthesized by neutralization of MeIm by HBF4 [13]: MeIm  HBF4 → MeImBF4

(5)

In the case of tertiary ammonium salt, it should be noted that some molecular species on the left hand side remains in the reaction product due to the equilibrium of reaction (5). A chloroaluminate salt is synthesized by the reaction of a chloride and aluminum trichloride [14,15]: EMImCl  nAlCl3 → EMIm(AlCl3)nCl

(6)

where the salts with n 1, n  1 and n  1 are called basic, neutral and acidic, respectively. Bromoaluminate salts are also obtained in the same manner using corresponding bromides [16]. Tetraalkylammonium fluorides with linear alkyl chains are isolated and the reactions with fluoroacids give the salts of the fluorocomplex anion [17]. On the other hand, alkylimidazolium and alkylpyridinium fluorides are unstable and have never been isolated as solvent-free solid. A monohydrate salt of 1-butyl-3-methylimidazolium fluoride (BMImF·H2O) has been recently reported as a decomposition product of BMImPF6 [18]. Fluorohydrogenates and solvated fluorides mentioned above are suitable starting reagents, in which the fluorobasicity is decreased to stabilize the salts. Reactions of EMIm(HF)2.3F and some fluoroacids have been recently reported to produce various RTMSs [19,20]: EMIm(HF)nF  MFm → EMImMFm1  nHF

(7)

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EMIm(HF)nF itself is a stable RTMS with a low viscosity and high conductivity as described below. In reaction (7), only a volatile by-product HF is formed to provide the salts with high purities. 3. PROPERTIES OF ROOM-TEMPERATURE MOLTEN SALTS RTMSs containing EMIm cations exhibit superior physical properties such as low viscosity and high conductivity, thus being the most extensively studied and reported so far. Physical properties of selected EMIm-based RTMSs are listed in Table 1. Physical properties of RTMSs of the other imidazolium cations are summarized in Ref. [6,21]. 3.1. Melting point and glass transition temperature

Generally, if the same cation is combined, a salt of a large and/or asymmetric anion tends to exhibit a low melting point. Fig. 3 shows the relationship between the size of inorganic anion with relatively high symmetries and melting point of EMImMFm1 (MFm1  BF4, PF6, AsF6, SbF6, NbF6, TaF6, WF7) [20], where the radius of the central atom in the fluorocomplex anion is plotted on the horizontal axis [22]. EMImPF6, EMImAsF6 and EMImSbF6 are isostructural with each other. For the hexafluorocomplex (MF6) salts, the melting point decreases linearly with increase in the size of the anion. EMImNbF6 and EMImTaF6, whose molar volumes are close to each other (190 and 187 cm3), exhibit very close melting points. For the EMIm salts of nonoctahedral anions, the relationship between the melting point and anion size for EMImWF7 is in accordance with that of the octahedral anions, whereas EMImBF4 exhibits a significantly lower melting point than that expected from the anionic volume. The tetrahedral BF4 anion is expected to have a different interaction with the cation from those of hexa- or heptafluorocomplex anions. EMIm(HF)2.3F exhibits a significantly low melting point compared with other EMIm salts. In the case of RTMSs of large organic anions like (CF3SO2)2 N, there is no clear relationship between the melting points and sizes of the anions. An asymmetric anion, (CF3CO)(CF3SO2)N, also gives salts with low melting points in combination with symmetric cations including nonaromatic cyclic alkylammonium cations, in which the salts usually exhibit high melting points [23,24]. 3.2. Conductivity and viscosity

Aqueous electrolytes usually possess higher conductivities than those of RTMSs, ranging from 102 to 103 mS cm1, for example, about 800 mS cm1 for 30 wt% H2SO4 (aqueous solution). The high conductivity is caused by the proton-hopping mechanism as well as the low viscosity of the solution. On the other hand, the conductivities of RTMSs are of the order of 100–101 mS cm1 (for example, 14 mS cm1 for EMImBF4 and 8.4 mS cm1 for EMIm(CF3SO2)2N).

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Table 1 Physical properties of EMIm cation-based RTMSs Melting point (°C)

Density at 25°C (g cm3)

Viscosity at 25°C (cP)

Conductivity at 25°C (mS cm1)

Ref.

EMImAlCl4

8

1.29

18

22.6

[14,15]

EMIm(AlCl3)2Cl

–

1.39

14

15.4

[14,15]

EMImNO2

55

1.27

–

–

[1]

EMImNO3

38

1.28

–

–

[1]

EMIm(CN)2N

12

1.08

17

27

[44]

EMIm(CN)3C

11

1.11

18

18

[45]

51

1.26

65

1.14

5

100

[11]

EMImBF4

15

1.28

32

13.6

[1,26]

EMImPF6

62

1.56

–

–

[47]

EMImAsF6

53

1.78

–

–

[20,48]

EMImSbF6

10

1.85

67

6.2

[20,49]

EMImNbF6

1

1.67

49

8.5

[19,20]

EMImTaF6

2

2.17

51

7.1

[19,20]

EMImWF7

15

2.27

171

3.2

[20]

EMImCF3CO2

14

1.29

35

9.6

[7]

EMImCF3SO3

10

1.38

43

9.3

[2,7,30]

EMIm(CF3SO2)2N

15

1.52

28

8.4

[7]

EMIm(CF3CO)(CF3SO2)N

2

1.46

25

9.8

[23]

EMIm(CF3SO2)3C

39

–

181

1.7

[7,30]

EMImFHF EMI(HF)2.3F

[46]

Although conductivities of organic electrolytes depend on the solute, solvent and concentration, they fall roughly in the order of 100–102 mS cm1. Therefore, one can say that conductivities of RTMSs are similar to or a little lower than those of organic electrolytes. The lower conductivities of these nonaqueous electrolytes are mostly caused by their high viscosities. Exceptionally high conductivities are found for some dialkylimidazolium fluorohydrogenates (⬃102 mS cm1).

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355

Fig. 3. Relationship between the size of anion and melting point for EMImMFm1. The horizontal axis gives the value of the radius of the central atom of the anion reported by Shannon [22] and the radius of seven-coordinated Mo(VI) is used for that of WF7 [19,20,22].

A simple inversely proportional relationship, which is known as Walden’s rule, is also observed between the molar conductivities and viscosities of imidazolium RTMSs,

λη  constant

(8)

where λ is the molar conductivity and η is the viscosity. Fig. 4 shows a Walden plot where the reciprocal molar conductivities of RTMSs composed of dialkylimidazolium cations and fluoroanions are plotted against their viscosities (DMIm, 1,3-dimethylimidazoium; PrMIm, 1-methyl-3-propylimidazolium; BMIm, 1-butyl-3-methylimidazolium; PeMIm, 1-methyl-3-pentylimidazolium; HMIm, 1-hexyl-3-methylimidazolium). Since the values are distributed over two orders of magnitude, both the axes are shown in logarithmic scales and a linear relationship is observed [11]. The viscosity is the most decisive factor for the conductivity of these molten salts. It should be noted that the plots of 1-alkyl3-methylimidazolium fluorohydrogenates (RMIm(HF)2.3F) also appear on the same line. As stated above, RMIm(HF)2.3F possess relatively high conductivities in RTMSs that arises from their low viscosities of these salts. The contribution of proton hopping among the fluorohydrogenate anions is ruled out based on the result of the pulsed-gradient spin-echo (PGSE) NMR studies of these molten salts, where the contribution of each ionic species is separately determined [25].

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Fig. 4. Walden plot of room-temperature molten salts composed of alkylimidazolium cation and fluoroanions including the present salts: (1) EMIm(HF)2.3F, (2) DMIm(HF)2.3F, (3) PrMIm(HF)2.3F, (4) BMIm(HF)2.3F, (5) PeMIm(HF)2.3F, (6) HMIm(HF)2.3F, (7) 1,3diethylimidazolium bis(trifluoromethylsulfonyl)amide, (8) EMImBF4, (9) DMIm(CF3SO2)2N, (10) 1-ethyl-3,4-dimethylimidazolium bis(trifluoromethylsul-fonyl)amide, (11) 1,3-dimethyl4-methylimidazolium bis(trifluoromethylsulfonyl)amide, (12) EMImCF3CO2, (13) 1,3-diethylimidazolium triflate, (14) 1,3-diethylimidazolium trifluoromethylcarboxylate, (15) 1-ethyl-3, 4-dimethylimidazolium triflate, (16) BMIm(CF3SO2)2N, (17) 1-etoxymethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide, (18) 1-ethyl-2,3-dimethylimidazolium bis(trifluoromethylsulfonyl)amide, (19) BMImCF3SO3, (20) 1-isobutyl-3-methylimidazolium bis (trifluoromethylsulfonyl)amide, (21) BMImCF3CO2, (22) 1-buty l- 3-ethylimidazolium trifluoromethylcarboxylate, (23) BMImCF2CF2CF3CO2, (24) 1-(2,2, 2-trifluoromethyl)-3-methylimidazolium bis(trifluoromethylsulfonyl)amide, (25) 1-butyl-3-ethylimidazolium perfluorobutylsulfate, and (26) BMImCF3CF2CF2CF2SO3. Viscosities and conductivities of the salts are obtained in the present study or summarized in the Ref. 6 and 21.

Temperature dependence of the conductivities and viscosities of RMIm(HF)2.3F is shown in Figs. 5 and 6. In general, the conductivities and viscosities of highly viscous RTMSs do not obey the simple Arrhenius’ law, and the plots of logarithmic value of conductivities and viscosities against the inverse of absolute temperature curve convexly and concavely, respectively (Figs. 7 and 8). In such a case, Vogel–Tamman–Fulcher (VTF) equation is usually employed instead of expressing the temperature dependence [26]:

κ  κ0T 1/2 exp [B / (TT0)]

(9)

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357

Fig. 5. Arrhenius plots of the conductivities of RMIm(HF)2.3F [11].

Fig. 6. Arrhenius plots of the viscosities of RMIm(HF)2.3Fs [11].

η  η0T 1/2 exp [B / (TT0)]

(10)

where κ and η are conductivity and viscosity, respectively. B and T0 are empirically determined constants, the latter being called an ideal glass transition temperature. RMIm(HF)2.3F whose conductivities and viscosities obey Arrhenius’ law over a wide temperature range is an exceptional case as a result of their extremely low viscosities. 3.3. Electrochemical window

A wide electrochemical window of RTMSs is one of the most important factors required for their applications as electrolytes of electrochemical devices

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Fig. 7. Arrhenius plots of the conductivities of selected RTMSs [7].

Fig. 8. Arrhenius plots of the viscosities of selected RTMSs [7].

with high power and energy. For aqueous solutions, theoretical electrochemical window is 1.23 V at the standard state. No drastic extension is expected, although overvoltages are more or less observed in practice depending on the electrolyte and electrode materials. On the other hand, organic solutions such as propylene carbonate and acetonitrile possess high electrochemical stabilities, leading to the electrochemical windows of more than 5 V. The electrochemical window of RTMS is usually measured by cyclic voltammetry using platinum or glassy carbon electrodes, as in the case of organic electrolytes. Fig. 9 shows cyclic voltammograms of platinum or glassy carbon electrodes in various RTMSs, where

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359

Fig. 9. Electrochemical windows of some RTMSs measured by cyclic voltammetry. W.E., glassy carbon for EMPyr(HF)2.3F and EMIm(HF)2.3F; and platinum for EMImTaF6 and EMImBF4 [11,12,19,20].

EMPyr denotes N-ethyl-N-methylpyrrolidinium cation. As mentioned in Section 2.1, nonaromatic cations such as aliphatic or nonaromatic heterocyclic ammonium cations are electrochemically more stable than aromatic cations. Fluoroanions such as (CF3SO2)2N and BF4, known as counteranions of supporting electrolytes of lithium battery and electric double-layer capacitor (EDLC), are also employed for the anions of RTMSs with high electrochemical stability. For instance, electrochemical window of N-butyl-N-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide (BMPyr(CF3SO2)2N) is reported to be about 6 V [27], whereas electrochemical windows of EMImBF4 and EMIm(CF3SO2)2N are about 4.5 V [7,20]. For fluorohydrogenate salts, cathodic limits are restricted by H2 evolution from (HF)nF anion and reduction of cation. The reaction occurring at the anodic limits are probably fluorination of the cations. Use of more stable alkylpyrrolidinium cations extends the anodic limit giving electrochemical windows of approximately 5 V [12]. 4. APPLICATION OF ROOM-TEMPERATURE MOLTEN SALTS TO ELECTROCHEMICAL DEVICES 4.1. Electric double-layer capacitor

The basic energy-storage mechanism of EDLC is not Faradaic, but the adsorption of the ions on electrodes made of materials with large surface areas such as activated carbon. A fast charge – discharge process gives a high power density. For the electrolyte of EDLC, a high conductivity and wide electrochemical window are required to realize high power and energy densities. Some reports on the application of RTMSs to EDLC are known today: EMIm(HF)2.3F [28], EMImBF4 [28–30], EMImNbF6 [29], EMImTaF6 [29], EMImCF3SO3

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[30,31], EMIm(CF3SO2)2N [30,31], EMIm(CF3SO2)3C [30], BMImBF4 [31], BMImPF6 [31] and BMPyr(CF3SO2)2N [31]. Nanjundiah et al. [30] reported capacitances of various RTMSs on a dropping mercury electrode with those of other electrolytes. There is a rough tendency for the capacitance in the sequence: aqueous solutions  RTMSs  organic solutions. Another important observation is that the capacitance (μF cm2) on Hg, glassy carbon and activated carbon are essentially the same, if the scan rate is chosen slow enough. Lewandowski et al. [31] prepared and tested a series of electrochemical capacitors based on activated carbon powders and RTMSs as electrolytes. The capacitances obtained for EMImBF4, BMImBF4, BMImPF6, EMIm(CF3SO2)2N and BMPyr(CF3SO2)2N are not significantly different (5.2–6.3 μF cm2 when activated carbon with the surface area of 870 m2 g1 is used). They pointed out that the existence of Faradaic processes such as redox reactions involving surface functional groups on the activated carbon or electrolyte impurities. EMIm(HF)2.3F with a high conductivity was also applied to EDLC [28]. Table 2 shows the comparison of some physical properties of electrolytes used for EDLC. Capacitance and resistance were measured by using coin type cells with activated carbon electrodes (coconut-shell charcoal). The capacitor using EMIm(HF)2.3F as the electrolyte possesses much lower resistance and larger capacitance than that of EMImBF4. However, as described in Section 3.3., the electrochemical window of EMIm(HF)2.3F is about 3 V and smaller than those of other RTMSs. In the practical EDLC with activated carbon electrodes the cell voltage in which charge process is stably performed is upto 2 V. The application of alkylpyrrolidinium and alkylpiperidinium fluorohydrogenates with larger electrochemical windows as electrolytes of EDLC is now under investigation. 4.2. Fuel cells

In the case of polymer electrolyte fuel cell (PEFC) using perfluoroalkylsulfonate membranes such as Nafion®, protons produced at the anode migrate in the Table 2 Physical properties of electrolytes and cell performances in EDLCs cells at 25°C, and 0.8 V [28] Electrolyte

Conductivity (mS cm1)

Viscosity (cP)

Capacitance (F cm3)

Resistance (Ω)

EMIm(HF)2.3F

100

4.9

11.1

6.6

EMImBF4

13

43

6.8

19.8

H2SO4 (35 wt%)

848

2.5

23.2

4.6

Et3MeNBF4/PC (1 M)

13

3.5

8.1

11.8

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361

electrolyte to react with oxygen, forming water at the cathode. Humidification of the electrolyte membrane is necessary to keep the ionic conduction of it, limiting the operation temperature of the system lower than 80°C. Operation of PEFCs at elevated temperatures, 100–200°C, enables reduction in the size of the cooling system and improvement of polarization behaviors, leading to the reduction of the use of costly precious metal catalysts. RTMSs are attractive candidates to realize the operation of the fuel cell at elevated temperatures. Some protic acids also have been applied for this purpose. Susan et al. [32] reported the combination of various amines and HN(SO2CF3)2. Although not many of the equimolar conditions of the amine and HN(SO2CF3)2 are liquids at room temperature, most of them melt below 130°C and their conductivities are of the order of 101 mS cm1 at 130°C. Details on the fuel cell operation using imidazole under the same concept have been reported separately [33]. The equimolar compound ImH(CF3SO2)2N (ImH, imidazolium) exhibits a high melting point of 73°C. Both in the imidazole- and HN(SO2CF3)2-rich composition ranges, the mixtures exhibit lower melting points than room temperature. Interestingly, imidazole-rich compositions exhibit higher conductivities. These high conductivities are explained by protonhopping conduction mechanism from Im to ImH. These systems are electrochemically active for H2 oxidation and O2 reduction at a Pt electrode under nonhumidifying conditions. The Pt electrodes immersed in EMIm(HF)2.3F respond to H2 and O2 gases to exhibit stable electrode potential. Using EMIm(HF)2.3F as an electrolyte, a fuel cell of a new concept has been proposed in which hydrogen is transferred via fluorohydrogenate conduction (Fig. 10) [34]. The cell exhibits an open-circuit voltage of higher than 1 V between H2 and O2 gas electrodes in EMIm(HF)2.3F at 25°C. Polarization behaviors of H2 and O2 electrodes in EMIm(HF)2.3F at 25°C are shown in Fig. 11. A significantly small overpotential of H2 electrode in EMIm(HF)2.3F is found compared with that in 30 wt% KOH aqueous solution, whereas the overpotential of O2 gas electrodes are comparable with each other. Since EMIm(HF)2.3F has HF dissociation pressure at higher temperatures,

Fig. 10. Principle of the fuel cell of EMIm(HF)2.3F [34].

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Fig. 11. Polarization behavior of (a) H2 and (b) O2 gas electrodes in EMIm(HF)2.3F, EMIm(HF)1.3F and 30 wt% KOH at 25°C [34].

Fig. 12. Polarization behavior of (a) H2 and (b) O2 gas electrodes in EMIm(HF)1.3F at 25 and 100°C, and 30 wt% KOH at 25°C [34].

EMIm(HF)1.3F, obtained by eliminating HF at 100°C, was employed as an electrolyte for the operation at 100°C. Open-circuit voltages of the cell using EMIm(HF)1.3F as the electrolyte at 25 and 100°C are again higher than 1 V. Polarization behaviors of H2 and O2 electrodes in EMIm(HF)1.3F at 25 and 100°C are shown in Fig. 12. For comparison with EMIm(HF)2.3F, polarization behaviors of H2 and O2 electrodes in EMIm(HF)1.3F at 25°C are also shown in Fig. 11. Since the rearrangement of H–F bonds in (HF)nF are involved in the electrode reactions, overpotentials of both the electrodes in EMIm(HF)1.3F are higher than those in EMIm(HF)2.3F due to the stronger (HF)2–F bond in EMIm(HF)1.3F than (HF)3–F bond in EMIm(HF)2.3F. However, polarization behaviors of both the electrodes are significantly improved at 100°C compared with those at 25°C.

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363

4.3. Lithium battery

Application of RTMSs to lithium battery started at an early stage of research on RTMSs [35]. The main purpose is to use them as substitutes for the organic solvents and improve the battery safety by taking advantage of the inflammable nature. RTMSs of BF4 and (CF3SO2)2N, combined with lithium salts of the same anions, have been mostly studied as shown in Table 3 [36,37]. Applications of the EMIm-based RTMSs to the electrolytes of the lithium secondary batteries have been reported by several authors; however, the weakness against reduction is unfavorable for this purpose. Cell performances of lithium secondary batteries using EMIm, TMPA (trimethylpropylammonium), PMPyr and PMPip salts, combined with (CF3SO2)2N, were examined using two-electrode cells. Comparison of the electrochemical windows is shown in Fig. 13 [38]. Here, Li(CF3SO2)2N was dissolved in the RTMS as a supporting electrolyte, and Li metal and LiCoO2 were used for negative and positive electrode materials, respectively. Cell test using PMPip(CF3SO2)2N exhibits the best performance among them, which is due to the high stability of PMPip cation against reduction. There are mainly two ways to use RTMSs, avoiding the cathodic reduction of them. One is the additives forming the solid electrolyte interface (SEI) film on the cathode that kinetically prevents the cathodic reduction; thionyl chloride is often used for this purpose [39]. The alternative solution is to employ other cathodes possessing a higher reduction potential instead of lithium or LiC6. Some lithiated metal oxides are used for this purpose. One of the examples will be described below. Another important problem to be solved is a high inner resistance of the cell. Nonaromatic cation-based salts are usually highly viscous to

Table 3 Lithium batteries using RTMSs of fluoroanions Electrolyte

Cathode material

Anode material

Ref.

EMIm BF4–LiBF4

LiCoO2

Li4Ti5O12

[36,40]

EMIm BF4–LiBF4

LiCoO2

Li-Al

[50]

TMPA (CF3SO2)2N–Li(CF3SO2)2N

LiCoO2

Li

[38]

PMPip (CF3SO2)2N–Li(CF3SO2)2N

LiCoO2

Li

[38]

PMPip (CF3SO2)2N–Li(CF3SO2)2N

Sn

Li

[38]

TEA (CF3CO)(CF3SO2)N–Li(CF3SO2)2N

LiCoO2

Li

[37]

EMIm (CF3SO2)2N–Li(CF3SO2)2N

LiCoO2

Li

[37]

EMIm (CF3CO)(CF3SO2)N–Li(CF3SO2)2N

LiCoO2

Li

[37]

Note: TMPA, trimethylpropylammonium; PMPip, N-propyl-N-methylpiperidinium; TEA, tetraethylammonium.

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Fig. 13. Electrochemical windows of some RTMSs for electrolytes of lithium batteries measured by linear sweep voltammetry. W.E., glassy carbon; C.E., Pt; R.E., Pt [23].

have low conductivities. Moreover, addition of lithium-supporting electrolyte also increases the viscosity of the electrolyte. Nakagawa [36,40] reported physical properties of the salt mixtures of EMImBF4 and LiBF4. The conductivity of the system decreases with the increase in the content of LiBF4. EMImBF4 containing 1.5 mol dm3 of LiBF4 exhibits the conductivity of higher than 100 mS cm1 that is still applicable to lithium battery. Melting point of the system also decreases with the increase in the amount of LiBF4 in the mixture. Thermal decomposition of EMImBF4 – LiBF4 mixtures occurs at much higher temperatures than that of ethylenecarbonate (EC)/γ -butyloractone (GBR). EC/GBR containing 1 mol dm3 LiBF4 releases the organic solvents at 100°C and decomposes at around 200°C, whereas EMImBF4 containing 1 mol dm3 LiBF4 is stable even at 300°C. Cell test of EMImBF4/ LiBF4 electrolyte was performed using Li4Ti5O12 as an negative electrode material. Because the intercalation/deintercalation potential of Li4Ti5O12 is around 1.5 V vs. Li/ Li, even EMIm cation-based salts, which are weak against reduction, can be used as electrolytes. The open-circuit voltage of the cell at the end of 1 h rest time before the third discharge with a current of 25 μA cm2 was 2.61 V. This cell showed good performance as long as for the low rate charge and discharge are performed. Lithium cation, small in size, is a strong Lewis acid and strongly interacts with the counteranion of the salt, leading to high melting point and viscosity. Fujinami and Buzoujima [41] reported a room-temperature molten lithium salt combined with a large perfluorotetraalkoxy-alkylaluminate anion that has the aluminum center bound to four oxygen atoms (Fig. 14). Two of the oxygen atoms are bound to fluoroalkyl groups and the others are bound to long ether chains. This weakly coordinating anion functions well to lower the melting point and promotes the dissociation of the cation and anion. Unfortunately, the viscosity of this Li-based RTMS is still very high and should be overcome in the next step for practical applications.

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Fig. 14. Structures of room-temperature molten lithium salts using large aluminate anions [41].

4.4. Solar cells

RTMSs are involatile, which is an important property for use in solar cells. Grätzel’s group [42] originally invented (CF3SO2)2N–based salts to prevent the loss of the electrolyte solvent by vaporization in their dye-sensitized solar cells. They also found that 1-hexyl-3-methylimidazolium iodide (HMImI) is a liquid at room temperature and forms the I/I3 redox couple when mixed with I2. The output of a solar cell is approximately proportional to the product of its electromotive force and maximum current output. The electromotive force is governed by the energy levels of the materials, and the current depends on the diffusion rate of the redox species when electron conduction in the semiconductor is fast. Therefore, RTMSs with low viscosities are preferable to facilitate the diffusion of the redox species.

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Fig. 15. I–V characteristics of a cell using EMIm(HF)2.3F and EMIm(CF3SO2)2N containing 0.9 M of DMHImI and 30 mM of I2 [43].

Cell performances of solar cells were compared using EMIm(HF)2.3F and EMIm(CF3SO2)2N as the electrolytes [43]. DMHImI (1,2-dimethyl-3-hexylimidazolium iodide) and I2 were used for the source of the I/I3 redox couple. From the I–V characteristic under illumination shown in Fig. 15, the maximum current density observed for EMIm(HF)2.3F is about six times higher than that of EMIm(CF3SO2)2N. This is due to the difference between the diffusion rates of I and I3 caused by the difference of the viscosity in the two RTMSs (5 and 28 cP, respectively). The diffusion coefficient of I3 in EMIm(HF)2.3F, calculated from the current density of the reduction peak in the cyclic voltammogram using a Pt electrode, is 4.3106 cm2 s1, whereas that in EMIm(CF3SO2)2N is 7.6107 cm2 s1. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9]

J.S. Wilkes and M.J. Zaworotko, J. Chem. Soc. Chem. Commun., (1992) 965. E.I. Cooper and E.J.M. O’Sullivan, Proc. of Eight Int. Symp. on Molten Salts, R. J. Gale, G. Blomgren, and H. Kojima (Eds.), The Electrochemical Society, Pennington, NJ, 1992, p.386. T. Welton, Chem. Rev., 99 (1999) 2071. K.R. Seddon, J. Chem. Technol. Biotechnol., 68 (1997) 351. P. Wasserscheid and W. Kein, Angew. Chem. Int. Ed., 39 (2000) 3772. R. Hagiwara and Y. Ito, J. Fluorine Chem., 105 (2000) 221. P. Bonhôte, A.-P. Dias, M. Armand, N. Papageorgiou, K. Kalyanasundaram, and M. Grätzel, Inorg. Chem., 35 (1996) 1168. R.P. Singh, S. Manandhar, and J.M. Shreeve, Tetrahedron Lett., 43 (2002) 9497. R. Hagiwara, T. Hirashige, T. Tsuda, and Y. Ito, J. Fluorine Chem., 99 (1999) 1.

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[10] R. Hagiwara, T. Hirashige, T. Tsuda, and Y. Ito, J. Electrochem. Soc., 149 (2002) D1. [11] R. Hagiwara, K. Matsumoto, Y. Nakamori, T. Tsuda, Y. Ito, H. Matsumoto, and K. Momota, J. Electrochem. Soc., 150 (2003) D195. [12] K. Matsumoto, R. Hagiwara, and Y. Ito, Electrochem. Solid State Lett., 7 (2004) E41. [13] M. Hirao, H. Sugimoto, and H. Ohno, J. Electrochem. Soc., 147 (2000) 4168. [14] J.S. Wilkes, J.A. Levisky, R.A. Wilson, and C.L. Hussey, Inorg. Chem., 21 (1982) 1263. [15] A.A. Fanning, Jr., D.A. Floreani, L.A. King, S.S. Landers, B.J. Piersma, D.J. Stech, R.L. Vaughn, J.S. Wilkes, and J.L. Williams, J. Phys. Chem., 88 (1984) 2614. [16] J.A. Boon, J.S. Wilkes, and J.A. Lanning, J. Electrochem. Soc., 138 (1991) 465. [17] K.O. Christe, W.W. Wilson, R.D. Wilson, R. Bau, and J.A. Feng, J. Am. Chem. Soc., 112 (1990) 7619. [18] R.P. Swatloski, J.D. Holbrey, and R.D. Rogers, Green Chem., 5 (2003) 361. [19] K. Matsumoto, R. Hagiwara, and Y. Ito, J. Fluorine Chem., 115 (2002) 133. [20] K. Matsumoto, R. Hagiwara, R. Yoshida, Y. Ito, Z. Mazej, P. Benkic E, B. Cemva, O. Tamada, H. Yoshino, and S. Matsubara, Dalton Trans., (2004) 144. [21] R. Hagiwara, Electrochemistry, 70 (2002) 130. [22] R.D. Shannon, Acta Cryst., A32 (1976) 751. [23] H. Matsumoto, H. Kageyama, and Y. Miyazaki, Chem. Commun., (2002) 1726. [24] H. Sakaebe and H. Matsumoto, Electrochem. Commun., 5 (2003) 594. [25] Y. Saito, K. Hirai, K. Matsumoto, R. Hagiwara, Y. Ito, and Y. Minamizaki, J. Phys. Chem. B, 109 (2005) 2942. [26] A. Noda, K. Hayamizu, and M. Watanabe, J. Phys. Chem. B, 105 (2001) 4603. [27] D.R. MacFarlane, P. Meakin, J. Sun, N. Amini, and M. Forsyth, J. Phys. Chem. B, 103 (1999) 4164. [28] M. Ue, M. Takeda, A. Toriumi, A. Kominato, R. Hagiwara, and Y. Ito, J. Electrochem. Soc., 150 (2003) A499. [29] M. Ue, M. Takeda, T. Takahashi, and M. Takehara, Electrochem. Solid State Lett., 5 (2002) A119. [30] C. Nanjundiah, S.F. McDevitt, and V.R. Koch, J. Electrochem. Soc., 144 (1997) 3392. [31] A. Lewandowski and M. Galinski, J. Phys. Chem. Solids, 65 (2004) 281. [32] M.A.B.H. Susan, A. Noda, S. Mitsushima, and M. Watanabe, Chem. Commun., (2003) 938. [33] A. Noda, M.A. B.H. Susan, K. Kudo, S. Mitsushima, K. Hayamizu, and M. Watanabe, J. Phys. Chem. B, 107 (2003) 4024. [34] R. Hagiwara, T. Nohira, K. Matsumoto, and Y. Tamba, Electrochem. Solid State Lett., in press. [35] J. Devynck, R. Messina, J. Pinggarron, B. Tremillon, and L. Trichet, J. Electrochem. Soc., 131 (1984) 2274. [36] H. Nakagawa, S. Izuchi, K. Kuwana, T. Nukuda, and Y. Aihara, J. Electrochem. Soc., 150 (2003) A695. [37] H. Matsumoto and H. Sakaebe, Polym. Prepr. Jpn., 51 (11) 2758 (2002). [38] H. Sakaebe and H. Matsumoto, Electrochem. Commun., 5 (2003) 594. [39] N. Koura, K. Etoh, Y. Idemoto, and F. Matsumoto, Chem. Lett., (2001) 1320. [40] H. Nakagawa, Youyuuenn Oyobi Kouonnkagaku, 47 (1) (2004) 19. [41] T. Fujinami and Y. Buzoujima, J. Power Sources, 119–121 (2003) 438. [42] N. Papageorgiou, Y. Athanassov, M. Armand, P. Bonhôte, H. Pettersson, A. Azam, and M. Grätzel, J. Electrochem. Soc., 143 (1996) 3099. [43] H. Matsumoto, T. Matsuda, T. Tsuda, R. Hagiwara, Y. Ito, and Y. Miyazaki, Chem. Lett., (2001) 26. [44] D.R. MacFarlane, J. Golding, S. Forsyth, M. Forsyt, and G.B. Deacon, Chem. Commun., (2001) 1430.

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[45] Y. Yoshida, K. Muroi, A. Otsuka, G. Saito, M. Takahashi, and T. Yoko, Inorg. Chem., 43 (2004) 1458. [46] K. Matsumoto, T. Tsuda, R. Hagiwara, Y. Ito, and O. Tamada, Solid State Sci., 4 (2002) 23. [47] J. Fuller, R.T. Carlin, H.C. De Long, and D. Haworth, J. Chem. Soc., Chem. Commun., (1994) 299. [48] A.B. McEwen, H.L. Ngo, K. LeCompte, and J.L. Goldman, J. Electrochem. Soc., 146 (1999) 1687. [49] C.H. Song, W.H. Shim, E.J. Roh, and J.H. Choi, J. Chem. Soc., Chem. Commun., (2000) 1695. [50] R.T. Carling, J. Fuller, W.K. Kuhn, J. Lysaght, and P.C. Trulove, J. Appl. Electrochem., 26 (1996) 1147.

Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 17

Fluorine-intercalated graphite for lithium batteries A. Hamwi, K. Guérin, and M. Dubois Laboratoire des Matériaux Inorganiques, Université Blaise Pascal de ClermontFerrand, UMR CNRS-6002, 63177 Aubière, France 1. INTRODUCTION: THERMODYNAMIC CONSIDERATIONS The very high value of the Gibbs free energy of formation of LiF (ΔfGo  587 kJ mol1), converted into electrochemical energy, potentially provides a very high emf of about 6.08 V (ΔfGo  nFEo, where n  1, F is the Faraday constant  96486 C mol1 and emf  Eo) for the galvanic cell obtained by combining F2 gas as the positive electrode and Li metal as the negative electrode: Li → Li  e 1  2 F2

 e → F

Li  12 F2 → LiF The calculated mass coulombic capacity is about 1410 mAh g1 of positive fluorine electrode and energy density is about 8576 Wh kg1. In practice, the usable part of this energy is smaller because of the energy loss due to many energetic phenomena developing in the cell during the discharge process (the so-called internal resistance), for example, those related to fluorine’s physical state in the positive electrode. Except for the gaseous state, the latter phenomena are considered as an important energy-consuming system. Unfortunately, using fluorine in the gas phase is, in practice, impossible not only because of its very high reactivity with most materials even at room temperature, but also because of its presence in a very small amount in the gas phase which, consequently, produces a very low energy in the system. In other words, the higher the amount of fluorine the higher the coulombic capacity. Thus, the use of a solid-state support for fluorine storage is required. Carbon materials can be recommended for two reasons: firstly, its lightness in order to achieve a high specific energy density of the

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cathode and, secondly, the large variety of bonds with different energy values between carbon atoms and fluorine atoms. Indeed, carbon reactivity with fluorine varies depending on the solid-carbon state forms, from amorphous carbon to highly crystallized forms such as graphite. Consequently, the carbon–fluorine interaction strength varies from very weak bond energies, as in the case of fluorine adsorbed on the carbon material surface, to very high bond energies such as the C–F covalent bond (the carbon atom is in sp3 hybridization state), to the C–F ionic bond of moderate energy. In the latter case, the ionic bond is between the F anion and some of carbon atoms (remaining in sp2 hybridization state) symbolizing the cation as in graphene layers of the graphitic structure. Generally, regardless of the carbon allotropic forms and the synthesis methods of the compounds with fluorine that will be described hereafter, the increase in the fluorine storage level generates an increase in the C–F bond energy. However, taking into account this C–F interaction energy consumption from the LiF formation energy, the energy that has to be converted produced by the electrochemical system, is less than that provided (previously calculated). In other words, the highperformance electrochemical system requires highly fluorinated materials with low C–F interaction energy. In general, the C–F bond dissociation energy (BDE) increases as the number of fluorine atoms bound to the same carbon atom in the given molecules increases (460, 497, 526, and 545 kJ mol1 for H3CF, H2CF2, HCF3, and CF4, respectively; Fig. 1 [1,2]. The C–F bond is highly polar and therefore, a large ionic contribution can be expected (0.4e). Calculations performed on fluoromethanes [3] predict an increase in the charge transfer from carbon to fluorine as the number of fluorine atoms bound to the carbon atom increases. Furthermore, the C–F bond strength increases from H3CF to CF4 while the bond length decreases (from 1.39 to 1.32 Å, respectively; Fig. 1) and then bond energies increase with the number of fluorine atoms in the molecule. The ionic C–F bond dissociation energy, supposed to be weak, was estimated by Di Vittorio et al. to be about 54 kJ mol1 at the most, which exists only in diluted fluorine – graphite intercalation compounds [4]. As fluorination temperature increases, both fluorine concentration and covalent character increase. Generally, solid-state carbon fluorides CFx are prepared by the direct reaction of fluorine with carbon materials. A temperature higher than 350°C is necessary if graphite and graphitized materials (for example, petroleum coke heat treated at 2800°C) are used. The higher the reaction temperature, the higher the fluorination level x (x  F/C, 0.5 x 1) of compounds (called CF(HT)) and the C–F covalent character, where the carbon atoms assume sp3 hybridization. Because of the decomposition of compounds, the reaction temperature does not exceed 600°C, the temperature at which the x ( F/C) value reaches 1. While amorphous or disordered (less graphitized) carbons (petroleum coke, carbon black, active carbon, etc.) [5,6] yield fluorides at all temperatures, the fluorination

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Ionic

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3.2

C-F in fluorinated graphite

3.0

2.6 Semi-ionic

2.4 2.2 2.0

C-F in Fluoromethanes

Covalent

C-F bond lenght (Å)

2.8

1.8 1.6 1.4 1.2 100

200

300 BDE

400

500

600

(kJ.mol-1)

Fig. 1. C–F bond length values as a function of their bond dissociation energy (BDE) for fluoromethanes and graphite fluorides.

level x can sometimes exceed one fluorine atom per carbon atom, indicating the formation of CF2 and CF3 groups. These latter groups should be less active than the CF group because the C–F bond energies are higher and the mass coulombic capacity is decreased. The higher the carbon reactivity (amorphous materials) and the reaction temperature, the higher the amount of CF2 and CF3 group formation. It is well known here that fullerenes (C60, C70), which are also considered as crystallized carbon forms, react easily at room temperature with fluorine to yield highly fluorinated compounds (x ⬇ 0.8), but they are not used as electrode material because of their high solubility in electrolytic solvents [7,8]. Therefore, the use of graphite (natural or artificial) and highly graphitized carbon (coke or other high-temperature heat-treated) as starting materials is more suitable because of the easy control of the reaction not only in pure fluorine atmosphere at various temperatures, but also in the presence of a catalytic atmosphere. Indeed, to enhance fluorine reactivity that further enabled the use of lower reaction temperatures, minute amounts of a volatile fluoride such as HF, AsF5, IF5, OsF6, WF6, SbF5, etc. [9,10], were introduced into the fluorine atmosphere, and fluorinated graphite compounds were thereby obtained from ambient temperature up to 100°C (called CF(LT)). All these compounds have a fluorination level (x  F/C) lower than 0.5, with the C–F bonds either ionic (weak bonding energies, x 0.25) or semi-ionic (or semi-covalent) involving stronger bonding energies, 0.25 x 0.5, but less than that corresponding to covalent ones.

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Moreover, it has been shown that irrespective of the stage number, sp2 carbon hybridization (i.e. planarity of graphene layers) is maintained despite the fluorine intercalation reaction into graphite [10,11]. Taking into account the estimated C–F lengths in the structure, previously proposed for semi-ionic compounds [12] of 1.49 and 2.00 Å, the energy value extrapolated from the C–F bond lengths and their corresponding BDEs should be 350 kJ mol1. C–F bond length in covalent CFx (C3C–F) should be close to 1.43 Å [1] (C–F BDE 460 kJ mol1) and in ionic CFx should be 3.07 Å (rC  rF  1.67  1.4) with estimated energy of about 54 kJ mol1 [4]. Unfortunately, in the latter case, the charge transfer is limited; it is at most about one electron for four carbon atoms (also x 0.25). In brief, Table 1 shows the roughly estimated values of C–F length and bond dissociation energy (BDE) and Fig. 2 indicates the usable energy part of the electrochemical system Li/CFx for different C–F bonding characters, both deduced from the experimental data [13–15]. The nature of C–F bonding in CF(LT) changes from ionic to semi-ionic with increase in fluorine content and decrease in stage number, and it has been well characterized by X-ray photoelectron spectroscopy, optical reflectivity measurement, and infrared spectroscopy. The formation of a C–F bond is demonstrated by XPS data [16]; indeed, for low fluorine content x 0.05, the binding energy value of the F1s electron is the same order as in metal fluoride with an ionic bond. When x increases up to 0.1, localization of the electron takes place due to the formation of a stronger C–F bond. This gives rise to a shift in the binding energy toward semi-ionic C–F bonding between fluorine and sp2hybridized carbon atoms. In accordance with the change of C–F bonding with respect to composition and stage number, CF(LT) varies from a metallic conductor to a semiconductor. At room temperature, the reactivity of fluorine against graphite is vastly improved by the presence of a volatile fluoride–anhydrous HF gaseous mixture [17]. The volatile fluorides MFn used were ClF3, BF3, IF5, BrF5, ReF6, WF6, Table 1 OCV of Li/CFx electrochemical system [13–15] and estimated values of C–F length and BDE for various CFx C–F bonding in CFx OCV range (V)a BDEb (kJ mol1 fluorine atoms) C–F length (Å)

Ionic

Semi-ionic

Covalent

4.5–5

3.7–4.0

3.2–3.4

150–104

230–201

278–258

⬃ 3.0

⬃ 1.5–2.0

⬃ 1.43

OCV of Li/F2  6.08 V, bBDE  F*[E°Li/F –E°Li/CF ].

a

2

x

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Covalent CFx

600

Semi-ionic CFx

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Ionic CFx

550 Energy (kJ.mol−1)

500 450 400 350 300 250 200 150 100 50 0 3.0

3.5

4.0 OCV (V)

4.5

5.0 BDE Usable energy

Fig. 2. Bond dissociation energy (BDE) and usable energy of CFx main types for 1 mol of fluorine atoms.

MoF6, and so on. At the first stage, highly fluorinated graphite compounds having the formula CFx(My ), with 0.5 x 0.9 and 0.06 y 0.02, were obtained. The interlayer distance Ic varied from 5.7 to 6.1 Å. It has been suggested that the planarity of the graphene layers is preserved and the C–F bonds have a semi-ionic character. The value of x also depends on the fluoride (MFn) used. The fluorination level is related to the Lewis acidity of the volatile fluoride and its interaction with HF. It seems that the higher values correspond to the weaker Lewis acidity of fluoride relative to that of HF [12]. The highest degree of fluorination was achieved using IF5 that is considered as a slightly weaker Lewis acid than HF. The semi-ionic character of the bond, which is considered to be intermediate between ionic and covalent, has been suggested by many physicochemical measurements (XRD, FT–IR, NMR, etc.). For example, 13 C–NMR spectroscopy has shown two carbon types: carbon atoms close to sp2 type with weak C–F interaction and carbon atoms with stronger C–F interaction. The conventional graphite fluorides and fluorine – graphite intercalation compounds are well described by many reviewers [17–20]. We will focus this chapter on a new family of compounds obtained from the fluorination posttreatment of graphite fluoride prepared at room temperature. Indeed, very recently, in order to remove the iodine-based impurities that are still present in the compounds CFx(Iy) and to achieve improved electrochemical performance, a heat treatment process at different temperatures up to 680°C under fluorine atmosphere was performed. The physicochemical characterizations formed by a

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combination of techniques (NMR, XRD, FT–IR, and EPR) and the evolution of the C–F bonding as a function of the post-treatment temperature (TFPT) will be discussed. The electrochemical performance of the resulting materials, used as electrodes in lithium batteries, will also be presented and discussed on the basis of the C–F bonding nature. 2. SYNTHESIS OF MODIFIED LOW-TEMPERATURE FLUORINATED GRAPHITE Graphite fluoride compounds were obtained by the reaction of graphite with a gaseous mixture of F2, HF, and IF5 at room temperature. The chemical composition for this product is around CF0.89I0.02H0.06. Experimental conditions were previously described [17]. RAW–CF(LT) was then post-treated under fluorine gas at temperatures between 100 and 680°C [21]. Color and composition of the post-treated products are summarized in Fig. 3. The color varies from green to white through yellow as the re-fluorination temperature increases and iodine fluoride content decreases. All these compounds are very stable in air and no attack on the glass flask is noticed. Surprisingly, no significant weight uptake was recorded during the re-fluorination of compounds even though the F/C molar ratio increased for re-fluorination temperatures higher than 400°C. This confirms the de-intercalation of the IFy species, as shown by 19F–NMR investigation and their substitution by fluorine atoms in the fluorographite matrix. The weight variation is low, except for the compound re-fluorinated at 150°C. Indeed, for this last compound, a high 8%

1.05

Brown

Green

Yellow

White 0.95

1.00 0.95

0.85

0.90

x inCFx

Ftot /C ratio

0.90

0.80

0.85

0.75

0.80 0

100

200

300

400

500

600

0.70

Fluorination post-treatment temperature(°C)

Fig. 3. Color and fluorination level of CFx for the various fluorination post-treatment temperatures (Ftot is calculated assuming that the weight variation is only due to fluorine atoms).

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weight loss is observed and is due to the loss of HF residues, which have been trapped in the graphite fluoride matrix in the form of fluoride–HF complex [12]. 3. EVOLUTION OF THE C–F BONDING DURING THE FLUORINATION POST-TREATMENT As demonstrated by several authors [18–19], NMR studies are well adapted to characterize graphite intercalation compounds and graphite fluorides because they can provide information about both the hybridization of the carbon atoms and the resonant nuclei of the intercalated species (e.g. HF molecules and IFy with y  5, 6 and 7 for IF5, IF6 and IF7, respectively). The solid-state NMR data, correlated with the data of X-ray diffraction, Fourier–transform infrared (FT–IR), and electron paramagnetic resonance (EPR) spectroscopies, are necessary to understand the processes occurring during the fluorination post-treatment of room-temperature fluorinated graphite (CF(LT)). 3.1. De-intercalation of the catalyst residues

As reported by Panich et al. [18], 19F–NMR spectrum of fluorine-GIC prepared at temperatures lower than 100°C is composed of two lines: a narrow line corresponding to free fluorine atoms and a broad line attributed to fluorine bonded to carbon atoms. The line-width of the latter is mainly due to dipolar coupling [18,19,22–24]. For diluted compounds (x 0.125) only the narrow line appears. When the fluorine content increases, the intensity of the broad line increases in comparison with the narrow one [18,22]. 19F–NMR spectra of hightemperature graphite fluorides ((CF(HT):(CF)n and (C2F)n) exhibit a broad line centered at about –180 ppm (relative to CFCl3); its width is also caused by dipolar interaction. For (C2F)n an additional weak narrow line, ascribed to weakly bound fluorine atoms was observed [18,23]. Moreover, using high-resolution 19 F–MAS and 13C–MAS associated to 19F–13C cross-polarization, the presence of CF and CF2 groups was suggested by Krawietz et al. [25]. Room-temperature fluorinated graphite (CF(LT):CFx(IFy)z) was studied mainly by static 19F–NMR [21,26,27]. These studies show a broad line (centered at about 180 ppm), attributed to C–F groups, and narrow lines at 5 and 59 ppm (denoted S1 and S2, respectively), attributed to residual IF5 groups [21,26–29] (Fig. 4). An additional narrow line is observed at 15 ppm (S3) (in the raw sample, only a shoulder appears); this chemical shift is characteristic of IF6 ion [21,26,27,30]. Although the content of the iodine fluoride species is low (I inferior to 2 at.%), their lines dominate the spectra because of their narrowness. The motional narrowing, due to the rotation of the intercalated iodine fluoride species, is responsible for the narrowness of the corresponding lines. The evolution of the NMR spectra with temperature allows the hypothesis of a translation movement of the intercalated species within the interlayer space to be rejected [27]. The

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S1

RAW CF(LT)

Fluorination post-treatment temperature

S3

400°C

S2 S4

450°C

500°C 550°C 600°C CF(HT)

400

200

0

−200

−400

−600

−800

δ / CFCl3 (ppm)

Fig. 4. Static 19F–NMR spectra of CF(HT) and of post-fluorinated CF(LT) samples for the various treatment temperatures. The spectra were recorded on a Bruker DSX300 spectrometer at 282.23 MHz.

doublet for IF5 group, i.e. S1 and S2 lines, is expected according to the molecular structure of IF5 [29]. The ratio S1/S2 of the peak surface is constant irrespective of the treatment temperature [21,27]. On the contrary, the ratio S3/S1 (i.e. SIF6 /SIF5) increases continuously indicating a thermal conversion of IF5 into IF6. These narrow lines of IF5 and IF6 disappear almost completely for TFPT  450°C: these species are de-intercalated from the fluorographite interlayers or are partly transformed into IF7 as evidenced by an additional line at 162 ppm [31] for TFPT in the range 400–500°C.

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Although IFy–GICs phases are unstable at high temperatures, they are still intercalated even at post-fluorination temperature as high as 450°C. This may be explained by the following two mechanisms of iodine fluorine trapping. (i) Due to the propagation of the re-fluorination process starting from the graphene edges and progressing toward the crystallite core. Such a mechanism traps the iodine species into the crystallite core since the fluorine atoms linked to the sp3 carbon atoms are oriented over and under the armchair carbon sheets and reduce the mobility of the iodine fluoride species within the fluorographite layers by steric hindrance (ii) Due to the inaccessibility of fluorine gas to the residual IFy–GIC phases due to their localization in the bulk. 3.2. From semi-ionic to covalent C–F bonding

On the 19F–NMR spectra (Fig. 4), the broad line centered near 170 ppm is typical of C–F bonds in fluorinated graphite compounds [18,28,32]. Its center of gravity shifts from 150 10 ppm to 180 10 ppm when the temperature of the post-treatment increases. This shift corresponds to the evolution of the covalent character of the C–F bonds in the fluorographite bulk. As a matter of fact, as also observed by Panich and Nakajima [33], the decrease in the chemical shift with increase in the treatment temperature is a consequence of the evolution of C–F bond nature from semi-ionic to covalent. Moreover, the width of this line decreases for post-treatment temperature  400°C. The line shape evolves toward that of (CF)n, which possesses a totally covalent C–F bond. The posttreatment, in fluorine gas at temperatures above 400°C, results both in the formation of covalent C–F bonds and in an enhancement of the structural order in comparison with the untreated compounds. When TFPT  400°C, the various components of the spectrum are well separated, and therefore, the spectrum of CF(LT) post-treated at 600°C displays at least four distinct lines at 70, 130, 190, and 280 ppm (Fig. 4). This underlines the different carbon–fluorine environment for re-fluorinated lowtemperature graphite fluorides. On the contrary, in the case of CF(HT), the broad 19F–NMR line could result from the superposition of several unresolved contributions. By low-frequency 19F–NMR [34], the ratio SF1/(SF1SF2)  SF1/SFtotal of the surfaces for the two characteristic lines (i.e. broad and narrow for fluorocarbon matrix (denoted SF1) and iodine species (SF2), respectively) was evaluated for samples post-treated in the range 150–450°C. Under low-frequency conditions, only one peak appears for all iodine species. The evolution of this ratio is shown in Fig. 5. The SF1/(SFtotal) ratio first slightly decreases for the sample post-treated at temperatures below 250°C, then at TFPT higher than 250°C the ratio strongly increases. At post-treatment temperatures close to 450°C, the narrow signal

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Fig. 5. Evaluation as a function of the fluorination treatment temperature of the ratio SC2/SC1 (䊊) (C1 and C2 lines are ascribed to the fluorocarbon rigid matrix and to graphitic carbons, respectively) and SF1/(SF1  SF2) (䊉) (i.e. surfaces relative to the C–F rigid matrix and to the IFy iodine species for SF1 and SF2, respectively). SF1 and SF2 are evaluated from low-frequency 19 F experiments that were recorded with a Maran Ultra (Resonance) spectrometer working at 18.9 MHz (experiments are performed at room temperature and recorded using a solid echo sequence, which allows a quantitative determination of the various contributions). The 13 C–NMR measurements were performed with a Bruker MSL300 spectrometer and a superconducting coil delivering a 7.05 T magnetic field (working frequency for 13C: 75.47 MHz).

disappears indicating that IFy groups disappear almost totally. So, the drastic increase in the ratio for TFPT  300°C is caused both by the formation of strong C–F bonds and by an important de-intercalation of IFy species. Fig. 6 displays the 13C–MAS–NMR spectra of post-treated CF(LT) that exhibit two lines, denoted C1 and C2 (with surfaces SC1 and SC2), located in the 84–88 and 135–137 ppm ranges (vs. TMS), respectively (Fig. 6). C1 line is characteristic of the fluorocarbon rigid matrix [12,18,35,36]. As proposed by several authors [33,35–37], C2 line is ascribed to graphitic type carbons. It should be noted that the chemical shift of pure graphite was measured at 119 ppm [35]. The ratio of the integrals SC2/SC1, which is still constant for TFPT 400°C, then drops for higher TFPT as shown in Fig. 5. This indicates a drastic change in

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Fig. 6. 13C–MAS–NMR spectra acquired with a spinning speed of 10 kHz of post-fluorinated samples for the various treatment temperatures.

the C–F bonding from semi-ionic (involving sp2 carbon atoms) to covalent (sp3 hybridization) when the samples are treated in the range 400–450°C. For the CF(LT) sample post-treated at TFPT  600°C and for CF(HT), the C2 line is no longer present. An increase in the post-treatment temperature does not lead to any variation in the chemical shift of the C2 line, which remains centered at 136.0 0.2 ppm irrespective of the temperature (Fig. 6). On the contrary, the chemical shift of the C1 line increases appreciably for TFPT  300°C after being constant for

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lower post-treatment temperatures. As a matter of fact, at TFPT  300°C, it changed progressively from 84.0 to 0.2 ppm, typical of semi-ionic C–F bonds to a value of 88.5 to 0.2 ppm (Fig. 7), which is characteristic of a mainly covalent C–F bond as in (CF)n compounds [28]. For TFPT between 400 and 600°C, the covalent character increases progressively in accordance with the 19F–NMR measurements. The evaluation of the 19F chemical shift for CF(LT) post-treated as a function of TFPT is also displayed in Fig. 7; recently, the study carried out by the magic angle spinning procedure [34] completed and confirmed the static 19F–NMR characterization which is described here. This evaluation of the C–F bonding from semi-ionic to covalent was also confirmed by the following XRD and FT–IR characterizations:

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(i) From XRD data (λ  1.5405 Å) [21,27], three main reflection lines centered around 14, 21 and 40.8–42.4°, were observed and assigned, respectively, to 001 [12] or 002 reflection as suggested elsewhere [12,20] (corresponding to the average interlayer spacing di ⬇ 6.00–6.60 Å), to the catalyst residual species that are still intercalated between the graphene layers (002 reflection of stage 1 IF5–GIC, d  4.24 Å [38]), and to 100 reflection of in-plane graphite

82 −220 80

0 100 200 300 400 500 600 Fluorination post-treatment temperature (˚C)

Fig. 7. Evaluation of the rigid fluorocarbon matrix with the fluorination post-treatment temperature of the 13C (䊉) (C1 in Fig. 6) and 19F (䊊) chemical shifts. δ19F is calculated from MAS–NMR experiments performed with a Bruker MSL 300 spectrometer associated to a superconducting coil giving a magnetic field of 7.049 T (working frequency for 13C and 19F : 73.4 and 282.23 MHz, respectively). A special cross polarization/magic angle spinning NMR probe for fluorine decoupling, with a 4 mm rotor was used (spinning speed of 12 kHz) [34]. Chemical shifts 13C and 19F were externally referenced to tetramethylsilane (TMS) and CFCl3, respectively.

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network (hexagonal system with corresponding reticular distances between 2.13 and 2.21 Å). The latter distances are associated to the C–C in-plane length for which the evolution as a function of the post-treatment temperature is displayed in Fig. 8. For TFPT 400°C, the C–C length (or its projection) remains close to 1.42 Å, a value close to that of pristine graphite (sp2 carbon hybridization). Beyond this strategic temperature, the C–C length increases and becomes closer to the C–C length for a sp3 hybridization (inplane projection  1.48 Å). (ii) For the re-treatment temperature below 400°C, on the basis of the FT–IR results (Fig. 9), the C–F bond mostly exhibits a semi-ionic character as evidenced by the strong band in the range 1127–1148 cm1 [17]. Its intensity continuously decreases when the re-fluorination temperature increases and it disappears above 500°C. At the same time, an additional band at 1216 cm1 attributed to covalent C–F bonding [17] appears for TFPT  350°C and increases continuously with temperature. It is necessary to note that for temperatures between 400 and 500°C, the C–F bonding exhibits semi-ionic and covalent nature. With regard to the evolution of the vibration frequencies ν (C–F) for the semi-ionic and covalent bonds as a function of the re-fluorination temperature, ν (C–F) for covalent bonds is constant at 1216 cm1 for re-fluorination temperature higher than 350°C. On the contrary, the frequency of the semi-ionic C–F shifts to the high frequencies (i.e. to a more covalent position) with an increase in temperature from 1127 to 1148 cm1 for the raw CF(LT) material and the sample post-treated at 450°C.

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Fig. 8. Evaluation of the C–C distances with the fluorination post-treatment temperature calculated from the X-ray diffraction data. The spectrum of conventional CF(HT) is added for comparison.

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Fig. 9. FT–IR spectra of room-temperature graphite fluoride post-treated at different fluorination temperatures in F2 gas. The spectrum of conventional CF(HT) is added for comparison.

3.3. Progressive fluorination limiting the conformational defects

EPR spectroscopy gives additional information on the conformational order of the post-treated materials and allows one to differentiate these samples to the conventional graphite fluoride CF(HT). Fig. 10 displays the EPR spectra of low (raw and re-fluorinated up to 600°C) and high-temperature fluorinated graphites. The origin of the main broad line (denoted as line A) was identified as carbon dangling bonds having a localized spin. Such spin carriers have been proposed for other fluorinated carbons obtained under F2 atmosphere : starting from natural graphite [24], amorphous carbon thin film [39,40] or nanosized graphite fluorides [41].

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Fig. 10. EPR spectra of the CF(LT) post-treated in F2 gas at various temperatures; the spectra of two CF(HT) obtained by direct fluorination at 600°C of natural graphite and of graphitized coke are also displayed. EPR spectra were recorded at room temperature using a X Band Bruker EMX spectrometer operating at 9.653 GHz.

Irrespective of the sample, the g-factor, which is typical of free radicals and localized structural defects, is close to 2.003 0.002. The narrow signal (line B) was ascribed to the spin carriers formed in the fluorographite layers to

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accommodate the charge of IF6 in a classical intercalation model. The broad EPR line cannot be simulated by a pure Lorentzian or Gaussian profile. It results from an unresolved hyperfine structure of the dangling bond electrons interacting with the neighboring fluorine nuclei. As a matter of fact, the wings of this structure can be observed for TFPT 200°C (Fig. 10). Dangling bond centers could adopt different configurations of surrounding fluorine atoms; this disorder leads to an unresolved super hyperfine structure (SHFS). During the re-fluorination treatment, this disorder is significantly lowered because the fluorine content increases and the degree of crystallinity is still high as observed from XRD data. Then, the local environment of each dangling bond in the highly fluorinated carbon matrix becomes more and more organized. Therefore, for fluorination posttreatment temperatures higher than 400°C, the super hyperfine structure with the fluorine nuclei becomes more and more resolved as TFPT increases (Fig. 10) and the broad line splits into seven lines. As shown by Panich et al. [24], the hyperfine interaction between dangling bond electron and six neighboring fluorine 1 nuclei (nuclear spin I  2 , n is the number of 19F nuclei) results in the splitting of the EPR spectrum into seven lines, i.e. (2nI  1)  7. This environment of the residual dangling bonds can be described as follows: it could consist of three fluorine atoms linked to three carbons in the armchair carbon sheet containing the dangling bond and three others in the adjacent sheet (the simulation of the signal for TFPT  550°C leads to the hyperfine parameter A  45 2 G, a linewidth ΔΗpp  36 G 2 G and g  2.003 0.001). Contrary to the sample re-treated at 600°C, another type of spin carrier (defects with a different environment) is present only in CF(HT) as evidenced by an additional line which is present on the EPR spectrum of these materials, denoted as line C with ΔHpp  20 1 G. The relative intensity of this additional line in comparison with that of the dangling bonds is dependent on both the starting graphite material and its fluorination process. This is emphasized by the EPR spectra of CF(HT) obtained starting from natural graphite or graphitized coke. In both cases, the SHFS is present but the intensity of line C is higher for the natural graphite than for graphitized coke (Fig. 10). Line C suggests the presence of additional conformational defects in the armchair carbon sheets. The spin density decreases as a function of the fluorination post-treatment [27] because of the reaction of the dangling bonds with F2 during the post-treatment: this density in the raw CF(LT) is high (close to 1020 2  1019, spins g1 corresponding approximately to 1 spin carrier for 200 carbon atoms), whereas after a treatment of CF(LT) at 600°C in F2 gas, it drops to a value close to 1  1018 2  1017 spins g1. It should be noted that the spin density of CF(HT) (resulting from graphitized coke) is equal to 12  1018 2  1018 spins g1 indicating a greater proportion of conformational defects than in the CF(LT) post-treated at this particular temperature.

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The evaluation of EPR experiments with temperature is different for lines A and B: the behavior of the latter is similar to that of weakly fluorinated CF(LT) [42] and other acceptor GICs [43,44], since its linewidth ΔHpp increases with decrease in temperature [27]; dipole–dipole interaction (proposed by Di Vittorio et al. [42]) and spin–orbit relaxation mechanism (Davidov et al. [43]) are involved in the relaxation of these materials. On the contrary, the width of line A is temperature-independent in the range 110–300 K with spin carriers behaving as isolated spins. Motional narrowing due to hopping motion of the dangling bond electrons at low temperature and/or due to surrounding fluorine motion at high temperature was proposed by Yokomichi et al. [39,40] for fluorinated amorphous carbon films. This phenomenon is not efficient in posttreated CF(LT). In brief, during post-treatment in pure F2 atmosphere of room-temperature graphite fluoride, C–F bonds have been formed exhibiting a bonding character more and more covalent with an increase in post-treatment temperature; the nature of the C–F bonding changes progressively from semi-ionic to covalent. The planarity of the carbon sheets is maintained (sp2 carbon hybridization) for TFPT 450°C. When TFPT is increased between 400 and 500°C, the C–F bonds develop a hybrid structure (both semi-ionic and covalent C–F bonds coexist) as the carbon skeleton that consists of chair-type structure along with some planarity. At the highest post-treatment temperatures (550–600°C), each carbon atom is then covalently bonded to a fluorine atom thereby increasing the fluorine content. This occurs after the removal of iodine fluoride species. The chair-type structure then becomes more and more regular and the residual dangling bonds behave as isolated spins and have a regular environment. Since the fluorination occurs due to the presence of intercalated iodine species, the thermal post-treatment of CF(LT) in F2 avoids the formation of conformational defects, contrary to a direct reaction of F2 gas with graphite at 600°C leading to the formation of CF(HT). Because of their low mobility in the fluorocarbon interlayer space, the iodine fluoride species could hinder the diffusion of the F2 molecules, shield the dangling bonds, and strengthen the planar configuration of the carbon sheets. The role of the iodine species is well evidenced by the strong correlation between the evolution of the NMR lines of IFy and the fluorocarbon matrix (Fig. 5): the removal of the iodine species and the change of the C–F bond occurred simultaneously showing their strong interactions. The recent works of Sato et al. [45] show that the progressive fluorination is not possible by a post-treatment in 5.0 MPa (⬃ 50 atm) fluorine gas of fluorine-GIC because the material burned to form fluorocarbon gases in the temperature range 300–400°C. These authors [45] show that the direct conversion of stage 1 fluorine-GIC into CFx occurred only with 0.1 MPa (⬃ 1 atm) F2 gas in the same temperature range also forming graphite fluoride with few defects: the original structure of the raw material with semi-ionic C–F bonds and planar sp2 carbon

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sheets is maintained for post-treatment temperature lower than 400°C, whereas covalent C–F bonds and sp3 carbon atoms are formed at higher temperature. Sato et al. [45] suggest that the post-fluorination is facilitated by the rearrangement of originally intercalated fluorine atoms in the fluorine-GIC. Therefore, the progressive post-fluorination allows the formation of modified graphite fluorides with various physical properties and various electrochemical performances as discussed in the following section. 4. ELECTROCHEMICAL STUDIES OF FLUORINATED GRAPHITE USED AS CATHODIC MATERIALS IN LITHIUM BATTERIES The use of carbon fluorides as cathode materials in nonaqueous primary lithium battery started at the beginning of the 1970s, the electrolyte used being composed of a 1 mol L1 lithium salt (LiX, X  ClO4, PF6 or BF4) dissolved in aprotic solvents (usually employed: propylene carbonate (PC), dimethylsulfoxide (DMSO), γ -butyrolactone (BL), tetramethylene sulfone (TMS), dimethoxyethane (DME), etc.). In the case of fluorinated graphite CFx, the theoretical specific capacity Qth (mAh g1) is given by the following equation: Qth  (xF)/3.6M, where F is the Faraday constant and M the molar mass of CFx. When a discharge current is applied, a departure from the OCV is observed both due to the ohmic drop in the electrolyte and electrode overpotential; it primarily occurs with the charge transfer, ionic diffusion, and/or phase transformation. In practice, the electric energy Epr supplied by the battery when discharged is related to the measured closed-circuit voltage (CCV) and to the electrode utilization (faradic yield ρ F): Epr 

冕

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CCV(z)F dz

0

where z is the cell reaction progress rate. 4.1. High-temperature graphite fluoride CF(HT)

Carbon fluorides CFx (with “x” close to unity) prepared at high temperatures (CF(HT)) have been studied actively as cathode materials in high-energy density lithium batteries [20,46,47]. It is expected that the OCV of the Li/CFx cell depends on the nature of the carbon–fluorine bonding. In fact, a strong covalent C–F bonding yields lower OCV than that measured in weaker C–F bonds found in purely ionic or in semi-ionic CFx materials (Fig. 2). This may relate to a higher CFx reduction overpotential with increase in C–F binding energy. Fig. 11 summarizes some average potentials and specific capacities of CFx prepared by direct fluorination of graphite at temperatures varying between 300 and 600°C, and used as cathodic materials in a lithium cell composed of a liquid electrolyte.

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Fig. 11. Average potential (䊏) and specific capacity () (calculated with cutoff voltage of 1.5 V) as a function of CFx synthesis temperature for lithium cells operated at room temperature with 1 mol L1 LiClO4–PC electrolyte for a current density of 0.5 mA cm2.

4.2. Low-temperature graphite fluoride CF(LT)

CF(LT) exhibits ionic [48] or semi-ionic [49] C–F bonds, in which the sp2 hybridization of the original graphite is maintained. The graphite fluorination level x strongly depends on the chemical nature of the catalyst fluorides. Fig. 12 displays some typical average potential and specific capacities of CF(LT) synthesized at room temperature using various catalysts. For instance, a high fluorination yield (i.e. x  0.8) was achieved in the presence of a IF5 and HF mixture [17,49–51]. When compared with conventional high-temperature CF(HT), these particular CF(LT) yield a similar discharge capacity of about 600 Ah kg1 but differs by having a higher discharge voltage (⬇ 2.9 V) (⬇ 2.0 V for CF(HT) under the same discharge conditions). This difference results in an energy density increase of 30%. Moreover, CF(LT) displays a higher rate capability of up to 5C, while CF(HT) hardly sustains a 2C rate [51]. The C rate corresponds to the recovering of the nominal capacity of the cell in 1 h. 4.3. Re-fluorinated CF(LT): an improvement in the electrochemical performance

When re-fluorination of CF(LT) is performed in the 200–400°C temperature range, the fluorine content x increases to reach a value in the range 0.9–1 [50]. However, re-fluorinated compounds contain minute amounts of IFy catalysts. The latter are highly undesirable as they adversely alter the cell’s shelf life (self-discharge) and increase its internal impedance. In some special applications such as in implantable medical devices, a battery should last about 15 years at human body temperature of 37°C. Therefore the self-discharge rate should be kept very low.

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Fig. 12. Average potential (䊏) and specific capacity () (calculated with cutoff voltage of 1.5 V) as a function of x in CFx for lithium cells operated at room temperature with 1 mol L1 LiClO4–PC electrolyte for a current density of 0.5 mA cm2.

Recently, Guérin et al. have reported an increase in the electrochemical performance of the CF(LT) compounds by a re-treatment under fluorine atmosphere at different temperatures of these post-treated CF(LT) [52]. Li/LiClO4-PC 1 mol L1/CFx cells were studied by galvanostatic measurements of 10 A kg1. Fig. 13 shows the galvanostatic discharges obtained with the fluorination posttreated CF(LT) series. Galvanostatic discharge under various current densities was carried out at room temperature. The discharge capacities detailed here have been obtained for a cell cutoff at 2 V. Fig. 14 shows the fluorination post-treatment temperature dependence of the average discharge voltage E and the discharge capacity that reached at 2 V. At the early stage of discharge (Fig. 13), the voltage drops abruptly for a relatively short time and then re-increases steadily to the main voltage plateau. The initial voltage drop, usually called the “delay effect”, is characteristic of Li/CFx cells and is ascribed to the low electrical conductivity of the active cathode. As the conductivity is improved with the carbon formation (resulting from CFx reduction), the voltage increases due to enhanced discharge reaction kinetics. The relative voltage drop increases with the post-treatment temperature but is still lower than that in CF(HT) even for CF(LT) post-treated under fluorine at the highest temperature. This suggests a lower resistivity of all the post-treated samples when compared with CF(HT). For fluorination post-treatment temperatures between 100 and 250°C, both average discharge voltage and discharge capacity remain constant (around 3.0 V and 570 Ah kg1, respectively). Again, the high discharge voltage results from the semi-ionic character of the C–F bond. The OCV varies with the voltage

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Fig. 14. Average potential (䊏) and specific capacity () (calculated with cutoff voltage of 2 V) as a function of the fluorination post-treatment temperature of CF(LT) for lithium cells operated at room temperature with 1 mol L1 LiClO4–PC electrolyte for a current density of 0.5 mA cm2.

plateau value but may result from a mixed potential of the post-treated CF(LT) and more active fluorine species. The energy density shown in Fig. 15 increases with the capacity and is already, for these fluorination post-treatment temperatures, 70% higher than that of CF(HT), whose energy density value is about

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Fig. 15. Specific capacity () (calculated with cutoff voltage of 2 V) and energy density (䊉) as a function of average discharge potential.

1000 Wh kg1. A discharge capacity of about 570 Ah kg1 is typical for low-temperature graphite fluorides [12]. For a theoretical CF1 composition, assuming all C–F sites are electrochemically active, a capacity close to 850 Ah kg1 can be found. This low discharge yield cannot be explained by inactive CF2 and CF3 groups, typically expected in CF(HT) [53] but by the fluorination level. In fact, CF2 and CF3 groups were not detected by FT–IR and 19F–NMR experiments on all the post-treated fluorinated graphite, except for a fluorination post-treatment at 600°C [27]. Moreover, despite a general composition of CF1 for post-treated CF(LT), some fluorine atoms arise from residual IFy species that contribute to the active mass whereas they are not electrochemically active. For a re-fluorination temperature between 300 and 550°C, the OCV and the average discharge voltage decrease gradually with increase in re-fluorination temperature (Fig. 14). This could be mainly explained by the C–F bond change toward a more covalent character. As shown in Fig. 15, the specific discharge capacity and energy density go through a maximum of 900 Ah kg1 and 2270 Wh kg1, respectively. This maximum falls at around 2.53 V average voltage that corresponds to CF(LT) post-treated at 550°C. The increase in the capacity correlates with the departure of iodine fluoride species and the concomitant increase in the fluorine content with re-fluorination treatment [21]. The energy density is particularly high for the hybrid graphite fluoride compounds (re-fluorination temperature between 400 and 500°C) and the two carbon hybridization states (i.e. sp2 and sp3) seem to favor good electrochemical performance. For a fluorination post-treatment temperature equal to 600°C, the average discharge voltage and the discharge capacity are lower than for CF(LT) post-treated at 550°C (Fig. 14). The increase in the inactive CF2 and CF3 groups

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identified by FT–IR and 19F–NMR measurements [27] could explain the lower capacity as in covalent CF(HT). The fact that despite a similar fluorination temperature (600°C), post-treated CF(LT) yields twice as high as energy density as CF(HT) is noteworthy (i.e. about 2000 vs. 1000 Wh kg1, respectively.) It should be noted that the discharge voltage varies monotonically from 3.15 to 2.3 V with re-fluorination temperature while keeping a high discharge capacity. Indeed, the discharge voltage can be correlated to C–F bond nature, which is a function of the fluorination post-treatment temperature and not of F/C ratio as is generally the case for low-temperature graphite fluorides [46]. The coexistence of sp2- and sp3-hybridized carbons with varying sp2/sp3 ratios that were discussed in the previous section, may explain such voltage evolution in the posttreated CF(LT) cathode materials. As the sp2-type carbons tend to be more electrochemically active and the sp3-type yield higher capacity, the post-treated CF(LT) materials offer a wide range of (voltage, capacity) pairs suitable for specific battery-operating requirements. Tuning of the fluorination post-treatment temperature is the key parameter in achieving such a specific material. 4.4. Reversibility of the electrochemical processes

The nature of the C–F bond in carbon fluorides synthesized at high temperature is covalent and causes a strong lithium fixation on fluorinated sites. Consecutively, during the first discharge, LiF is formed irreversibly [46]. Therefore, covalent carbon fluorides cannot be used in secondary lithium batteries. The physical and chemical studies of compounds, synthesized at room temperature and then re-fluorinated, have demonstrated the semi-ionic nature of the C–F bond for specific post-fluorination temperature and the lithium reversibility has been investigated again. Indeed, Yazami et al. [49] undertook a comparative study of the electrochemical behavior of the carbon fluorides, prepared by direct fluorination at high-temperature CF(HT) and by fluorination at low-temperature CF(LT) using catalysts, in a lithium/solid polymer electrolyte in order to determine the reversibility of the electrode reaction. Cyclic voltammetry with decreasing scan rate (from 1000 to 1 mV min1) was applied at the same temperature (80°C) to two cells, one with CF(HT) and the other with CF(LT). This clearly emphasizes that, whereas in the case of CF(HT), the voltammograms do not exhibit reversible behavior, when CF(LT) is used, associated reduction and oxidation peaks were present, which denote reversible behavior. Proceeding on the reversibility of lithium intercalation into graphite fluorides, all compounds (CF(LT)) post-treated under F2 at different temperatures were studied by cyclic voltammetry in Li/LiClO4-PC 1 mol L1/CFx cells between 2.0 and 4.5 V with a sweeping rate of 6 mV min1. As exemplified in Fig. 16 by CF(LT) post-treated at 200°C, a reduction peak of lithium associated with an oxidation one, centered at 3.0 and 3.2 V, respectively, is present for the fourth cycle as in the first

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ones (Fig. 16a). The same behavior was observed in the case of CF(LT) post-treated at 400°C (Fig. 16b), whereas no reduction peak is noticed for a fluorination posttreatment temperature of 500°C even during the first cycle (Fig. 16c). The current quantities involved in the reversible process are low and they decrease when the

Intensity (mA)

0.10 0.05 0.00 −0.05 −0.10

2.0

2.5

3.0

3.5

4.0

4.5

4.0

4.5

4.0

4.5

E (V) vs. Li+/ Li

(a)

Intensity (mA)

0.045 0.030 0.015 0.000 −0.015 −0.030 2.0

2.5

2.0

2.5

(b)

3.0 3.5 E (V) vs. Li+/Li

0.10

Intensity (mA)

0.05 0.00 −0.05 −0.10 −0.15 −0.20 −0.25 (c)

3.0

3.5

E (V) vs. Li+/Li

Fig. 16. Fourth cyclic voltammograms of [Li/LiClO4–PC 1 mol L1/CF(LT) post-treated] at TFPT  200°C (a), TFPT  400°C (b) and TFPT  500°C (c) (sweeping rate 6 mV min1).

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fluorination post-treatment temperature increases. More generally, a reduction peak associated with an oxidation peak appears for all CF(LT) post-treated at temperatures below 450°C. Such behavior can be correlated with the fact that for such posttreatment temperatures the planarity of the graphene sheets is maintained. For temperatures above 450°C, the planarity of the graphene sheets is partly broken and LiF is inevitably formed in the working electrode consuming all the fluorine atoms upon cycling. In this case, the sample behaves as the conventional CF(HT) [49], exhibiting a voltammogram without a well-defined peak, in particular in the oxidation wave (Fig. 16c). The reduction process occurs mainly at potentials lower than 2.5 V. These results on the electrochemical behavior of CF(LT) confirm the possible reversibility of the intercalation of lithium ions. This opens a new investigation field for fluorinated graphite. 5. CONCLUSION: TO DESIGN MATERIALS WITH AN ADJUSTED CAPACITY – VOLTAGE COUPLE Although lithium batteries using CFx are still the focus of intensive work [54], more attention has been paid during these last years on fluorinated graphite prepared at room temperature. The post-treatment of low-temperature graphite fluoride in fluorine gas in the range 150–680°C result in new derivatives of high-energy density associated with average discharge voltage ranging between 3.15 and 2.35 V and specific capacity between 600 and 900 Ah kg1. The dual nature of the C–F bond of post-treated CF(LT) and its evolution as a function of the fluorination posttreatment temperature, which are evidenced by a combination of characterizations (19F and 13C–NMR, FT–IR, XRD, EPR), make it possible to design materials with a (capacity, voltage) couple adjusted to particular requirements and lead to outstanding electrochemical performance as cathodes in primary lithium batteries. Such materials can also be suitable in secondary lithium batteries. ACKNOWLEGMENTS The authors wish to thank Pr F. Masin (Université Libre de Bruxelles, Belgium), Dr R. Yazami (INPG, LEPMI, St. Martin D’Hères, France), Dr. J. Giraudet, Dr. Z. Fawal (Université Libanaise, Faculté des Sciences III, Tripoli, Liban), Pr. P. Hoggan, and Dr. J.P. Pinheiro for their cooperation and fruitful discussion in this work. REFERENCES [1] D.R. Lide (Ed.), Handbook of Chemistry and Physics, 83rd edn., CRC Press, Boca Raton, FL, 2002. [2] S.S. Chen, A.S. Rodgers, J. Chao, R.C. Wilhoit, and B.J. Zwolinski, J. Phys. Chem. Ref. Data, 4 (1975) 441.

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A. Hamwi et al. D. Peters, J. Chem. Phys., 38 (1963) 561. S.L. Di Vittorio, M.S. Dresselhaus, and G. Dresselhaus, J. Mater. Res., 8 (1993) 1578. N. Watanabe, R. Hagiwara, and T. Nakajima, J. Electrochem. Soc., 131 (1984) 1980. A. Morita, N. Eda, T. Ijima, and H. Ogawa, Power Sources, Vol. 9, J. Thompson (Ed.), Academic Press, New York, 1983, p. 435. A. Hamwi, C. Latouche, V. Marchand, J. Dupuis, and R. Benoit, J. Phys. Chem. Solids, 57 (1996) 991. D. Claves, J. Giraudet, A. Hamwi, and R. Benoit, J. Phys. Chem. B, 105 (2001) 1739. I. Palchan, D. Davidov, and H. Selig, J. Chem. Soc., Chem. Commun., 12 (1983) 657–658. T. Mallouk and N. Bartlett, J. Chem. Soc. Chem. Commun., (1983) 103–105. R. Hagiwara, M. Lerner, and N. Bartlett, J. Chem. Soc. Chem. Commun., (1989) 573. A. Hamwi, J. Phys. Chem. Solids, 57 (1996) 677. A. Hamwi, M. Daoud, J.C. Cousseins, and R. Yazami, J. Power Sources, 27 (1989) 81. R. Yazami and P. Touzain, Solid State Ionics, 9 (1983) 489. M.J. Root, R. Dumas, R. Yazami, and A. Hamwi, J. Electrochem. Soc., 148 (2001) A339. T. Nakajima, In: Fluorine–Carbon and Flouride–Carbon Materials: Chemistry, Physics, and Applications, T. Nakajima (Ed.); Marcel Dekker, New York, 1995, p 11. A. Hamwi, M.Daoud, and J.C. Cousseins, Synth. Met., 26 (1988) 89. A.M. Panich, Synth. Metals, 100 (1999) 169. H. Touhara and F. Okino, Carbon, 38 (2000) 241. T. Nakajima and N. Watanabe, Graphites Fluorides and Carbon–fluorine Compounds, T. Nakajima (Ed.), CRC Press, Boca Raton, FL, 1991, p. 84. K. Guérin, J.P. Pinheiro, M. Dubois, Z. Fawal, F. Masin, R. Yazami, and A. Hamwi, Chem. Mat., 16 (2004) 1786. A.M. Panich, T. Nakajima, H.M. Vieth, A.F. Privalov, and S.D. Goren, J. Phys.: Condens. Matter, 10 (1998) 7633. A.M. Panich, T. Nakajima, and S.D. Goren, Chem. Phys. Lett., 271 (1997) 381. A.M. Panich, A.I. Shames, and T. Nakajima, J. Phys. Chem. Solids, 62 (2001) 959. T.R. Krawietz and J.F. Haw, Chem. Commun., 19 (1998) 2151. A. Hamwi, M. Daoud, D. Djurado, J.C. Cousseins, Z. Fawal, and J. Dupuis, Synth. Metals, 44 (1991) 75. M. Dubois, K. Guérin, J.P. Pinheiro, Z. Fawal, F. Masin, and A. Hamwi, Carbon, 42 (2004) 1931–1940. H. Selig, W.A. Sunder, M.J. Vasile, F.A. Stevie, P.K. Gallagher, and L.B. Ebert, J. Fluorine Chem., 12 (1978) 397. R.K. Heenan and R. Robiette, J. Mol. Struct., 55 (1979) 191. K.O. Christe and W.W. Wilson, Inorg. Chem., 28 (1989) 3275. N. Bartlett, S. Beaten, L.W. Reeves, and E.J. Wells, Can. J. Chem., 42 (1964) 2531. C.A. Wilkie, G.-Y. Lin, and D.T. Haworth, J. Solid State Chem., 30 (1979) 197. A.M. Panich and T. Nakajima, Mol. Cryst. Liq. Cryst., 340 (2000) 77. J. Giraudet, M. Dubois, K. Guérin, J.P. Pinheiro, A. Hamwi, W.E.E. Stone, P. Pirotte, and F. Masin, accepted in J. Solid State Chem. H.A Resing, J. Milliken, D.D. Domingues, and L.E. Iton, Proc. 17th Biennal Conf., Carbon, Lexington, Kentucky University, 1985. T. Mallouk, B.L. Hawkins, M.P. Conrad, K. Zilm, G.E. Maciel, and N. Bartlett, Phil. Trans. R. Soc. Lond. A, 314 (1985) 179. E.W. Hagaman, D.K. Murray, and G.D.Del Cul, Energy & Fuel, 12 (1998) 399. S. Mouras, A. Hamwi, D. Djurado, and J.C. Cousseins, Rev. Chimie Minérale, 24 (1988) 572. H. Yokomichi, T. Hayashi, T. Amano, and A. Masuda, J. Non-Cryst. Solids, 227 (1998) 641.

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[40] H. Yokomichi and K. Morigaki, J. Non-Cryst. Solids, 266 (2000) 797. [41] K. Takai, H. Sato, T. Enoki, N. Yoshida, F. Okino, H. Touhara, and M. Endo, Mol. Cryst. Liq. Cryst., 340 (2000) 289. [42] S.L. Di Vittorio, T. Enoki, M.S. Dresselhaus, G. Dresselhaus, M. Endo, and T. Nakajima, Phys. Rev. B: Cond. Mater. Mater. Phys., 46 (1992) 12723. [43] R. Davidov, O. Milo, I. Palchan, and H. Selig, Synth. Metals, 8 (1983) 83. [44] M. Murata and H. Suematsu, J. Phys. Soc. Jpn., 51 (1982) 1337. [45] Y. Sato, S. Shiraishi, Z. Mazej, R. Hagiwara, and Y. Ito, Carbon, 41 (2003) 1971. [46] R. Yazami, Chemistry, Physics and Applications of Fluorine–Graphite and Fluoride–Carbon Compounds, T. Nakajima (Ed.), Marcel Dekker, New York, 1995, pp. 251–281. [47] N. Watanabe, Solid State Ionics, 1 (1980) 87. [48] H. Touhara, Y. Goto, N. Watanabe, K. Imaeda, T. Enoki, H. Inokuchi, and Y. Mizutani, Synth. Metals, 23 (1988) 461. [49] R. Yazami and A. Hamwi, Solid State Ionics, 28 (1988)1756. [50] A. Hamwi and R. Yazami, Flourine-containing carbonaceous substances, their preparation, and use as battery electrodes, WO Patent, 97/41061 (1997). [51] A. Hamwi, M. Daoud, and J.C. Cousseins, Synth. Metals, 30 (1989) 23. [52] K. Guérin, R. Yazami, and A. Hamwi, Electrochem. Solid State Lett., 7(6) (2004) A159. [53] Y. Kita, N. Watanabe, and Y. Fujii, J. Am. Chem. Soc., 101 (1979) 3832. [54] A.G. Ritchie, C.O. Giwa, P.G. Bowles, J. Burgess, E. Eweka, and A. Gilmour, J. Power Sources, 96 (2001) 180.

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 18

Battery application of graphite intercalation compounds Yoshiaki Matsuo Department of Materials Science and Chemistry, Graduate School of Engineering, University of Hyogo, 2167, Shosha, Himeji Hyogo, 671-2201, Japan 1. INTRODUCTION Graphite intercalation compounds are very useful for battery applications because of the high electrical conductivity and their two-dimensional layered structure, which enables further intercalation of foreign ions such as lithium ions. Therefore, they have been used for the electrode materials of batteries such as lithium primary battery [1–6], lithium-ion battery [7,8], alkaline battery [9,10] and thermo cells [11–14]. Graphite intercalation compounds are classified into three categories from the viewpoint of chemical bonding between graphite and intercalated species as shown in Fig. 1. In donor and acceptor type intercalation compounds, graphite is positively and negatively charged, respectively, and it is bound to intercalated species via ionic bonding. The intercalated species include alkaline metal, alkaline earth metal, transition metal, etc. to form donor-type intercalation compounds. Among them, it is well known that lithium-intercalated graphite has been used as an anode of lithium-ion battery. Halogens, acids, oxides, etc. form acceptor-type intercalation compounds and some of them have been tested as cathode-active materials of lithium primary battery and electrode material of alkaline cell. On the other hand, fluorine and oxygen with high electronegativities are known to form covalent-or semicovalent-type graphite intercalation compounds when they are allowed to react with graphite under appropriate conditions. They are poly(carbon monofluoride), (CF)n, poly(dicarbon monofluoride), (C2F)n, and graphite oxide (abbreviated as GO or it is sometimes called graphitic acid). In these compounds, the planarity of carbon sheet of graphite is completely lost and the carbon layer is thought to consist of cyclohexane-like zig-zag carbon array. These materials are insulators; however, in the presence of conducting additives they are used as cathode-active materials of lithium primary battery [1,5].

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Yoshiaki Matsuo

Acceptor type Acids Metal oxide Metal halides etc

Thermo cell

Donor type Alkaline metals Alkaline earth metals etc

Graphite (Residual carbon)

Δ Surface modification

Δ

Alkaline cell

Semi-covalent type fluorine (CxF, x<3)

Covalent type Graphite fluorides ((CF)n, (C2F)n) Graphite oxide

Anode active material of lithium ion battery Cathode active material of lithium primary battery

Fig. 1. Classification and battery application of graphite intercalation compounds.

The battery performance of graphite intercalation compounds are greatly affected by the structural parameters of the host graphite. These include the crystallite sizes along the a-axis (La) and the c-axis (Lc), surface area, pore structure, interlayer spacing (d(0 0 2)), etc. The important structural parameters described in this chapter are shown in Fig. 2. Note that the lattice parameters and crystallite sizes of natural graphite are a0  0.246 nm, c0  0.671 nm, Lc(0 0 2)  100 nm and La(110)  100 nm. One of the ways to modify these parameters of carbons is the thermal decomposition of covalent-type of graphite intercalation compounds with strong bonding between carbon and intercalated species. The crystallite sizes, interlayer spacings, etc. of the obtained residual carbons are modified during decomposition, probably due to the formation of defects as a result of escape of carbon with intercalated species as gaseous products such as fluorinated hydrocarbons, carbon dioxide or carbon monoxide, exfoliation of carbon sheet and dissociation of carbon–carbon bondings caused by vigorous and rapid exothermic reaction. These could lead to the improvement of battery performance when the resulting carbons are used as an active material or a precursor of active material of batteries, because the diffusion of ionic species, for example Li ion, in the active materials will be highly improved. In this chapter, recent advances on the preparation and electrochemical properties of covalent-type graphite intercalation compounds and residual carbons prepared via their thermal decomposition are summarized.

Battery application of graphite intercalation compounds

c

399

c

d(110) d(002) c0

a

c0

a0

b

a

b

La

Lc

Fig. 2. Lattice parameters, diffraction planes and crystallite sizes of graphite.

2. COVALENT-TYPE GRAPHITE INTERCALATION COMPOUNDS The structure and properties of graphite fluorides are well known as summarized in the literature [15]. The carbon layers of the graphite fluorides consist of an infinite array of trans-linked cyclohexane chairs without aromatic nature. (CF)n is regarded as a first-stage compound, whereas (C2F)n is a second-stage compound. They are used as cathode materials of lithium primary battery and lubricant. On the other hand, the structure of GO is still unknown, though it was first synthesized in the 1850s [16]. GO is prepared by oxidizing highly graphitized carbons in concentrated acids using strong oxidizing reagents such as potassium chlorate, manganese peroxide or electrochemical method [17–19] and then by treating the intermediate oxidized graphite with a large excess of water. The composition of GO widely varies, depending on the synthetic methods and crystallinity of the starting carbons. The O/C and H/C ratios are normally in the range of 3–4 and 2–3. Fig. 3 shows the typical X-ray diffraction pattern of GO with the composition of C8O3.3H2.4, which was prepared based on the Brodie’s method from natural graphite powder in fuming nitric acid using potassium chlorate as

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Yoshiaki Matsuo

Intensity (A.U.)

Graphite oxide

10

Graphite

15

20 2/deg. CuK

25

30

Fig. 3. X-ray diffraction patterns of graphite oxide, together with that of graphite.

an oxidizing reagent, along with that of pristine graphite. The diffraction peak shifted from 2θ  26.5 to 15°, which shows the increase in interlayer spacing from 0.3354 to 0.63 nm. It has been known that the interlayer spacing of GO changes in the range between 0.6 and 1.1 nm, depending on the ambient humidity [6]. The chemical and spectroscopic analyses showed that GO has various oxygen containing functional groups such as phenolic hydroxyl, carboxyl and ether groups, which are well summarized by Boehm [20]. Two types of structure models have been proposed, one is stage 1 type [20–29] and the other is stage 2 type [30,31]. Many researchers have proposed stage 1 type structure model like (CF)n; however, Nakajima et al. [30] proposed a stage 2 type structure model based on the finding that the Ic values of fluorinated GO at relatively low temperature were rather similar to that of (C2F)n but not (CF)n. It has been suggested that the structure of GO is the intermediate of the compounds with ideal compositions of C8O2 and C8(OH)4. GO was used as an alternative of graphite fluorides for the cathode of lithium primary battery; however, the lower thermal stability of GO was unfavorable [32]. Improvement of thermal stability and battery performance have been achieved by fluorinating GO at relatively low temperatures [33,34]. Recently, it has been reported that various chemical species are intercalated into the layer of GO including various polymers, cationic surfactants and n-alkylamines [35–58], and used as a cathode active material of rechargeable lithium battery [41,49], an adsorbent for non-ionic organic molecules [37,50], photochemical reaction media [51–54] and matrices of photofunctional molecules [55–58]. Thermal decomposition of GO has been studied from the viewpoint of graphitization process by several researches [59–65]. However, the properties of GOs pyrolyzed at relatively low temperatures have not been characterized sufficiently and the studies of the application of these carbons was quite limited.

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401

3. STRUCTURE AND REACTIVITY OF CARBONS PREPARED FROM COVALENT-TYPE GRAPHITE INTERCALATION COMPOUNDS 3.1. Residual carbons from covalent-type graphite intercalation compounds as precursors of (CF)n [66,67]

Various carbons prepared from the thermal decomposition of (C2F)n and the intermediate oxidized graphite were used as precursors of (CF)n. The thermal decomposition of (C2F)n started at 480°C and completed at 600°C in an argon atmosphere, when the temperature was elevated by 5°C/min. The residual carbon contained 4% of fluorine. Another residual carbon was obtained by decomposing the intermediate product obtained during the preparation of GO as follows. To a mixture of natural graphite, sodium nitrate and sulfuric acid, potassium permanganate was added slowly and stirred. Then, water was gradually poured into the solution for about 10 min, during which the temperature raised up to 180°C by the heat of hydration. It was supposed that graphite oxide formed at first and decomposition occurred successively. SEM observation of the above carbons indicated the exfoliation along the c-axis. Fig. 4 shows the relationship between crystallite sizes of residual carbon and the amount of added KMnO4 as an oxidizing reagent. Both the crystallite sizes La and Lc decreased with the increase in the amount of KMnO4. The decrease in crystallite size with increase in the amount of KMnO4 was larger in La than in Lc; however, La was always larger than Lc. The BET surface area was 20–30 times than that of original graphite. The increase in reaction time leads to the decrease in crystallite size. When these carbons were fluorinated, the reciprocal of the half-width of (002) line ((002)1) of the resulting (CF)n was larger than that obtained from natural graphite. The (002)1 values are proportional to the crystallite size along the c-axis. They were very small compared with the values of (CF)n

Crystallite size/nm

50

La

40 30 Lc 20 10 0 0

10 20 30 Amount of KMnO4 / g

40

Fig. 4. Relation between crystallite sizes of residual carbon and amount of added KMnO4.

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Yoshiaki Matsuo

prepared by direct fluorination ((002)1  0.3–0.78). Table 1 shows the summary of experimental results for carbons prepared via graphite oxide. 3.2. Carbons prepared via pyrolysis of graphite oxide under reduced pressures or hydrogen gas flow [62–65]

Fig. 5 shows the X-ray diffraction patterns of GO prepared by the Brodie’s method and that pyrolyzed at 300°C under vacuum for 10 h (hereafter PGO300V). In order to avoid vigorous exothermic reaction and exfoliation, the rate of increase in temperature was very low, that is 1°C/min. The composition of PGO300V determined from elemental analysis of C, H and O was C20O. The diffraction peak appeared at 2θ 15° for graphite oxide shifted to the higher angle, at 21° after pyrolysis, indicating that the d spacing of the obtained carbon was surprisingly large, that is, 0.404 nm. The diffraction peak of PGO300V was relatively sharp, compared with carbons obtained from heat treatment of polymer precursors. This suggests Table 1 Experimental results for (CF)n prepared via graphite oxide Synthetic conditon of graphic oxide

Residual carbon

Graphite fluoride

KmnO4 (g)

Reaction time (min)

Crystallite size, Lc(nm)

Fluorination temperature (C)

F/C

1A

5

20

200

400

0.89

2A

20

20

170

400

0.90

3A

30

20

130

400

1.03

4A

40

20

85

400

1.00

3B

30

60

120

400

0.88

3B

30

60

120

350

0.90

3B

30

60

120

338

0.79

Intensity (A.U.)

Sample

5

10

15

20 25 2/deg. CuK

Fig. 5. X-ray diffraction pattern of PGO300V (C20O).

30

35

40

Battery application of graphite intercalation compounds

403

high order of the orientation of carbon layers. This type of carbon with a large interlayer spacing has been reported by Maruyama [59] 50 years ago; however, the structure and properties have not been characterized sufficiently. Fig. 6 shows the TEM image of PGO300V. Many defects were observed in the carbon layers; however, the orientation of carbon layers was relatively high, as indicated by the X-ray diffraction data. This would be because graphite oxide was prepared from natural graphite with high layer regularity and this was considerably maintained even after decomposition. Fig. 7 shows the Raman spectra of PGO300V. Two broad peaks at 1350 and 1580 cm1 due to D and G bands were observed. Reflecting the fact that a large number of defects exist, the crystallinity of the surface of PGO300V is low. As the temperature for pyrolysis increased, the interlayer spacing decreased gradually and became almost similar to that of graphite. Raman spectrum of PGO900V was still very similar to that of PGO300V. Removal of oxygen atoms bonded to carbon atoms leads to the decrease in interlayer spacing. However, sufficient energy was not

0.4nm

Intensity (A.U.)

Fig. 6. TEM image of PGO300V (C20O).

1250

1350

1450

1550

Raman shift /

Fig. 7. Raman spectrum of PGO300V (C20O).

cm-1

1650

1750

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Yoshiaki Matsuo

Intensity (A.U.)

provided for reorientation of carbon atoms in order to reduce defects within carbon layers at these heat-treatment temperatures. When GO was pyrolyzed under H2 gas flow, similar product with large d spacings was obtained. However, the Raman spectra of the carbon atoms obtained at high temperatures as shown in Fig. 8 were considerably different from that obtained under vacuum. Three peaks at 1350 (D band), 1580 (G band) and 1620 cm1 (D band) were observed for all the samples and as the temperature increased, the relative intensity of D band was more pronounced. This indicates that the crystallinity of the carbon from pyrolysis of GO under H2 gas flow was higher than that obtained under vaccum. This could be because unstable carbon atoms reacted with hydrogen and were removed from the sample. In addition, as shown in Table 2, the carbon atoms obtained under H2 gas flow contained a small amount of hydrogen that might terminate the edge of the active carbons.

1250

(C) (B)

1350

1450 1550 Raman shift / cm-1

1650

(A) 1750

Fig. 8. Raman spectrum of PGO900H, PGO950H and PGO1000H.

Table 2 Composition, H/C ratio and d(002) values of carbon obtained via pyrolysis of graphite oxide under hydrogen gas flow at various temperatures Temperature (ºC)

Composition

H/C

d(0 0 2) (nm)

300

CO0.035H0.023

0.023

0.404

500

CO0.016H0.050

0.050

0.348

700

CO0.0028H0.063

0.063

0.339

900

CO0.0007H0.084

0.084

0.337

925

CH0.059

0.059

0.337

950

CH0.056

0.056

0.336

1000

CH0.038

0.038

0.334

Battery application of graphite intercalation compounds

405

The layer regularity of these carbon atoms obtained from pyrolysis of GO is very high. However, many defects still exist within the layer as observed for non-graphitized carbons. 4. ELECTROCHEMICAL PROPERTIES OF GRAPHITE INTERCALATION COMPOUNDS 4.1. Cathode properties of residual carbon atoms prepared from covalent graphite intercalation compounds [66,67]

Fig. 9 shows the discharge curves of (CF)n prepared from the residual carbon atoms via (C2F)n at 450°C and GO at 400°C, in comparison with (CF)n and (C2F)n obtained by direct fluorination of natural graphite at a constant current density of 0.5 mA cm2. Higher discharge potentials than that of conventional graphite fluorides were observed for (CF)n prepared from both residual carbons, while the flat discharge potential was maintained. The discharge potential of (CF)n prepared from residual carbon via GO was 0.2 and 0.5 V, higher than that of (CF)n and (C2F)n obtained by direct fluorination of natural graphite, respectively. A large number of lattice defects would be responsible for the lower overpotential. Fig. 10 shows the relationship between cathode overpotential of (CF)n electrodes as a function of (002)1. The overpotential rapidly decreases with decrease in (002) 1, especially in the range of less than 0.4 deg1. In particular, the change in overpotentials among the samples prepared under the same conditions except the oxidation time by KMnO4 is almost independent of (0 0 2)1. This indicated that there is another factor that controls the overpotential in addition to the crystallite size Lc. Fig. 11 shows the change of overpotential as a function of the narrow peak in 19F-NMR spectra that corresponds to the concentration of the lattice defects, that is, the fluorine species weakly bonded to polynuclear

Potential / V vs.Li/Li+

4 3

(C)

2

(D)

(A) 1 0

(B) 0

200

400 Capacity/mAh/g

600

800

Fig. 9. Discharge curves of graphite fluoride cathodes (0.5 mA cm2, 1 M LiClO4–PC, 25°C): (A) (CF)n prepared from Madagascar natural graphite at 600°C; (B) (C2F)n prepared from Madagascar natural graphite at 350°C; (C) (CF)n prepared from residual carbon via (C2F)n at 450°C; (D) (CF)n prepared from residual carbon via graphite oxide.

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Yoshiaki Matsuo

-0.65

Overpotential/V

-0.75 -0.85 -0.95 -1.05 -1.15 -1.25

0

0.2

0.4

0.6

0.8

 (002)-1/ deg-1

Fig. 10. Cathode overpotentials of (CF)n electrodes as a function of (0 0 2)1 (0.5 mA cm2, at 25% discharge). (CF)n prepared via graphite oxide; (CF)n prepared by direct fluorination. -0.55

Overpotential / V

-0.65

-0.75

-0.85

-0.95

-1.05 0.04

0.045

0.05 0.055 I(narrow) / I(wide)

0.06

Fig. 11. Change of cathode overpotential as a function of the ratio of narrow-to-wide peak intensities in 19F-NMR spectra for graphite fluoride.

aromatic carbon rings. As the relative intensity of the narrow peak increased, the overpotential almost decreased linearly. It was reported that during the discharge of graphite fluoride, solvated lithium ions are intercalated into graphite fluoride, making a thin layer of a ternary graphite intercalation compound that decomposes by the subsequent disproportionation reaction [68]. When there are many defects in graphite fluorides, the decomposition of the intermediate discharge product is facilitated and the diffusion of Li ion in the diffusion layer leads to the decrease in overpotential. The low-temperature fluorination of residual carbon also brings about the decrease in overpotential as a result of the increase in

Cathode potential / V

Battery application of graphite intercalation compounds

407

3.5 3 (C)

2.5 2 1.5 0.01

(A) 0.1

1 10 100 Current density/Acm-2

1000

(B) 10000

Fig. 12. Galvanostatic polarization curves of graphite fluorides at 25% discharge: (A) (CF)n prepared from Madagascar natural graphite at 600°C; (B) (C2F)n prepared from Madagascar natural graphite at 350°C; (C) (CF)n prepared from residual carbon via graphite oxide at 400°C.

the number of lattice defects. Fig. 12 shows the galvanstatic polarization curves of graphite fluorides prepared from carbon via graphite oxide (3B) in comparison with the curves of (CF)n and (C2F)n obtained by direct fluorination of natural graphite; these were obtained between 10 and 30% discharge. The most noticeable fact is the remarkable increase in power density, that is, the power density of the former sample was around 25 times than that of conventional graphite fluorides working at a potential of 3.0 V. The contribution of specific surface area was negligible. 4.2. Anode properties of carbon atoms obtained via pyrolysis of graphite oxide in lithium-ion battery [62–65]

Various kinds of carbonaceous materials have been tested for the anode of lithium-ion battery as an alternative for commercially used graphite with a theoretical capacity of 372 mAh g1. These are more or less disordered and some have imperfections within them such as stacking disorder of carbon layers the socalled turbostratic disorder and unorganized or buckled layers. Two types of carbonaceous materials are known, one is “soft carbon” that is graphitizable when heat-treated at high temperatures and the other is “hard carbon” that is not graphitizable even at high temperatures. It has been known that both La and Lc values increase as the heat-treatment temperature of soft carbon increases, while d(002) value decreases to 0.3354 nm of graphite [69,70]. The capacity of carbonaceous materials used as anodes of lithium-ion battery as a function of heattreatment temperature is shown in Fig. 13. The capacity of the carbon treated above 2400°C is in the range of 300–370 mAh g1, which is similar to that of graphite. When the heat-treatment temperature is between 1800 and 2400°C, the capacity is relatively low. The capacity greatly increases with the decrease in heat-treatment temperature and exceeds that of graphite. Fig. 14 shows the charge – discharge curves of carbon obtained from pyrolysis of graphite oxide at 300°C under H2 gas flow (hereafter PGO300H), at a

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1000 Some hard carbons

Capacity/mAh/g

800

Most soft carbons

600 400 200 0 500

1000

1500 2000 2500 Heat treatment temperature /°C

3000

Fig. 13. Relationship between reversible capacity and heat-treatment temperature of various carbonaceous materials: 䊉, soft carbon atoms; O, hard carbon atoms.

Potential / V vs.Li/Li+

3.5 3 2.5 2 1.5 1 0.5 0

0

500

1000 Capacity/mAh/g

1500

2000

Fig. 14. Charge – discharge curves of PGO300 at a constant current density of 20 mA g1.

constant current density of 20 mA g1. The discharge potential almost monotonically decreased during the first discharge, which is typical for non-graphitized “soft carbons.” The discharge capacity reached 1700 mAh g1, however, the charge capacity was low, that is, 400 mAh g1. This indicates a large irreversible capacity due to solvent decomposition and unextracted lithium ions. The color of the sample changed from black to blue at 0.6 V and then to gold at around 0 V. This suggests that the charge transfer from lithium ion to carbon occurred as is observed for graphite used as an anode of lithium-ion battery and intercalation of lithium ion occurs as low as 0.6 V. At a capacity below 0.6 V (1200 mAh g1), a large amount of lithium ions would be intercalated into PGO300H, as expected from Fig. 13. The interesting feature of this carbon was the structural change during charge – discharge measurement. Fig. 15 shows the X-ray diffraction patterns of the sample at various potentials. The measurement started within 5 min after the sample was

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Intensity (A.U.)

(G) (F) (E) (D) (C) (B) (A) 10

15

20 25 2/deg. CuK 

30

35

Fig. 15. X-ray diffraction patterns of PGO300H before (A) and after discharged to various potentials: (B) 0.834 V, (C) 0.600 V, (D) 0.439 V, (E) 0.290 V, (F) 0.057 V and (G) 0.007 V.

removed from the cell without washing the residual electrolyte, in order to avoid the decomposition of the products when exposed to ambient atmosphere. The diffraction peak at 2θ  22° observed for the pristine sample shifted to 2θ 19.6° when the potential reached 0.60 V. The increase in the interlayer spacing was about 0.05 nm, which was comparable with that observed for stage 1 LiC6 (ca. 0.04 nm). Further increase in the interlayer spacing was observed with decrease in potential. At 0.007 V, the interlayer spacing became 0.522 nm; this was 0.118 nm larger than that before discharge. This was almost twice of that observed at 0.6 V. It has been suggested that intercalation of lithium ion solvated by organic molecules occurs at a rather higher potential region than that of lithium ion. In addition, the increase in the interlayer spacing should be more than 0.39 nm, considering the size of lithium ion solvated by ether or carbonate. This value is much larger than that observed for the present sample and excludes the intercalation of lithium ion solvated by propylene carbonate into PGO300H. Therefore, it would be reasonable to realize that lithium ions lie as double layer between carbon layers of PGO300H. Fig. 16 shows the charge – discharge curves of PGO300H with various charge cutoff potentials. The charge capacity was very small suggesting that deintercalation of lithium was rather difficult from PGO300H samples charged to 0–0.6 V. The existence of sp3 carbons in the layer of the material prepared from the pyrolysis of GO might cause the formation of Li–C bonding with considerable covalent nature, which was strong enough against electrochemical deintercalation of lithium. However, the carbon lithiated at 0.057 V was not stable under an ambient atmosphere and decomposed giving amorphous carbon and Li2CO3. The large increase in interlayer spacing, about 0.12 nm, was commonly observed for carbon atoms obtained by the pyrolysis of GO under H2 gas flow below 700°C, when they were charged to 0 V. As the interlayer spacing of the

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Yoshiaki Matsuo charge (A)

(B)

discharge

(C) (D)

Potential / V vs.Li/Li

+

4 3 2 1 0 2000

1500

1000 500 Capacity / mAh/g

500

Fig. 16. Charge – discharge curves of PGO300H obtained with various cut-off potentials (1 M LiClO4–PC, 20 mA/g, 25°C).

carbon atoms decreased, the diffraction peak of lithiated carbon obtained at 0 V shifted to higher angle. On the other hand, the increase in interlayer spacing of PGO900H charged to 0 V was only 0.04 nm, which was similar to that observed for lithiated graphite, that is, LiC6. The interlayer spacings of these carbon atoms having large interlayer expansion by lithium intercalation were larger than 0.339 nm. The interaction between adjacent carbon layers would be weaker than that of graphite with 0.3354 nm of interlayer spacing. This could lead to the intercalated lithium interacting with only one carbon layer in PGO samples, forming double-layer structure, while lithium ions are located at the center of benzenelike hexagonal ring of two adjacent carbon layers in graphite. 4.3. Anode properties of carbon atoms obtained via pyrolysis of graphite oxide in propylene carbonate-based electrolyte solution

Other interesting results on PGO samples, prepared under H2 gas flow, were the anode properties in 1 M LiClO4–propylene carbonate solution, when it was prepared at relatively high temperatures under H2 gas flow. Propylene carbonate (PC) with a high salt solubility and low melting point is favorable for lithium-ion battery. However, when it is used for electrolyte solution with graphite as an anode, exfoliation of graphite occurs before lithium insertion. This phenomenon is explained as follows: intercalation of large lithium ions solvated by PC molecules occurs, leading to the decomposition of graphite layer along with the propylene evolution as a result of reduction of PC at around 0.9 V vs. Li/Li [70]. Ethylene carbonate that is usually used as an electrolyte of lithium battery is also inserted into carbon layer. However, the decomposition products easily form surface-protective film on the surface of graphite (solid electrolyte interface, SEI). The SEI film prevents further decomposition of graphite and EC. In order to

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avoid the decomposition of PC, surface treatment of graphite and addition of organic or inorganic additives to PC have been conducted. The former approach includes coating of carbon with amorphous nature [71]. Many additives with higher reduction potentials than that of PC to form SEI layer and larger salvation energies for lithium ion to reduce the solvated PC molecules have been reported, such as CO2, N2O, SO2, Sx2, ethylene sulfite, vinylene carbonate, chloroethylene carbonate, catechol carbonate, 12-crown-4, butyrolactone derivatives, etc. [72–86] Fig. 17 shows the charge – discharge curves of PGOH900, PGOH950 and PGOH1000 in 1 M LiClO4–propylene carbonate solution at a constant current of 20 mA g1. A short plateau indicating the reduction of PC was observed at 0.9 V; however, the potential reached 0 V after the long plateau at around 0.2 V without any evolution of propylene. It is known that co-intercalation of solvated Li ion and exfoliation of carbon layer are not observed for non-graphitized carbons. Although the interlayer spacing of PGO was very close to that of graphite as mentioned in Section 3.2, the surface structure such as surface area was not similar to that of graphite but that of carbon with lower crystallinity. This would avoid the co-intercalation of solvated lithium ion and subsequent exfoliation of carbon layer as was observed for carbon-coated graphite.

Potential / V vs.Li/Li+

1.5

1 (B) 0.5 (C) 0

(A)

0

500

1000 Capacity / mAh/g

1500

Potential / V vs.Li/Li+

3 2.5

(C)

(B)

( A)

2 1.5 1 0.5 0

0

100

200 300 400 Capacity / mAh/g

500

600

Fig. 17. Charge – discharge curves of (A)PGO900H, (B)PGO950H and (C)PGO1000H. (1 M LiClO4–PC, 20 mA/g, 25°C).

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4.4. Electrochemical hydrogen storage in carbons obtained via pyrolysis of graphite oxide [87]

Hydrogen storage is one of the important issues to develop an environmentally friendly energy system using hydrogen gas. Several reports have suggested that a large amount of hydrogen gas was physically adsorbed in carbon materials with controlled nanostructure such as carbon nanotubes and carbon nanofiber [88–90]; however, the results were not reproducible and contained significant experimental error. Electrochemical hydrogen storage into carbon materials has also been reported [91–97]. Fig. 18 shows the cyclic voltammogram of multiwalled carbon nanotube (MWCNT) containing nickel powder in 6 M KOH solution at 10–50 mVs1 [92]. A couple of redox peaks was observed at 0.9 and 0.6 V in both cases and the current peak was enhanced in the presence of MWCNT. The result suggests that hydrogen atoms formed as the result of

300

Current / mA

200 100 0 -100 -1.2

-0.7

-0.2

-200 -300 -400 Potential / V vs. Hg/HgO

(a) 6 4

Current / mA

2 0 -1.2 -2

-0.7

-0.2

-4 -6 -8 (b)

Potential / V vs. Hg/HgO

Fig. 18. Cyclic voltammograms of MWCNT electrodes at different sweep rates (1, 10 mV s1; 2, 20 mV s1; 3, 50 mV s1) (a) nickel electrode and (b) MWCNTs-Ni electrode.

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electrochemical reduction at nickel surface, moved to MWCNT and adsorbed according to the following mechanism: Ni  H2O  e

y

NiHad  MWCNT y

NiHad  OH

(1)

MWCNT-Had  Ni

(2)

Nickel as a conducting additive is believed to play an important role for storing hydrogen in the above carbon atoms. Charge – discharge measurement indicated that 0.8 wt% of hydrogen was stored at the first cycle. Purified single-walled carbon nanotubes (SWCNT) showed the highest electrochemical hydrogen capacity of 2.9 wt%. However, in these reports, it has not been clear where hydrogen is stored. Theoretical calculation has been reported by several groups and hydrogen molecules are stored in the tube or between tubes. In both sites, hydrogen is stored in the empty space surrounded by carbon sheets. Carbon film obtained via pyrolysis of graphite oxide with a large space in the interlayer spacing and relatively high regularity of the orientation of carbon layers has given some important information for this issue. Fig. 19 shows the cyclic voltammogram of PGO300V thin film on Ni substrate by spin coating of colloidal solution of GO. Ni substrate was covered with PGO300V film and epoxy resin; therefore, it was not exposed to electrolyte solution. Accordingly, reaction (1) would not occur in this system. A small anodic peak at 0.98 V vs. Hg/HgO, which is indicated by an arrow in Fig. 21, was observed before the vigorous evolution of hydrogen gas appeared above 1.1 V vs. Hg/HgO. This peak position was similar to that observed for hydrogen storage in MWCNT shown above [92]. Fig. 20 shows the X-ray diffraction patterns of PGO300V, potentiostatically reduced at 1 V vs. Hg/HgO for 0.25–48 h and then dried at 60°C overnight after washing with water, together with that of the

0

Current/mA

-500 -1000

-1500

-2000 -1.2

-1

-0.8

-0.6

-0.4

-0.2

0

Potential / Vvs.Hg/HgO

Fig. 19. Cyclic voltammogram of PGO300V film electrode in KOH solution at 5 mV s1.

Yoshiaki Matsuo

Intensity (A.U.)

414

*

*

*

*

10

15

20 2 / deg. Cu K

25

pristine A 0.25h

B

1h

C

24h

D

48h

E

30

Fig. 20. X-ray diffraction patterns of (A) PGO300V film and those potentiostatically reduced at 1 V vs. Hg/HgO for (B) 0.25 h, (C) 1 h, (D) 24 h and (E) 48 h. The asterisks indicate the potassium bicarbonate formed from residual KOH. -1.2 Potential / V vs. Hg/HgO

5th -1 1st -0.8 discharge

-0.6 -0.4 -0.2 0

charge 0

200

400 Capacity / mAh/g

600

Fig. 21. Charge – discharge curves of PGO300 in 6 M KOH solution during 1st and 5th cycles.

pristine sample. With increase in the duration of electrolysis, the intensity of the diffraction peak at 2θ  21.5° decreased and a new peak appeared at 2θ  20.4°. The increase in the interlayer spacing was 0.03 nm, which was much smaller than that observed for lithium-intercalated PGO (up to 0.12 nm) as shown in Section 4.2. This peak became very broad when PGO300V was reduced for 48 h. Characteristic peaks due to C–H stretching vibration appeared at around 2900 cm1 only for the reduced sample in IR spectrum of the film reduced for 48 h. This clearly indicates that hydrogen was introduced into PGO300V by electrochemical reduction and considering the increase in interlayer spacing, some amount of hydrogen could exist between carbon layers of PGO300V. Fig. 21

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shows the charge – discharge curves of PGO300. Although a long plateau was observed at around 1.1 V vs. Hg/HgO during discharge, the potential increased monotonically. The charge capacity was 52 mAh g1 (0.2 wt% of hydrogen), indicating that hydrogen atoms introduced into PGO300V can be electrochemically extracted. The large irreversible and low-reversible capacity could be due to hydrogen evolution at PGO300V electrode at around 1 V vs. Hg/HgO and formation of strong C–H bonding, which is hardly oxidized by electrochemical treatment. In addition, the above phenomenon was not observed for PGO400V with smaller interlayer spacing of 0.39 nm, indicating that a large empty space is inevitable for hydrogen storage. This well corresponds to the theoretical prediction based on quantum calculation, although the size of empty space of PGO300V seems to be small. REFERENCES [1] [2] [3] [4] [5] [6] [7]

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 19

Fluoride-based electrolytes and their applications for intermediate temperature ceramic fuel cells Bin Zhua and Bengt-Erik Mellanderb a

Chemical Engineering and Technology / Chemical Reaction Engineering, Royal Institute of Technology (KTH), S-100 44 Stockholm, Sweden

b

Physics and Engineering Physics, Chalmers University of Technology, S-412 96 Gothenburg, Sweden 1. INTRODUCTION The current fuel cell (FC) technologies still have to meet critical challenges in order to reach a commercial success. Two major obstacles are the lack in fuel flexibility and the high costs for the components and systems. It may, for example, be noted that commercially available fuels are based on hydrocarbons, e.g. natural gas, petrol-based fuels, coal gas as well as biogas, while the existing fuel cell technology in most cases uses pure hydrogen as fuel. Among all FC technologies, the solid oxide fuel cell (SOFC) can utilize a number of different fuels. The conventional SOFCs use a ceramic electrolyte, e.g. yttria-stabilized zirconia (YSZ), and operate at high temperature, typically 1000°C. The high operating temperature puts very high demands on the materials and technology, which poses a major challenge for the further development of SOFCs into the market. In order to develop cost-effective SOFCs, much effort has been devoted to obtain a lower operating temperature (below 800°C) and intermediate-temperature solid oxide fuel cells (ITSOFCs, 400–700°C) have been developed either by using thin film technology to reduce the operating temperature of the yttria-stabilized zirconia electrolyte [1–5] or by using alternative electrolyte materials with high-ionic conductivity at reduced temperature, such as various doped ceria compounds, etc. [6–10]. All these efforts have, however, limitations due to the deficiency of technology and the stability of the material. Ceramic fuel cells (CFCs) are sometimes used as a more general term for fuel cells based on ceramic materials, which have the desired properties. For advanced ceramic fuel cells, the following properties are desired: (i) a wide fuel flexibility for many fuels, e.g. gaseous fuels, such as H2, different hydrocarbons, natural gas,

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biogas, liquid fuels, alcohol, gasoline, diesel as well as fuels that are originally in solid form of gasified coal and biomass; (ii) suitability for a wide range of applications for both stationary and mobile areas, e.g. power generation in both small and large scales, electrical vehicles, transportation and telecommunication devices, etc.; (iii) low costs for all components of fuel cell and stack. A fuel cell with these properties requires an operating temperature range that is optimized regarding material properties as well as reaction kinetics. A promising type of fuel cells is the so-called intermediate temperature ceramic fuel cells (ITCFCs). Compared with the high-temperature SOFCs, ITCFCs may be manufactured at a considerably lower cost. ITCFCs may also operate on readily available hydrocarbon fuels (e.g. coal gas, natural gas and biomass fuels for stationary power generation; alcohols (methanol, ethanol, etc.) or gasoline/diesel for tractionary power applications), and they may thus have a significant importance for a sustainable and environmentally responsible development in all parts of the world. The ability to use the existing infrastructure of hydrocarbon fuels is the most economic solution for the next generation (20 years) until a hydrogen-based infrastructure and economy can be created. The ability to operate on available logistical hydrocarbon fuels will soon be in fact, the most important advantage of this technique. ITCFCs may thus provide a new generation of fuel cells due to its inherent advantages. Due to the shortage of functional electrolytes, the development of advanced CFCs operating at intermediate temperatures has proceeded a long time ago, along a winding road. For example, in the late 1980s, the Swedish Energy Ministry started an innovative research project led by Prof. A. Lundén on intermediate temperature fuel cells based on oxyacid salts. The attempts to employ Li2SO4 as the electrolyte in a fuel cell can even be traced to the 1970s when Prof. Lundén’s group made the first fuel cell tests and obtained a current density of about 10 mA cm2 [11]. Later it was discovered that this fuel cell performance and the current output was contributed by proton conduction [12]. In order to improve the performance and the electrical and mechanical properties, further investigations of various salt-oxide composites were based on e.g. sulphates, phosphates, and nitrates [13]. The power density was improved by an order of magnitude to the 0.1 W cm2 level at 600°C during this period. In the late 1990s, new types of salts that did not contain oxygen were investigated. It is commonly believed that the oxygen element or oxygen ion acid groups are the origin for the proton conduction in salts. However, these salts, e.g. halides (chlorides and fluorides) were discovered to be fairly good electrolytes for fuel cells [14–16], even comparable with the oxyacid salt electrolytes. The possibility of proton or oxygen ion conduction in these materials opened a new interesting subject regarding both fundamental and applied research. The halide-based materials studied include pure chlorides, MClx (M  Na, Ba, Sr, x  1,2) and fluorides, MFx (M  Na, K, Ca, Ba, Sr, Pb, La, x  1–3) as well as solid solutions, e.g., MXNX2 (M  Li, Na, K, N  Ca, Ba, Sr, X  Cl, F). In addition, halides

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containing CaH2, e.g. MXxCaH2 (x  1,2) and MXNX2CaH2 and their composites with an oxide, e.g. Al2O3 were used as electrolytes and displayed a high proton conduction and successful fuel cell operation [17]. The anionic (F ion) conductivity has been extensively studied in the past since many fluorides are archetypical examples of fast anionic conductors [18–20]. In this chapter we instead focus on possible proton and oxygen ion conduction in fluoride-based electrolytes that may be of interest for fundamental and applied research. Also, our focus is to develop new advanced CFCs for intermediate temperatures. 2. MATERIALS AND EXPERIMENTS 2.1. Cell and ancillary materials

The following materials were used: (i) the pure fluorides, MFx (M  Li, Na, K, Ca, Sr, Ba, La, x  1–3) as well as MF-M Fx solid solutions (MLi, Na, K, M  Ca, Sr, La, x  1, 2, 3) (Aldrich) as electrolyte. Fluoride composites were prepared using alumina, Al2O3-type E (Merck, A.R.), as filler material. (ii) Platinum (Leitplatin 308A, Hanau, Germany) or silver (Leitsilber 200, Hanau, Germany) pastes or oxides with a layered rock salt structure, such as LiNiO2 and LiCoO2 as electrodes. The Ni- and Co-based oxides were used for the fuel cell anode and cathode, respectively. (iii) Stainless steel was used for the gas distribution/current collector plates. 2.2. Phase structure and microstructure analysis

Two types of structural analysis were mainly used to characterize the fluoride phases and microstructure. Powder X-ray diffraction (XRD) is the common analysis method to determine the crystal structure and electron spin resonance (ESR) or paramagnetic resonance (EPR) is probably one of the most powerful tools of solid-state materials, since it can be used to determine the nature of vacancies, defects or charge compensations. 2.3. Electrochemical analysis

The following electrochemical measurements are especially important for research on the fluoride-based electrolytes and applications with regard to proton or oxygen ion conduction, electrical properties and fuel cell applications. 2.3.1. Hydrogen concentration cell

The hydrogen concentration cell used in our investigations was constructed as (H2, anode) Pt/electrolyte disc/Pt (5% H2 in Ar, cathode), i.e. the electrolyte sample with platinum paste (Leitplatin 308A, Hanau, Germany) electrodes was placed in a cell providing different hydrogen partial pressures at the electrodes.

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The electric motive force (EMF) of this concentration cell can be measured. If we consider a simple case when the partial pressures of the residual water and oxygen contents in the gases can be kept the same at both gas sides, the situation can be simply described by the Nernst equation. The proton transport number can thus be calculated as Eobs 2F t(H)    ln(P/P) RT

(1)

where Eobs is the measured EMF value and Pand P are the hydrogen partial pressures at the electrodes. 2.3.2. Oxygen concentration cell

Using commercial O2 and Ar mixed O2 gases, the oxygen concentration cell was constructed as (O2, cathode) Pt / electrolyte disc / Pt (2% O2 in Ar, anode) The function is similar to that of the hydrogen concentration cell. If the partial pressures of the residual water and hydrogen contents in the gases can be kept the same on both the sides, the oxygen ion transport number can be determined by the Nernst equation, for oxygen concentration cell it is Eobs 4F t(O2)    ln(P/P) RT

(2)

where Eobs is the measured EMF value and P and P are the oxygen partial pressures at the electrodes. 2.3.3. Fuel cell studies and conductivity measurements

The I–V characteristics, i.e. the current density versus voltage, determined for hydrogen or oxygen concentration cells as well as for FCs can be used to characterize the transport properties of the electrolyte material and to estimate the ion conductivity [21]. In some respects, the conductivity obtained from the FC study may more correctly depict the electrical properties of the materials in practice, especially when considering that the measurement is made under the equilibrium conditions of different gas atmospheres resulting in different material stoichiometry. The conductivity obtained from FC measurements reflect the total conductivity for all transported ions, e.g., protons and oxygen ions as well as other reversibly transported species, e.g., hydride ions, since in the FC process all these ions are source ions provided by the reversible gas electrodes. 2.2.4. Quantitative study of proton and oxygen ion conduction

Measuring the quantities of water produced by the operating FC can be a useful method to separate different contributions to the ion transport. In our measurements, water was collected using a liquid-nitrogen cold trap. A constant

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current output was maintained from the FC for a certain time period and the water formed on both anode and cathode sides of the FC was collected and weighed using an accurate analytical balance (precision 106 g). In this way, proton and oxygen ion contributions may be separated since QH  2FWH2O(c)/MH2O

(3)

QO2  2FWH2O(a)/MH2O

(4)

where Q is the transported charge, WH2O (c) or WH2O (a) is the amount of water obtained at the cathode or anode side (in g), MH2O is the molar weight of water and F is the Faraday constant, (96,490 C/mol). In experiments on fluoride systems, it is also important to distinguish and determine the strong intrinsic F ion conduction. 3. FLUORIDE-BASED ELECTROLYTES 3.1. Structure properties

Most fluorides studied for the CFCs possess the fluorite structure. This structure exists also for a wide range of solid solutions based on NFx–MFy (N  Li, Na, K, M  Ca, Ba, Sr, La; x, y  1–3) systems. Some two-phase composites with other structure, e.g. LiF–MgF2, have been discovered to be highly conducting electrolytes as well. 3.1.1. XRD and phase structure

Many fluorides and their solid solutions have the fluorite structure, where the cations are situated at the lattice positions of a face-centered cubic lattice and 1 1 1 the F ions at (4 , 4 , 4 ) of the lattice constant from each cation, (see Fig. 1). In other words, each cation is located at the centre of a cube with 8 F ions at the corners, while each F ion is at the centre of a tetrahedron of cations. Solid solutions can be formed with some combinations of fluorides, such as LiF, NaF, CaF2, BaF2, etc. [22–24]. These solid solutions usually have a fluoritetype structure. Our studies also include composite fluoride systems, where LiF–MgF2 system is a typical example. Fig. 2 shows the XRD patterns for the LiF–MgF2 system, where the phases for LiF and MgF2 can be identified. The structure of LiF is NaCl type, and the structure of MgF2 is tetragonal. 3.1.2. ESR studies and microstructure

The ESR study was carried out for NF–MF2 (N  Li, Na, K, M  Ca, Sr, Ba), because of their high potential in CFC applications. In these studies, polycrystalline powder ceramic samples were used. Fig. 3 shows the ESR spectra of the NaF–CaF2 system treated in oxygen atmosphere. No ESR signal is detected for NaF or CaF2, but a signal appears for NaF–CaF2 samples. There are two basic

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Fig. 1. The fluorite structure.

M gF2

LiF(15%)

LiF(25%)

I

LiF(45%) LiF(60%)

LiF(80%)

LiF

20

25

30

35

40

45

50

2θ (Degrees)

Fig. 2. XRD pattern for mixed LiF–MgF2 system.

signals with effective g values of 1.9628 appearing at around a central field of 3200 G and g  2.1443 (at 3400 G). ESR spectra for Na LiF–CaF2 could be caused by unpaired trapped charge carriers, e.g. electrons or holes. The former, low-g-valued signal, would be related to an electron trapped in the F vacancy, the

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ESR signal

1NC 2NC

geff2

geff1

0

6NC 3.5NC

3000

3500

4000

H (Gauss)

Fig. 3. ESR spectra for NaF–CaF2 system treated in oxygen atmosphere.

so-called F centre; the latter to interstitial F ions, the so-called Vk centres. But the ESR signals related to the latter were not observed, since in NF–CaF2, doping with monovalent ions predominantly results in F vacancies; see the basic defect chemistry description in the following section. From Fig. 3 it can be seen that the signal intensity shows a strong dependence on the composition. The maximum intensity appears for the sample containing 10 mol% NaF. The ion doping for the NaF–CaF2 system takes place up to a certain NaF content. Above this composition, the system is a mixture of two fluoride phases. The 10 mol% NaF composition seems to be a limit for forming solid solutions; this causes the maximum intensity of the low-g-valued ESR signal. Above this level, the number of the effective F ion vacancies would not be increased. On the contrary, the increase in NaF, i.e. the decrease in CaF2 content, directly causes a relative decline in the vacancy concentration in the total NaF–CaF2 samples used for the ESR measurements, corresponding to the smaller percentage of CaF2 involved. The composition dependence of the intensities observed for the low-g-valued ESR signal suggests a strong relation to the F centres due to the fact that the fluorine ion vacancies act as the predominating defects. The low-g-valued ESR signal for NaF–CaF2 can be interpreted when the interaction between the trapped electron and the surrounding ions in CaF2 is considered. An unpaired electron is trapped at a negative ion vacancy surrounded by a tetrahedron of positive ions (also six nearest-neighbour F ions). The main magnetic interactions of an unpaired electron in CaF2 would be the hyperfine interactions with surrounding nuclear moments (the nonmagnetic Ca2 is not involved), i.e. F nuclei with a nuclear spin I  1/2. These hyperfine interactions would produce (2I  1)6  26 spin states distributed among 6  2I  1  7 energy levels for the F centre. The theoretical intensity ratios for this case are: 1:6:15:20:15:6:1 [25–27]. Therefore, the high-g-valued signal may be actually a

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degeneracy spectrum consisting of seven hyperfine structure signals caused by the spin–orbit interaction among F centre and nearest six F ions. The high concentrations of the F centres may cause a degeneracy of the signal. After r-irradiation, some F centres can be annealed or converted, then well-resolved resonance lines of the hyperfine structure was observed. This signal can be well interpreted by the hyperfine interaction between the unpaired electrons (trapped in the fluorine ion vacancies produced by the single valence Na replacing Ca2) and the surrounding NN (near-neighbour) fluorine ions. The doping and mixing as well as annealing treatment of the NaF–CaF2 make a strong inhomogeneous crystalline electrical field and a drastic broadening effect and degeneracy resonance line for the low-g signal related to the F centres. The nature of the high-g-valued signal is interesting. After treatment in hydrogen, we found that the high-g-value signal depended strongly on the sample compositions, while the low-g-valued signal remained unchanged. Fig. 4 shows ESR spectra for various NaF–CaF2 samples treated by hydrogen. The linewidth of the resonance for the high-g-valued signal shows a maximum at 35% mol NaF mixed with CaF2 (3.5NC), where the conductivity reaches the maximum as reported before [21]. The coincidence of the conductivity maximum and ESR signal width suggests that the high-g-valued ESR signal depends on the ionic transport properties. Further study and characterization to be carried on for this signal may provide us more useful information concerning the ion transport mechanism. The two major signals observed for CaF2-based electrolytes are of different nature. The low-g-valued signal has a strong structural character due to unpaired electrons trapped in fluorine ion vacancies, i.e. F centres. It reflects mainly the ion environment in CaF2 with the strongest density for the low NaF content. On

3.5NC

ESR signal

1NC MFx 0

6NC

2NC

1000

2000

3000

4000

5000

6000

H (Gauss)

Fig. 4. ESR spectra for NaF–CaF2 system treated in hydrogen atmosphere.

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the other hand, the high g signal shows a strong structureless feature and dependence on the atmosphere where the samples are treated; the nature and identification of this is not clear. A possible interpretation obtained from combining these results with the results of the electrical study is, however, suggested. The nature of this signal depends on the interfacial ‘‘structure’’ or interfaces between NaF and CaF2 phases. The composition coincidence between the maximum conductivity and the ESR signal linewidth suggests that this high g signal depends on the ion, more possibly with the hydride ion/hydrogen or proton transport. 3.2. Defect chemistry in fluoride-based electrolytes

Most fluorides are well known as anionic F ion conductors. Therefore, it is necessary to have a deep insight into the defect chemistry to understand the origin of protons or other possible ions, e.g. oxygen ion conduction in fluoridebased conductors. Both anions and cations can be involved in defect formation, i.e. the lattice of fluorides is able to dissolve different species such as O2, OH, Cl and probably S2 [28–30]. In addition, it is known that interactions with the fluoride can occur in gaseous H2O, hydrogen and H2S as well as in O2 and S2 at higher temperatures. From a defect chemistry point of view, both proton and oxygen ion conduction may be generated with different atmospheres as proposed below: in H2O:

H2O(g)  VF  2FFx  OF  2HF.F [V.F]  [HF.F]  [OF]

in O2:

1/2O2(g) 2FFx  OF  V.F  F2(g) [V.F]  [OF]

in H2:

1/2H2(g)  FFx  h.  HF.F [HF.F]  n

In MF2–NF (MCa, Ba, Sr, N Li, Na) systems, N dopant can generate additional F-ion vacancies NF → NM  V.F  F Fx [NM]  [V.F], where F Fx is a F ion in the lattice, V.F a fluorine vacancy, HF.F, a proton attached on the lattice fluorine ion; OF is a O2 at the F site and NM, a N ion at the M2 site. The replacement of monovalent N ions by bivalent M2 creates F-ion vacancies, which causes an enhancement of the conductivity. The defect structure

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and chemistry provide a reasonable explanation for the ionic transport mechanism and the observed phenomena in the fuel cells. From the defect chemistry point of view, it is easier to generate the H defect than the O2 defect in these composites. Protons are thus generally the major defects existing in the halide-based, especially, fluoride-based composite electrolytes studied in this work. Protons are faster mobile species than oxygen ions in the halide-based composites. Also, proton conduction is predominating in these gas concentration cells and fuel cells, due to the basic nature of the halidebased composite electrolytes used in fuel cell applications. 3.3. Electrical properties: Hybrid proton and oxygen ion conduction?

To study this subject, hydrogen and oxygen concentration cell as well as fuel cell measurements were performed as described in Section 2. The results obtained from the hydrogen or oxygen concentration cell measurements [30] may not be directly compared with each other, because the material stoichiometry strongly depends on the in situ gas atmosphere. Indeed the employed gas atmosphere, e.g. the hydrogen, oxygen or fuel cell gas atmospheres, can change the sample’s stoichiometry, ionic conductivity and transport properties based on consideration of the defect chemistry described in the above section. In the hydrogen or oxygen concentration cells, the concentration of the mobile species, e.g. H or O2, is not the same for either cases. For example, in the hydrogen concentration cell, only H is available from the external resources, while O2 ions are not available; the converse is true for the oxygen concentration cell. In each specific case, the H or O2 is missing, but they are both mobile if they can be sufficiently generated from the supplied gas resources. On the other hand, we may not judge tH tO2  1 for individual measured results from the hydrogen and oxygen concentration cell. The key point here is that the studied materials do not possess H and O2 inherently, since they are only generated from the provided external gas resources; see Section 3.2 above. The different proton and oxygen ion conduction behaviour may reflect the conduction and transport properties: (i) protons may be transported in the electrolyte and crossover the interfaces between the electrolyte and electrodes smoothly, resulting in a plateau in the discharging curves, while oxygen ions are hindered, causing a large polarisation and significant degradation of the cell voltage during discharge; (ii) protons are much faster than oxygen ions in transportation, so related proton transport processes having no polarization and the proton conductivity is much higher than that of oxygen ions. A basic consideration of proton and oxygen ion conduction in fluorides could be built as follows: protons and oxygen ions may form various dipoles/pairs and associates with anions, both lattice and interstitial F, or vacancies, resulting in unusually high mobilities and conductivities in these materials. Many of the defects are possible, z y e.g. MM , V.F, H.i, Fi, OF, Oi (My is the cation; y  1–3; z  or ; V, vacancy;

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i, interstitial). These defects can actually become oxygen ion or proton carriers by forming various pairs, e.g. H–F dipoles or (HVF)x and (OFVF)x associates. Protons can move flexibly in the lattice with the aid of different mechanisms, e.g. the Grothuss mechanism or channel/chain mechanism, when protons are formed as F–H or F–H–F dipoles/associates. In addition, both vacancy and interstitial mechanisms are possible for the proton and oxygen ion transport. In the fuel cell atmosphere, the situation is different from that in the hydrogen or oxygen concentration cells, since in the fuel cell both H and O2 are available from external hydrogen (anode) and oxygen (cathode) resources. The situation here is more complicated due to the different mechanisms for H and O2 defect generation and transportation. There is obviously a significant difference between oxygen ions and protons. The O2 can replace the F anion and occupy regular lattice sites, acting as normal ionic dopant. On the contrary, protons have very special qualities, being a bare nuclei with extremely small dimensions, 105 Å in diameter, and showing strong polarisation ability, etc. Protons are too small to occupy regular lattice sites. Instead they are easily attached or combined with anion groups or form proton – vacancy pairs. Therefore, in the fuel cell case, there is a competition between the H and O2 defects. From the defect chemistry point of view, the H defect is more easily generated than the O2 defect. Thus the H defect is usually predominating over the O2 formation, making the proton transport number reasonably much higher than that of the O2, while in the hydrogen and oxygen concentration cell case, only one defect formation process exists. As long as the external gas resources are available and the time for the defect forming process is sufficient, the general defect chemistry should be followed. The transport properties are therefore significantly different for the fuel cell, hydrogen and oxygen concentration cells. During the fuel cell operation, water was observed at both anode and cathode sides, which indicates that these fluoride and hydrofluoride-based electrolytes have both oxygen ion and proton conduction. By measuring the water quantity formed at both anode and cathode sides, about 90% of the water was found to form at the cathode side (caused by the proton conduction) and only 10% of water was formed at the anode side (O2 conducting) when using composite electrolytes with alumina fillers. 3.4. Ion conduction in fluoride composites

Some fluoride two-phase systems, e.g. LiF–MgF2, showed a dramatic conductivity enhancement compared with that of the individual phases, four orders of magnitude or more. For the LiF–MgF2 system, a high conductivity was observed in a wide composition region from about 18–60 mol% LiF. The composition dependence of the conductivity shows two peaks with rather similar conductivity values. The LiF–MgF2 system showed a maximum conductivity of about 102 S cm1 at 600°C for compositions between 20LiF–80MgF2 and

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60LiF–40MgF2 in molar ratio. Accompanying the conductivity enhancement, the activation energy decreased significantly for the two-phase material, from about 1 eV for single-phase LiF and MgF2 to an average of 0.4 eV for the LiF–MgF2 system. In a wide composition region, the activation energy is nearly unchanged, which may be an indication of an interfacial conduction mechanism in the grain boundary region of the two-phase material. A common route to obtain high ionic conduction in fluorides is to make fluoride solid solutions based on two or more single fluoride materials. For many fluoride solid solutions either fluorine ionic vacancies or interstitial fluorine ions are created, which significantly enhance the ionic conductivity compared with that of the single fluorides. However, in case of the LiF–MgF2 system, the process of Li ions occupying Mg2 sites to form a solid solution does not seem to be significant, since the LiF–MgF2 system always displays two separate phases according to the XRD measurements. The conductivity enhancement for the LiF–MgF2 system could probably be explained by a high conductivity in the interface layers between the grains, as has been observed for many composite materials [31–36]. The conventional two-phase composite system contains one phase with a moderate conductivity and another that is insulating. For this type of composite system there is usually only one conductivity peak. The double peaks observed for the LiF–MgF2 system may be due to two conducting components, LiF and MgF2, that are involved. 3.5. Hydride ion conduction ?

In the LiF–MgF2 system, the effects of hydride ions (H) were first discovered. Hydrogen can be incorporated into the fluoride structure either in the positively charged form, proton (H), which is bonded to the fluoride anions or as the negatively charged hydride ion (H) replacing the F ions in the lattice sites due to a rather similar ionic radius for the H and F ions. A rather low cell open-circuit voltage (OCV) was discovered, less than 0.9 V, in this system compared with the other fluoride system which regularly shows a cell voltage above 1.0 V. This phenomenon may be due to the H ion conduction that contributes to some degree in the LiF–MgF2 system, which is much less in the other fluoride systems. Introducing H conduction may help to understand the observed OCV phenomena, which will be discussed later. 4. HYDROFLUORIDE-BASED ELECTROLYTES 4.1. Defect chemistry and ionic conduction

In order to investigate the effects of coexisting protons in a fluoride lattice, hydrofluoride solid solutions are the best systems. Hydrogen bonds and acid – base properties of these and other protonated species should play a central role for proton transport in hydrofluorides.

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Proton conduction in hydrofluorides and their solid solutions has been discovered in a wide range of materials [37,38]. For example, consider a typical example, the simple hydrofluoride system: LiF–CaH2. The monovalent metal fluorides, e.g. LiF, may generate hydride ionic vacancies for the hydrofluorides. For example, LiF → LiCa  V.H H Hx [LiCa  ]  [V.H ], where H Hx is a H ion in the lat 2 tice, LiCa  is a Li ion at the Ca site and V.H a hydride ion vacancy. The hydride ion vacancies are responsible for the H and H conduction in the LiF–CaH2 (and its Al2O3) composite electrolytes, since protons can be transported via the H ion vacancies as they do via F vacancies in the fluorides. In addition, one should also address the influence of oxygen ions. It is actually unavoidable to introduce oxygen impurities in the materials during preparation. Oxygen ions may generate additional anionic, (H or F) vacancies in the hydrofluorides according to O2 → OA  VA. , where A is an anion, (H or F); hence oxygen impurities will also cause a vacancy mechanism for ionic transport. The composition dependence of the ionic conductivity for the LiF–CaH2–Al2O3 samples shows that the conductivity increases with the CaH2 content (or with the H vacancies) up to 15 mol% CaH2 [38]. A competition between proton and H ion conduction exists. The H ion conductivity may increase with CaH2 content above 15 mol% while the proton conductivity neither increases or decreases further with CaH2 content due to the trapping effect from the higher concentration of H ions. As a consequence, a nearly constant conductivity level or a slight peak approximately between 45 and 50% CaH2 is observed. On the other hand, the LiF–CaH2(–Al2O3) electrolytes are not behaving in the same way as the LiF–BaF2–CaH2(–Al2O3) electrolytes. In the latter case, F ionic vacancies may dominate during the preparation of the LiF–BaF2 solid solution via defect chemistry LiF → LiBa  V.F  F Fx and [LiBa]  [V.F], where F Fx is a F ion in the lattice, LBa is a Li ion at the Ba2 site, and V.F is a fluorine vacancy. The replacement of monovalent Li ions by the bivalent Ba2 creates F ion vacancies, resulting in an increase in the conductivity. In the LiF–CaH2(–Al2O3) case, the replacement of Li ions by the bivalent Ca2 creates H ion vacancies that represent the major defects, resulting in increase, in both H and proton conductivity. Based on the above proton, hydride and oxygen ion conduction may coexist in the hydrofluoride-based electrolytes. Nevertheless, proton conduction has been proved to be dominant over all other charge carriers for the hydrofluoridebased electrolyte fuel cells that will be discussed later. 5. FUEL CELL FUNDAMENTALS: PROTON, HYDRIDE AND OXYGEN ION CONDUCTING FUEL CELL PROCESSES From the point of view of ionic formation and transport in H2/air fuel cell devices, the most mobile and transported ions in fluorides are protons and possibly oxygen

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ions. In some cases, hydride ion (H) is possible in two-phase fluoride composites, while in the hydrofluoride electrolytes, protons (H) and hydride ions (H) are most mobile. Both can be provided by reactions at the hydrogen electrode and undergo their respective FC process. For the normal proton conducting process, the reaction at the hydrogen electrode is H2 → 2H 2e In this case, the electrons are left at the hydrogen electrode whose potential becomes negative. The hydrogen (fuel) loses electrons at the hydrogen electrode and forms protons, which move towards the oxygen (air) electrode through the electrolyte under the force of the proton concentration gradient. When the external circuit is connected, protons along with the incoming electrons from the external circuit combine with O2 at the air side to undergo the electrode reaction process and, as a result, water is formed at the air electrode. In this FC process, an electromotive force is formed and protons are moving from the low potential area (hydrogen electrode zone) to the high potential side (air electrode zone) in the electrolyte. The principle of the H conducting electrolyte in this FC process is shown in Fig. 5(a). The corresponding FC reactions are: at the anode, H2 → 2H 2e; 1 at the cathode, 2H 2 O2  2e → H2O; and the overall reaction of the cell. 1 H2 2 O2 → H2O. The theoretical OCV value for this cell process is about 1.2 V at room temperature. In the H ion conducting case, the situation is a little different. The hydrogen receives (instead of loses) electrons at the hydrogen electrode and converts them into hydride ions, H, i.e., H2  2e → 2H This reaction makes the hydrogen electrode positive since electrons are taken away from the hydrogen electrode to form the hydride ions. The movement of H ions from the high potential area (hydrogen electrode zone) to the low potential area (air electrode zone) in the electrolyte constitutes the fuel cell electromotive force. Fig. 5(b) displays the principle of the H conducting electrolyte. The relevant fuel cell processes are: at the hydrogen electrode, H2 2e → 2H; at the 1 1 air electrode, 2H 2 O22e → H2O; and the overall reaction, H2  2 O2 → H2O. Similar to the proton conducting FC case, water is produced at the air electrode. The electromotive force of the H ion conducting FC makes the H ions to be transported in the same direction as the protons, as shown in Fig. 5(a) and (b). The sheer difference between them is that the polarity of the hydride ionic fuel cells is opposite to that of the proton-type fuel cells.

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Air electrode (+)

433

Fuel electrode (-)

Air (O2) H+

H2

H2 O (a)

Air electrode (-)

Fuel electrode (+)

Air (O2) H-

H2

H2 O (b)

Air electrode (+)

Air (O2)

Fuel electrode (-)

O2-

H2 H2 O

(c)

Fig. 5. Illustration of H (a), H (b) and O2 (c) fuel cell processes.

In the oxygen ion conducting case, the oxygen receives electrons at the air 1 electrode and becomes O2 ions, i.e., 2 O2  2e → O2, making the air electrode positive since the electrons are consumed during this electrode reaction. The oxygen ions thus move from the air electrode (high potential) to the hydrogen electrode (low potential) forming the oxygen ion fuel cell electromotive force. Correspondingly, water is formed at the hydrogen electrode, i.e. 1 O2H22e → H2O, and the overall reaction is H22 O2 → H2O (see Fig. 5(c)). In summary, proton and hydride ion conduction (transportation) can produce water at the air electrode (cathode), but in the oxygen ion conduction case, water is formed at the hydrogen electrode (anode). Thus, depending on which side of electrode (the air or the hydrogen side) water is formed, information is obtained whether H, H or oxygen ion conduction is occurring. A simple judgement can be made as to reaction or ion (H or H) is dominating based on

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the signs of the cell potential. There is, however, always a thermodynamic equilibrium between the reactions for the hydrofluoride electrolyte fuel cell since two processes for H and H ions may exist simultaneously. The measurements on the fuel cell OCV are directly related to the transported ionic species. In the case of proton or oxygen ion transport, the cell OCV has the same polarity. For example, in the fluoride electrolyte fuel cells the OCV values are usually between 1.0 and 1.2 V. However, in the case of combined H and H conduction, the measured cell voltage is a sum of the net opposite potentials. A lower OCV value (0.8–1.0 V) with the polarity of the proton conduction was obtained for the hydrofluoride-based electrolyte fuel cells. It may be concluded that the H conduction should be predominating over the H conduction, although there is an influence from H ions to reduce the cell voltages. The hydride ions, on the one hand, has a negative effect to reduce the fuel cell voltage, causing a low-power output; on the other , by contributing H ionic conduction and transportation to the fuel cells, the total conductivities involved in the electrode reaction and fuel cell process are increased, so that the total current output can be enhanced. 6. FLUORIDE- AND HYDROFLUORIDE-BASED MATERIALS AS POTENTIAL ADVANCED CFC ELECTROLYTES Various alkaline and alkaline earth fluorides have been investigated as electrolytes for fuel cells. The required improvement in the current densities of the fuel cells was made by preparing MF–MF2–Al2O3 (M Li, Na, K and M Ca, Ba, Sr) composite electrolytes. Fig. 6 shows I–V characteristics for the LiNiO2 anode supported fuel cells with various alkaline and alkaline earth fluoride electrolytes. For the LiF–BaF2–Al2O3 electrolyte, the fuel cells showed a higher performance with a peak power density of about 110 mW cm2 at a current density of about 230 mA cm2 and at 750°C in the best case. As a comparison, using the electrolyte supported cell with the same electrolyte, the Pt anode, and the Ag cathode the cell reached a maximum power of only 28 mW cm2 under a current density of about 60 mA cm2 and a voltage of 0.5 V at the same temperature. The LiF–MgF2 demonstrated the best results, about 130 mW cm2 for a current density of about 300 mA cm2 at 780°C in the 50 mol% LiF 50 mol% MgF2 electrolyte case. Hydrofluoride-based electrolytes have also been studied extensively, involving simple hydrofluorides, LiF–CaH2 and its composite with Al2O3, and fluoride solid solution MF–MF2–CaH2 (M Li, Na, K, M  Ca, Ba, Sr) or the composites with Al2O3. For simple hydrofluorides like LiF–CaH2 (with alumina), a peak power of 152 mW cm2 at 400 mA cm2 was achieved at 750°C as the best fuel cell performance, which is inserted in Fig. 6 for comparison.

Fluoride-based electrolytes and their applications for intermediate temperature ceramic fuel cells

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P

1.0

Cell voltage (V)

0.8 100 0.6

0.4

50

0.2

BL 0

200

Power density (mWcm-2)

150

LCH

ML

0 400

600

Current density(mAcm-2)

Fig. 6. Typical I–V and I–P characteristics for fuel cells using fluoride-based, LiF–BaF2– Al2O3, LiF–MgF2, and hydrofluoride-based, LiF–CaH2–Al2O3 electrolytes. The LiNiO2-support technique was used for the fluoride-based electrolytes and Pt electrodes for the hydrofluoride electrolytes, respectively. The fuel cells were operating with H2 and O2 for the fluoride electrolytes and with H2 and 2% O2 in Ar for the hydrofluoride electrolyte, in all cases under a pressure of 1 atm.

7. ADVANCED CFCS BASED ON HYBRID FLUORIDE AND CERIA COMPOSITES The latest development concerns hybrid conductors based on proton conducting fluorides and oxygen ion conducting ceria, typically SDC (samarium-doped ceria). Fig. 7 shows the I–V characteristics for 80 wt% SDC20 wt% (40 mol NaF60 mol CaF2) electrolytes operated at 650°C. The fuel cell reaches power densities of about 300 mW/cm2, double that of the fluoride-based electrolyte fuel cells and this is achieved at a significantly lower temperature of 650°C, compared with 750°C in the case mentioned above. These new composite materials are thus promising and more developments are currently undertaken. 8. SUMMARY The extensively studied fluoride-based electrolytes for fuel cells possess a cubic fluorite structure, but the nonfluorite structured LiF–MgF2 two-phase composite is also interesting with its demonstrated higher ion conductivity. The two major ESR signals observed for fluoride-based electrolytes are of different nature. The low-valued g signal is strongly related to the ion environment in the fluoride structure with strong structural characteristics. The high

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Bin Zhu and Bengt-Erik Mellander

Cell voltage (V)

300 0.8 200

0.6

0.4 100

Power density (mW/cm2)

1.0

0.2

0.0

0 0

200

400

600

800

Current density (mAcm-2)

Fig. 7. Typical I–V and I–P characteristics for the fuel cell using the ceria-fluoride composite electrolytes at 650°C.

valued g signal shows a strong structureless feature but depends on the ion transport properties. Hydrogen and oxygen concentration cells as well as fuel cell measurements show that the fluoride-based materials have high proton and oxygen ion conduction with significantly different behaviour, strongly depending on the composition and gas atmosphere employed, which causes different stoichiometry and the subsequent diverse ionic transport properties. From the defect chemistry point of view, the H defect is easier to generate than the O2 defects in these composites. Protons are thus in general the major defects existing in the fluoride-based composite electrolytes more extensively studied. Protons are faster than oxygen ions in the fluoride-based composites, also proton conduction is predominating in the gas concentration cell and fuel cell measurements. The conjunction of proton conducting fluoride and oxygen ion conductors (ion-doped ceria) creates another new type of functional fluoride-ceria composite electrolytes, which has demonstrated a potential for future advanced ceramic fuel cell applications. REFERENCES [1] [2] [3] [4] [5] [6]

S. de Souza, S.J. Visco, and L.C. DeJonghe, Solid State Ionics, 98 (1997) 57. T. Hibino, H. Tsunekawa, S. Tanimoto, and M. Sano, J. Electrochem. Soc., 147 (2000) 1338. E.P. Murray, T. Tsai, and S.A. Barnett, Nature, 400 (2000) 648. P. Seungdoo, M.V. John, and J.G. Raymond, Nature, 404 (2000) 265. D Perednis and L.J. Gauckler, Solid State Ionics, 166 (2004) 229. B.C.H. Steele, P.H. Middleton, and R.A. Rudkin, Solid State Ionics, 40/41 (1990) 388.

Fluoride-based electrolytes and their applications for intermediate temperature ceramic fuel cells [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21] [22] [23] [24] [25] [26] [27] [28] [29] [30] [31] [32] [33] [34] [35] [36] [37] [38]

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J.P.P. Huijsmans, F.P.F. van Berkel, and G.M. Christie, J. Power Sources, 107 (1998) 71. C. Milliken, S. Guruswamy, and A. Khandkar, J. Electrochem. Soc., 146 (1999) 872. N. Maffei and A.K. Kuriakose, J. Power Sources, 75 (1998)162. R. Doshi, Von L. Richards, J.D. Carter, X.P. Wang, and M. Krumpelt, J. Electrochem. Soc., 146 (1999) 1273. B. Heed, B. Zhu, B.-E. Mellander, and A. Lundén, Solid State Ionics, 46 (1991) 121. A. Lundén, B.-E. Mellander, and B. Zhu, Acta Chem. Scand., 45 (1991) 981. B. Zhu and B.-E. Mellander, J. Power Sources, 52 (1994) 289. B. Zhu, J. Power Sources, 84 (1999) 39. B. Zhu, I. Albinsson, and B.-E. Mellander, Solid State Ionics, 135 (2000) 503. B. Zhu, Mater. Res. Bull., 35 (2000) 47. B. Zhu and X.R. Liu, Electrochem. Commun., 2/1 (2000) 10. D.R. Franceshetti, Solid State Ionics, 5 (1981) 613. D.R. Franceshetti, J. Schoonman, and J.R. MacDonald, Solid State Ionics, 5 (1981) 617. J.-M. Réau, S. Matar, G. Villeneuve, and J.-L. Soubeyroux, Solid State Ionics, 9/10, (1983) 563. B. Zhu, Electrochem. Commun., 1 (1999) 242. R.W. Ure, Jr., J. Chem. Phys., 26 (1957) 1363. M. O’Keeffe, Science, 180 (1973) 1276. J. Schoonman, Solid State Ionics, 5 (1981) 71. J. Arends, Phys. Stat. Sol., 7 (1964) 805. A.M. Stoneham, Proc. Roy. Soc. A., 306 (1968) 369. H.W. den Hartog and J. Arends, Phys. Stat. Sol., 22 (1967) 131. P. Fabry and E. Siebert, The CRC Handbook of Solid State Electrochemistry, P.J. Gellings and H.J.M. Bouwmeester (Eds.), CRC Press, Boca Raton, 1996, p. 329. S. Kumata, N. Miura, N. Yamazoe, and T. Seiyama, Chem. Lett., 6 (1984) 981. B. Zhu, Solid State Ionics, 145 (2001) 371. C.C. Liang, J. Electrochem. Soc., 120 (1973) 1289. P. Hatwig and W. Weppner, Solid State Ionics, 3/4 (1982) 249. K. Shahi and J.B. Goodenough, J. Solid State Chem., 42 (1982) 107. J. Maier, J. Phys. Chem. Solids, 46 (1985) 309. A. Bunde, W. Dieterich, and H.E. Roman, Phys. Rev. Lett., 55 (1985) 5. B. Zhu, Z.H. Lai, and B.-E. Mellander, Solid State Ionics, 70/71 (1994) 125. B. Zhu and X.T. Yang, Electrochem. Commun., 1 (1999) 411. B. Zhu and X.R. Liu, Electrochem. Commun., 2/1 (2000) 10.

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 20

The use of Nafion® as electrolyte in fuel cells Madeleine Odgaard IRD Fuel Cells A/S, Kullinggade 31, DK-5700 Svendborg, Denmark 1. INTRODUCTION The Proton-exchange membrane fuel cells (PEMFC) are electrochemical devices that efficiently convert chemical energy of the fuel directly into electrical energy. They operate like batteries and are similar in characteristics and components; the principle of operation is described in more detail in Section 2. Several types of fuel cells exist and are classified after the type of electrolyte used: e.g. solid oxide fuel cell (SOFC), molten carbonate fuel cell (MCFC), phosphoric acid fuel cells (PAFC) and PEMFC. The operating temperature of the fuel cells is determined by the electrolyte used. This chapter covers only the PEMFC; the explanation and application of the other fuel cell types is covered in several books and reviews published in the last decade[1–3] as well as historical reviews [4,5]. The development of PEMFC, also called as the solid polymer fuel cell (SPFC), has as the name indicates, been strongly related to improvements in performance of the polymer electrolyte membrane. The use of an ion-exchange membrane as electrolyte was first suggested by Grubb in 1957 [6,7], and the first fuel cell system based on a sulphonated polystyrene electrolyte was developed by General Electric in the 1960s for NASA for application as an on-board power source in the Gemini space program [8]. The fact that the polystyrene sulfonate membrane was not electrochemically stable, and only limited power density (less than 50 mW/cm2) was achieved with an excessive noble metal loading per cm2 of electrode (10–40 mg Pt/cm2) resulted in a focus on alkaline fuel cells (FCs) and only academic interest in the PEMFC. A major breakthrough in PEMFC was the invention of a perfluorinated sulfonic acid membrane (PFSA) introduced by E.I. du Pont de Nemours and Company in the 1960s under the name Nafion®, for application in the chlorine alkali industry. The advantage of the Nafion® membranes is their chemical stability. A specific membrane, developed to ensure the chemical stability toward

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strong sodium hydroxide and wet chlorine, which is a strong oxidising agent, also gave an increased lifetime to the PEMFC application. Development over the last 20 years, starting with work performed from the mid-1980s (under contracts by the Canadian Department of National Defense, DND), resulted in significant improved PEMFC performance [3]. Development of a similar perfluorinated membrane from Dow Chemical, tested in PEMFC at Ballard Power Systems gave improved power density. Groups at Los Alamos National Laboratory achieved a further step in the development by improving the utilisation of catalyst, i.e. the noble metal loading needed to obtain high power densities could be reduced. Today power densities close to 1 W/cm2 using low metal loading per cm2 electrode (⬃0.3 mg Pt/cm2) is demonstrated. Lifetimes of over 50,000 h of operation using the Nafion® membrane have been shown [9]. Presently, Dupont’s Nafion® is one of the most advanced commercially available proton-conducting polymer material and is the preferred electrolyte material for both hydrogen (H2-PEMFC) and direct methanol fuel cells (DMFC). PEMFC is seen as a system of choice for automotive systems for stationary application such as micro Combined Heat and Power (CHP) generators as well as for portable applications, e.g. power back-up systems. The PEMFC operates at low temperatures (60–80°C), allowing quick start-ups and immediate response to changes in the power demand. The use of fuel cells gives several advantages compared with conventional power generator systems. They offer a source of electrical energy that is continuous, environmentally safe, and additional benefits include low maintenance, excellent load performance, etc. In fuel cells, the chemical energy is directly converted into electricity, without preliminary conversion into heat. Consequently, this conversion is not limited by the Carnot cycle and 100% efficiency can be achieved. DMFC technology, using methanol derived from biomass or other renewable energy sources, gives the same advantages as PEMFC technology, e.g. highenergy efficiency, low or zero emissions. Although a lower performance than hydrogen fuel cell is achieved, the DMFC has a number of additional advantages: it does not have hydrogen storage problems, its fuel-supply infrastructure is cheaper than for hydrogen, and emissions are significantly lower than for petrolor diesel-fuelled generators. 2. PEMFC PEMFC is related to some of the earliest electrochemical discoveries. The first description of the fuel cell principle dates back to 1839; published by Friedrich Schoenbein and with the first demonstration of a fuel cell by William Grove in 1843 [10]. In spite of the fact that all essential fuel cell components and their various functions has been known for a long time, we still lack a lot of basic knowledge both at the molecular and nanostructural level. We also have to solve

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many of the basic materials and material-processing problems, essential for achieving the goal of a commercial PEM fuel cell without using rare and expensive materials and with long lifetime. Strong industrial and scientific efforts are being made in USA, Japan, and Europe to commercialise the PEMFC technology that converts primary energy into electricity and heat with a low or no pollutant emission within the next 10 years. 2.1. Principle of operation

PEMFC converts hydrogen and oxygen electrochemically into electrical power, heat and water. The nucleus of the fuel cell is the membrane electrode assembly (MEA), where the electrochemical reaction takes place. MEA consists of the ion-conducting polymer electrolyte membrane sandwiched between the anode and cathode: each containing a macroporous diffusion backing and active catalyst layer. Hydrogen is split into protons and electrons at the anode (the negative electrode). The polymer electrolyte membrane placed in the centre allows protons to pass from the anode to the cathode (positive electrode), while the electrons pass through the external circuit to the cathode. The electrons combine at the cathode with the protons that have crossed the membrane and with oxygen from the air, forming water (Fig. 1). The electrodes contain platinum or platinum alloys usually made of a porous mixture of carbon-supported platinum and platinum ruthenium alloy for cathode and anode, respectively, and ionomer (i.e. electrolyte) (Fig. 2). The catalyst particles must have contact to both protonic and electronic conductors and furthermore, these must have passage for reactants to reach the catalyst sites. The contacting point of the reactants, catalyst and electrolyte is conventionally referred to as the three-phase interface. The catalysts are usually deposited as nanoparticles on the high-surface area carbon. During fuel cell operation, a complex flow of reactants and reaction products exists in the pores of the electrodes. The cathode pores must allow gaseous oxygen to reach the catalyst surface and support efficient removal of the product, water, to prevent flooding of the backing or catalyst layer. The anode pores must provide efficient transport of fuel from the flow field to the catalyst surface. The current is proportional to the rate at which the reaction occurs. The performance is thus often quoted in terms of current density (current per cm2) at a fixed voltage. As long as fuel and oxygen (typically air) are supplied, the fuel cell will continue to produce electricity and heat. The only by-product when fuelled with hydrogen is water. A single PEM fuel cell (Fig. 3) in operation generates a voltage 1 V. Therefore, to obtain sufficiently high-voltage levels for a specific application the fuel cells are stacked in series. They are electrically connected by bipolar flow

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Madeleine Odgaard The electrode reactions occur according to the following scheme:

Anode: 2H2 → 4H + + 4e−,

E 0a = 0.00 V

Cathode: 4H+ + 4e− + O2 → 2H2O, E 0c = 1.23 V Cell: 2H2 + O2 → 2H2O,

E 0cell = 1.23 V

With respect to cell thermodynamics, the above electrochemical reactions should proceed spontaneously because the overall reaction has a negative free energy change (ΔG). The theoretical potential for the electrochemical reactions can be expressed by the Nernst equation, where the reversible cell potential E 0 is, when all components are in their standard states, RT ⎡ v⎤ ln ∏ ai i ⎥ zF ⎢⎣ i ⎦ However, in practice, the cell potential is a combined effect of the thermodynamics, kinetics, mass transport and ohmic resistance, and combination of anode and cathode losses reduces the cell voltage from the theoretical maximum of 1.23 V. The polarisation curves obtained for a fuel cell is thus divided into the three regions as shown below.

E theo = E 0 −

General Fuel Cell Performance ηactivation : Electrode kinetics - activation overpotential ηohm : Ohmic loss - contact resistances ηconc : Mass transport - limiting current Efficiency = E

cell/Etheoretical

1.23

E[V]

ηohm ηconc

100 % Efficiency

ηactivation

0 Current [A]

Fig. 1. PEM FC electrode reactions and general fuel cell characteristic.

plates that also distribute the fuel to the MEA within each cell. The electrical efficiency, ηel, is defined as the fuel cell voltage relative to the standard potential (E 0  1.23 V) of the cell reaction. The power density of a PEMFC depends on the operating conditions. A typical performance of a PEM single fuel cell using a Nafion® 112 membrane operating at 70°C and 2.5 bar absolute pressure is shown in Fig. 4. 2.3. The role of the polymer electrolyte

The major role of a polymer electrolyte in a fuel cell is to provide ionic conductivity. The electrolyte must also prevent the passage of electrons, act as a

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Proton Exchange Membrane~ 25-200 μm thick Anode Pt/Ru catalyst: XRD crystallite size ~ 3nm Surface area ~60m2/g

~10 μm thick 10nm-1μm pores

Cathode Pt catalyst: XRD crystallite size ~ 1-3nm Surface area ~30m2/g

~10 μm thick 10nm-1μm pores

Fig. 2. Schematic view of membrane electrode assembly (MEA).

Fig. 3. Presentation of a single-cell PEM FC.

barrier to the reactants and maintain chemical, thermal and mechanical stability. Chemical stability of the membrane in the oxidative and reductive electrochemical environments within the fuel cell is crucial. The membrance must be chemically inert to withstand possible degradations such as peroxide radical attach on the polymer end groups. In the operating fuel cell any gas or fuel permeating through the

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PEM FC SINGLE CELL performance: Tcell = 70°C, λair /λH2 = 1.5/2.0, ΔpAir = ΔpH2=1.50 bar (g)

0.70 0.63

0.8

0.56

0.7

0.49

0.6

0.42

0.5

0.35

0.4

0.28

0.3 0.1 0

0.21

Cell voltage [V] η el Cell power density [W]

0.2

0

0.1

0.2

0.3

0.4

0.5 0.6 0.7 0.8 0.9 Current Density[A/cm2]

1

1.1

1.2

0.14

Power Density [W/cm2]

Cell Voltage [V]

1 0.9

0.07 1.3

0.00 1.4

Fig. 4. Single cell performance obtained using a Nafion® 112 membrane. The cell operating at 70°C, at a pressure of 2.5 bar abs. using hydrogen and air with stoichiometry of 1.5 and 2.0, respectively.

membrane, often quoted as fuel crossover, is equivalent to an internal current (short) reducing the cell voltage. A reasonable mechanical strength and moderate dimensional changes when the electrolyte membrane is used for making the MEA, incorporating the MEA into a stack and during fuel cell operation is needed. 3. PROPERTIES OF THE NAFION® MEMBRANE 3.1. Introduction

PFSA membrane reviewed in this chapter deals with the Nafion® brand, developed and introduced by DuPont in the 1960s [11,12]. To date, Nafion® is one of the most extensively studied and most advanced commercially available proton-conducting materials. Nafion® was originally developed for use in the chlorine alkali industry, and the largest application of Nafion® membranes was in the early 1980s [13]. Nafion® membranes have many applications due to their high chemical and electrochemical stability, low permeability to reactant species, selective and high ionic conductivity and ability to provide electronic insulation. Apart from the use in chlorine alkaline cells, Nafion® membrances covers water electrolysis, gas separation, sensors, dehydration/hydration of gas streams, recovery of precious metals, salt-splitting and fuel cells [14–16]. Due to the fact that Nafion® is an expensive membrane, many attempts have been made to develop alternative materials. To date, Nafion® is still the preferred polymer electrolyte material for fuel cell application. This chapter aims at describing the structure and properties of the Nafion® membrane from a fuel cell application point of view. A clear understanding of the structure of Nafion® and

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its relation to the electrochemical properties is necessary for a full utilisation of the membrane and for future developments of the membrane. Transport of protons and associated water molecules is known to be strongly dependent on the membrane microstructure, although this is not well understood. 3.2. Chemical and physical properties 3.2.1. Perfluorinated membranes in general

In general, the perfluorinated membranes consist of a polytetrafluoroethylene (PTFE) backbone and side chains with acidic functionality (Fig. 5). The success of Nafion® has led to development of several variants of PFSA membranes since the 1960s by DuPont and other companies. The synthesis and preparation of the membranes developed by the different companies vary from place to place and has been continuously changed and improved. In the mid-1980s, significant improvements in fuel cell performance was achieved with a membrane developed by Dow Chemical [17] and tested at Ballard in fuel cells [18]. The Dow membrane is similar to Nafion® but prepared with shorter side chain, i.e. reduced equivalent weight EW 800 [19], corresponding to a higher SO3H concentration. The process and manufacturing of the Dow membrane were complicated and expensive. The rights to the Dow membrane were later taken over by DuPont. Other ion-exchange perfluoropolymer membranes, with the so-called long side chains were developed by companies such as Asahi Chemical that produces the Aciplex® membranes [20] and Asahi Glass Company, which produces the Flemion® membrane [21]. A very significant improvement in fuel cell performance was achieved by reducing the membrane thickness. The advantage gained with the thinner membranes has led to a new approach for making thin perfluorinated membranes. This implies impregnating a Nafion® solution into micro- or macroporous PTFE fabric, introduced by W.L. Gore and Associates under the tradename GoreSelectTM [22]. Many innovative attempts to develop alternative proton conducting membranes for fuel cell applications have taken place. The main motivations being (a) reduced membrane cost, (b) improved performance at higher temperatures, (c) lower requirement of humidification and (d) reduced methanol permeability. −(CF 2 − CF 2) x − (CF 2− CF) y (O−CF 2−CF) n−O−(CF 2) p−SO 3H CF3

Fig. 5. The general structure of the Nafion® membrane, with x  5 – 13, p  2 and y, n  1.

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The membranes vary from perfluorinated, partially fluorinated or nonfluorinated aliphatic polymers to polymers with an aromatic backbone (see Chapter 21 and [9,23–27]). 3.2.2. The Nafion® Membrane

The starting point for the DuPont technology is the perfluorination of the different monomers. The synthesis of Nafion® is based on the copolymerisation of tetrafluoroethylene (TFE) and a functional fluorinated monomer (vinyl ether) [28]. The sulphonic acid group functionality is introduced through the functional sulphonyl fluoride groups (SO2F). The general chemical structure of the Nafion® perfluorosulphonic acid ionomer polymer is given in Fig. 5. The length of the side chain, the composition of the polymer backbone and the processing of a film determine the final properties of the polymer electrolyte membranes [14]. The perfluorinated backbone provides chemical and mechanical stability, the ether groups provides flexibility, while the sulphonic acid groups yields high ionic conductivity [23]. The acid groups are fixed to the polymer and cannot leach out, while the counterions (H) are free to migrate and can be readily exchanged with other ions, according to the general reaction scheme –SO3–M1  M2 y –SO3–M2  M1 The ion-exchange capacity (IEC) of the polymer is directly related to the equivalent weight (EW). The EW is defined as the molar mass of the polymer per sulphonic acid group, 1 EW   IEC The “concentration” of the fixed ionic groups in the membrane determines the hydration of the polymer and hence the ionic conductivity and selectivity. As the length of the side chain decreases, i.e. lower value of EW, the number of sulphonic acid group per mass increases, resulting in higher conductivity. The desired EW is achieved by varying the ratio of vinyl ether monomer to PTE. The Nafion® is commercially available in different forms such as extrusion- or dispersion-cast nonreinforced membranes, powders, tubes and solutions. 3.2.2.1. Nafion® PFSA membranes

The first commercially available Nafion membrane was Nafion® 120 (1200 EW, 250 μm thick), prepared by an extrusion-cast membrane manufacturing process [29]. Since then DuPont has been active in developing the membrane’s performance by varying the EW and thickness. To date, the Nafion® membranes are manufactured and are available in different EW in the range between 1500 and 800, corresponding to ion-exchange capacities in

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Table 1. Nafion® properties, all values taken with membrane conditioned at 23°C, 50% relative humidity (RH). (MD – machine direction; TD – transverse direction [30]) Membrane Type

Typical thickness (μm)

N-111

25

N-112

51

NE-1135

89

N-115

127

N-117

183

NE-1110

254

Other properties Conductivity (S/cm)

0.083

Acid Capacity (meq/g)

0.89

Specific Gravity

1.98

Tensile Strength, max. (Mpa) Tear Resistance –— Initial (g/mm)

43 in MD, 32 in TD 6000 in MD, TD

the range 0.6–1.25 meq/g; the thickness is in the range 25–250 μm. Table 1 show the various membranes and their physical properties [30]. Nafion® 120 was followed by Nafion® 117 (1100 EW, 180 μm). Nafion® 115, 112, 111 and 105 are the latest developed membranes for fuel cell application. The most common membranes used today in H2 PEMFC is the thinnest Nafion® 112 and 111, and Nafion® 117 is most preferred for membrane DMFC (Section 4). To meet the demand of low-cost membranes, DuPont has developed a solution-cast manufacturing process for high volume production, aiming for automated MEA process technologies [31,32]. 3.2.2.2. Nafion® PFSA polymer dispersions

The early work by Raistrick [33] and Wilson and Gottesfeld [34,35] showed that significant improvement of fuel cell performance could be achieved by incorporating the electrolyte into the catalyst layer as a solution. This increased the utilisation of the catalyst and thereby to a considerable reduction of the noble metal loading needed [36]. Today dissolved Nafion® in terms of a dispersion is commonly used in the catalyst layers to prepare thin-film electrode layers [37,38]. DuPont has patented a solvent- [39] and water-based [40] high-pressure processes to convert the Nafion® polymer into dispersions. The dispersions have solid contents ranging from 5 to 20% by weight.

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3.2. Structure and electrochemical properties 3.2.1. Morphology

Nafion® membranes have been extensively studied to clarify the correlation between morphology and proton transport. Using a number of techniques, including small- and wide-angle X-ray scattering and neutron scattering [19,41–48], electron microscopy [49–51], AFM [52], NMR [53–58], ESR, IR and Raman spectroscopy [59–61], several models have been proposed to explain the structure of the Nafion® membranes [62–64]. Eisenberg [65] presented the existence of ion cluster and the most common model for Nafion® was proposed by Gierke et al. [62] and Hsuand Gierke [66,67]. The sulphonic acid group along with the adsorbed water in this model create the conductive pathway through the membrane. The acidic groups and the water molecules are separated from the perfluorinated backbone by forming ion clusters: inverted micelles (40–50 Å in diameter) are connected by narrow cylindrical channels (10–20 Å in diameter). Although substantial work in the characterization of Nafion® exists, and the presence of ionic clustering has been established, the characteristic properties of these domains as function of water content are not well understood [68,69]. It is generally accepted that fully hydrated Nafion® combines distinct regions within the membrane. There is a hydrophobic region containing the fluorocarbon backbone and a hydrophilic ionic region containing the sulphonic acid functional groups [46]. In the presence of water, this gives rise to some hydrophobic/hydrophilic nanoseparation. The cluster formation is facilitated by the flexible ether-linked side chains on the sulphonic groups, which have the spatial freedom for aggregation to form a hydrophilic network. The size of the clusters is dependent on the water content and the equivalent weight of the membrane. A high level of hydration produces large cluster dimensions [70]. When Nafion® membranes dehydrate, the size of the water clusters within the polymer microstructure decrease, leading to narrowing of the interconnecting channels [71]. An intermediate region exists between the two phases with some of the character of both the regions (Fig. 6). While the hydrophobic domain provides the membrane with mechanical stability and prevents the polymer from dissolving in water, the domain is responsible for the transport of protons and water [46]. 3.3.2. Proton Conductivity

The most important electrolyte membrane property for the PEMFC application is the proton conductivity. The membranes must provide an ionic pathway for the transfer of protons produced from the anode to the cathode. Low proton conductivity of the membrane results in high ohmic resistance. Nafion® is an excellent proton conductor with conductivity equal to a 1.0 M sulphonic acid. The Nafion® polymer behaves similar to an aqueous acid in the sense that it is limited to operate

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-

-

NAFION -(CF2-CF2)n-CF-CF2O-(CF2-CF-O)m-CF2-CF2-SO3H CF3

1 nm

: -SO3: protonic charge carrier : H2O wide channels more separated less branched

good connectivity small -SO3- /SO3separation pKa ~ -6

Fig. 6. Schematic representation of the Nafion® microstructure (Reprinted with the permission from Kreuer [46]. J. Mem., Copyright 2002, Elsevier.).

at temperatures below the boiling point of water. The conductivity of Nafion® is 0.08 S/cm at 25°C and 100% relative humidity (RH) [30]. The conductivity increases with temperature, giving a conductivity of 0.2 S/cm at 80°C and 100% RH. The proton conductivity of the polymer membrane has been studied extensively in the last 20–30 years, using AC impedance spectroscopy and DC techniques [20,45,54,70,72–86]. The most commonly used Nafion® membranes in PEMFC have a nominal dry thickness of 25–50 μm and are from the 1100 EW Series [70]. The reason for using thinner membranes is to achieve higher fuel-cell performance. The reduced ohmic resistance of the Nafion® membrane results in a significantly reduced slope in the pseudolinear region of the fuel cell performance curve, i.e. cell potential vs. current density. Single-cell performance with different Nafion® membranes from the 1100 EW series measured in a Ballard MARK5E single cell by Walsh et al. [70], is shown in Fig. 7. Several groups have examined the effect of temperature on the conductivity of Nafion® in detail. In general, the conductivity increases with temperature

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Fig. 7. The MEA performance obtained with Nafion® 1100 EW series of membranes with different thickness the cell is at 80°C and operating on H2O2 at 300 kPa abs. and 1.5/10.0 stoichiometry with full internal humification. (Reprinted with the permission from Walsh et al. [70]. Copyright 2002, The Electrochemical Society.)

but is also very dependent to the variation in water content, i.e. proton mobility, and hence the specific resistance is related to the hydration level [45,72,74,75,87–90]. At elevated temperatures, the membranes tend to dehydrate and the conductivity decreases. Another aspect of proton conductivity is a more fundamental understanding of the proton transport mechanism. The hydrated hydrophilic domains provide the high conductivity and is much dependent on the presence of water. Due to the hygroscopic nature of the sulphonic acid group, the Nafion® membrane adsorbs water. The water in the membrane exists in at least three different stages (Fig. 8). The first 2–3 water molecules (per sulphonic acid group) adsorbed by the polymer interact almost entirely with the acidic groups forming its primary hydration shell. These first water molecules are tightly bound to the ions. The next water molecules entering the matrix are more weakly bound and capable of solvating the protons, and the final stage is bulk water filled in the open-pore structure. Comprehensive reviews and papers covering the proton transport properties and mechanism of Nafion® combining experimental data with nonequilibrium statistical mechanical transport models has been published [46,68,69,91,92]. Proton transfer in solid polymer electrolytes follows two principle mechanisms: a Grotthus-like hopping and a “vehicle” mechanism [93]. The mechanism of proton conduction in fully hydrated perfluorinated membranes has been suggested by Kreuer et al. [92] to be comparable with that of liquid water. The proton

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22 NAFION 117, T = 300 K 20 18 water as second phase

n = [H2O]/[-SO3H]

16 14 12

loosely bound water

10 8 6 4 2

primary hydration of -SO3H

0 0.0

0.1

0.2

0.3

0.4

0.5 0.6 p / po

0.7

0.8

0.9

1.0

Fig. 8. Water absorption isotherm for Nafion® (Reprinted with the permission from Kreuer [92]. Copyright 1997, Elsevier.).

conductivity diffusion coefficient Dσ roughly follows the water diffusion coefficient DH O for Nafion® as a function of temperature and water content (Fig. 9). 2 The protons are mainly located in the central part of the hydrated hydrophilic nanochannels. In this region, the water is bulk-like and the proton transport is similar to protons in liquid water. Proton transfer is thought to occur along the hydrogen bonds of water network and Grotthus hopping becomes an important transport mechanism. For decreasing water contents, where the water is more bound through hydrophilic/ electrostatic interactions with the –SO3 group, the proton mobility decreases, and becomes dependent and co-operational with the water movement. This mode of proton movement is sometimes described as the “vehicle mechanism” since movement of protons occurs as complexes, e.g. H3O. 3.4. Water and methanol transport: a technological aspect

Water transport inside perfluorinated membranes is a complex phenomenon. The water balance of the membrane has been extensively studied [94–100] as the protonic diffusivity and conductivity of the membrane are dependent on the hydration level. In practical fuel cells, the membrane is typically operated in a partially dehydrated form.

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1e-4

80

60

DH2O

40

170 meV

20

NAFION 117

173 meV D / (cm2s-1)

1e-5

205 meV n = 10 238 meV

n=5

1e-6 n=3

DH2O Dσ 2.6

2.8

3.0 (1000/T) /

3.2

3.4

K-1

Fig. 9. Proton conductivity diffusion coefficient Dσ and the water diffusion coefficient DH O 2 for Nafion® as a function of temperature and water content. (Reprinted with the permission from Kreuer [92]. Copyright 1997, Elsevier.)

Factors affecting the water balance in PEMFC are (a) water adsorption from the vapour phase, (b) electroosmotic drag and (c) backdiffusion of water. The transfer of protons from anode to cathode is associated with the electroosmotic drag of water in the same direction. Each proton will drag at least one water molecule in water-vapour-saturated membranes and more (⬃2–2.5) in liquid-water-saturated membranes [74,95]. Without the addition of water to the anode, this would lead to dehydration of the membrane and reduced membrane conductivity. The dehydration is partially balanced by backdiffusion of water produced in the cathode. The water balance is highly dependent on the operating conditions of the PEMFC, such as temperature, pressure, humidity of the gases and cell current. Effective operation of PEMFC thus requires delicate control over the water balance in the entire stack. In contrast to H2-fuelled PEMFC, the DMFC anode is supplied with a liquid methanol/water mixture maintaining full hydration of the electrolyte membrane. A shortcoming of the Nafion® membrane related to the direct methanol fuel cell is its high methanol permeability, which drastically reduces the performance [101]. Water is needed for anode reaction in DMFC (Section 4). This implies that pure methanol cannot be used, but a mixture of methanol/water is necessary.

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Methanol mixes very easily with water and readily diffuses through the Nafion® membrane as well as other commercially available proton-exchange membrane. This results in low fuel utilization, low cell voltage and excess catalyst loading (Section 4.3). 3.5. Other properties facing the fuel cell requirements

Besides being a good proton-conducting membrane, the fuel cell must also serve as an electronic insulator. Any gas or fuel permeation from anode to the cathode through the polymer causes an internal chemical short circuit in the fuel cell and decreases the power and energy densities. If hydrogen/methanol are simultaneously present at the cathode of the fuel cell, reactions run parallel and the net current is the sum of the anodic oxidation and the cathodic oxygen, and the “fuel crossover” is often expressed as mass flow, and converted into the corresponding parasitic electronic current using Faraday’s law [102]. Investigations of mass-transport parameters through Nafion® membranes, such as permeation, diffusion and solubility of different gases and methanol, has therefore been carried out by several authors [72,78,103–105]. Gas permeation is found to be a function of the relative humidity or the water content in the membrane, and the permeation can be separated into a solubility and diffusive component in the aqueous and the hydrophobic perfluoro phase of the membrane [104]. It is also found that permeability through the water-saturated sample is higher than that through the sample equilibrated with water vapour. From the literature, it is known that when Nafion® takes up water from the liquid phase, the number of moles of water per mole of sulphonic acid is significantly higher than the number obtained by equilibration with saturated water vapour. This so-called Schroeder’s paradox was explained by Zawodzinski et al. [72], as due to the difficulty in condensing vapour within the pores of the membrane. Due to the varying water contents different permeability is expected. The pronounced dependency of the permeability also indicates that transport of the gases primarily takes place through the hydrophilic regions in Nafion® [103]. To commercialize PEMFC technology, one of the most critical criteria is to achieve long lifetime of the overall system, especially for the residential stationary power market, requiring durable performance for many years of operation. Nafion® has demonstrated durability in fuel cells up to 50,000–60,000 h of operation [23]. Stability under oxidative and reductive conditions and their relation to the corresponding fuel cell stability is difficult to establish, since the origin of membrane degradation is not fully understood. Evidence of membrane thinning and fluoride detection in the product water indicates that the polymer undergoes chemical attack. Peroxide radical attach on polymer end groups [13,106] with residual H-containing terminal bonds is generally believed to be the principle degradation mechanism.

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4. APPLICATION AND PERFORMANCE OF NAFION® IN FUEL CELLS 4.1. Introduction

The most unique feature of the fuel cell technology is its flexibility; the fuel cell is modular and the power is therefore scalable, with the potential to find application in products over a wide power range from few mW to MW. In contrast to liquid acid electrolytes, Nafion® membranes have the advantage of being a solid polymer, which makes them easy to handle. The technology has now reached a stage in its development where commercial exploitation seems to be a reality, although some basic material and material-processing problems are yet to be solved, being the significant focus over the last 20 years on adapting the PEMFC for use in consumer and industrial applications. Papers and reports including surveys and summaries of the current technological and commercial status have recently been published [24,107–112]. The PEMFC is seen as a system of choice for automotive systems, for stationary application such as micro Combined Heat and Power (CHP) generators as well as for portable applications, i.e. power back-up systems. The PEMFC operates at low temperatures (60–80°C), allowing quick start-ups and immediate response to changes in the power demand. The use of fuel cells has several advantages compared with conventional power generator systems. They offer a source of electrical energy that is continuous, environmentally safe and additional benefits include low maintenance, excellent load performance, etc. An analysis of the current energy consumption and associated emissions in the industrialized countries shows significant contributions from heating and power plants in the form of CO2 and SO2. This calls for not only the development of more advanced heat and power technologies, but also the use of alternative fuel supplies. Advanced technologies will help in solving this problem through, for example, the use of excess electricity from renewable energy sources for the production of hydrogen via the electrolysis of water. A lifecycle analysis of hydrogen thus produced leads to global-warming emissions, which are a whole order of magnitude less than for fossil fuel produced hydrogen, even after considering the effort needed to produce these systems and to transport the energy produced. However, such an environment friendly technology is possible, but has not been developed so far. PEMFC today are therefore fuelled by hydrogen from reformed fossil fuels that are at present the most realistic possibility. Fuel cells have the potential to reduce the fuel consumption. They have a higher efficiency than even most efficient combustion alternatives, resulting in fuel savings between 15 and 50%, depending on the application, and a significant reduction in CO2 emission. Another attractive solution is to use liquid methanol as fuel. DMFC technology, using methanol derived from biomass or other renewable energy sources, gives the same advantages as PEMFC technology, e.g. high-energy efficiency and

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low or zero emissions. Although it has a lower performance than the hydrogen fuel cells, the DMFC technology has a number of other advantages: it does not have hydrogen storage problems and is easy to store, its fuel supply infrastructure is simpler than for hydrogen and emissions are significantly lower than petrol- or diesel-fuelled generators. Eliminating the need for a reformer also facilitates technical simplifications, and thereby reduces the size and weight of portable powergenerator systems. This leads the DMFC well to applications, where the energy density must be high. Fuel cells can substitute batteries entirely or can be used in a fuel cell/battery hybrid, where the fuel cell acts as the battery charger. 4.2. The PEMFC stack

A single PEM fuel cell in operation (Fig. 2, Section 2.1) delivers a voltage

1 V. To obtain practical useable voltages fuel cells are stacked in series. High current is realised by enlarging the active area of the membrane electrode assembly. The cell typically consist of bipolar plates, which are pressed against the membrane electrode assembly. The bipolar plates, providing the electrically conductive path for the generated current to the adjacent cell, also serves as distributor of the reactant fuel and oxidant to the entire surface of the electrodes having a manifold of grooves. The choice of material for commercial PEMFC stacks to some extent, is dictated by several factors to some extent depending on the specific market. These factors not only concern performance, like current conduction, heat conduction and mechanical strength, but also lifetime and cost issues. The bipolar plate material differs from sheet metal, graphite foil, and graphite polymer composites, which are all potentially low-cost materials. By volume, the bipolar plates make up the largest part of the PEMFC, and for some applications, lowdensity materials are crucial for maximising the power to weight ratio (kW/kg). For applications aiming long lifetime, graphite materials provide good chemical resistance, whereas metals have shown some limitations due to low corrosion in the harsh electrochemical environment inside the fuel cell stacks [24,113,114]. Other stack components include seals, cooling elements, current collectors and end plates. End plates give the fuel cell stack mechanical stability and enable sealing of the components by compression. Fig. 10 shows an example of a fuel cell stack made up of 70 cells connected in series. The power output of a PEMFC, using MEAs based on Nafion® membrane, is highly dependent on the experimental conditions. Gas pressure and utilisation are the most important factors assuming a perfect gas humidification and an optimal stack temperature. Elevated pressure is preferential from a performance point of view because of an easier water management especially at high currents. Fig. 11 shows the performance of a 70-cell PEMFC stack run operated at a cell temperature of 70°C, 85% fuel utilization and 50% oxygen utilization at ambient pressure and elevated pressure of 2.5 bar.

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Fig. 10. A 70-cell PEM FC stack by IRD Fuel Cell A/S designed for stationary applications such as uninterruptible power supplies (UPS) and combined heat and power (CHP). The bipolar plate is graphite based for long lifetime. The stack end plate-manifolds have a polymerbased sandwich construction which, together with polymer frames around the single cells, ensures thermal insulation and thermal management of the stack.

Effective water and heat management have a major impact on fuel cell performance [115,116]. Thermal management is required to remove the heat in order to prevent excessive operating temperatures, i.e. prevent dehydration of the Nafion® membranes. Due to chemical degradation, Nafion® is limited to a maximum operating temperature of 130–140°C. Nafion® exhibits the highest proton conductivity when fully hydrated, this implies humidification of the gases. Increasing the temperature above 100°C requires a pressurised system to maintain a water-saturated environment. The water management is not only to ensure the membrane remains fully hydrated, but also to prevent excessive water accumulation within the cells and stack manifolds. Water flooding can block the oxygen transport to the electrode. 4.3. Application of PEMFC

Worldwide a number of prototype fuel cells and fuel cell systems have been built and demonstrated. The common driving force for applying fuel cell

The use of Nafion® as electrolyte in fuel cells

IRD Fuel Cells: PEMFC70 2.00 SN 003 Polarisation curve Tstack = 70°C, λair = 2.0, λH2 = 1.5, 100 % RH

70 60

457

6000

5000

50 40 3000 30

Power [W]

Voltage [V]

4000

2000 20 Voltage @ 1.5 bar (g) Voltage @ 0.0 bar (g) Power @ 1.5 bar (g) Power @ 1.5 bar (g)

10 0 0

10

20

30

40

50

60

70

80

90

1000

0 100 110 120 130 140

Current [A]

Fig. 11. The voltage and power density of a 70-cell IRD Fuel Cells A/S PEM FC stack.

technology is the overall environmental benefit to conserve energy with high efficiency, and reduce chemical and noise pollution on both a local and a global scale. Some applications will benefit from the use of the waste heat, and others consider the main advantage being their potential use to back-up or replace batteries, allowing extended runtime for portable electronic devices due to high energy density. With the rapid developments within this field, recent analysis and extensive reports covering the fuel cell industry in a series of market and regional surveys is available on Fuel Cells Today [www.fuelcelltoday.com], which is a free global internet portal. The biggest challenge for the developers is a substantial reduction in the cost of the fuel cell systems. The long-term target price of EU (FP6) for stationary PEMFC systems is 100 €/kWe and 50 €/kWe for automotive PEMFC systems. The present high cost is mainly ascribed to the use of materials such as the Nafion® membranes, the catalyst materials and also components like graphitebased bipolar plates that are expensive due to the present piece by piece production and the limited number of component manufactures. The stringent requirements in terms of compactness, high-energy density, performance stability and low cost will change the research direction towards optimising the different aspects of PEMFC system.

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4.3.1. Automotive applications

Governmental legislation such as emission regulations especially in regions with air-quality problems, e.g. urban areas, has created considerable interest in using fuel cells for providing power for buses. In terms of volume the automobile market represent the largest opportunity for high volume production of fuel cells and systems. A challenge for the Nafion® price development is to meet the requirement of low price in the automotive market. To meet the EU price target of 50 €/kWe for the total PEMFC system, it is anticipated that the membrane electrode assembly including the catalyst must be available for less than 10 €/kWe. The first step in this development is the hybrid electric vehicle (HEV), where the fuel cells provide the base load and batteries providing the peak power. The second step is the fuel cell vehicle (FEV) where the fuel cells provide all the primary motive power. The thinnest Nafion® membranes are commonly used in fuel cell stack for automotive application due to high power density demand; the thin membranes provide the lowest resistance. A consequence of the high power density is a shortening of the fuel cell lifetime. PEMFC perform best when pure hydrogen is used as the fuel. The simplest and most efficient vehicle will use hydrogen gas, supplied from a liquid or gaseous reservoir on-board. Small scale, on-site production of hydrogen to provide fuel for buses in large city areas have been demonstrated [117]. For practical applications, on-board generation of hydrogen by reforming natural gas, hydrocarbons and alcohols appear more realistic in the near future [118,119]. The reformatted gases produced by these methods contain, besides hydrogen and CO2, small amounts of CO as by-product. CO has a strong tendency to adsorb on the catalyst surface, blocking the sites needed for the hydrogen reaction to occur [120–123]. For automobile application, a shortcoming of PEMFC based on Nafion® membranes is that their range of operation is limited to temperatures conventionally around 60–80°C. The incentive to develop higher temperature membranes is not only enhanced CO tolerance and reaction activity, but also easier thermal management with potential to lower the additional volume for the total system. 4.3.2. Stationary application

The use of fuel cells to generate heat and electricity to single households has started only lately and is virtually non-existent in the European market. An example of a 2 kW micro-CHP unit is shown in Fig. 12. The unit is designed to operate in parallel with the existing power grid. The power is sized for an average heat needed in a north European single-family house. The key component in the CHP system is the PEMFC stack, using MEAs based on Nafion® membranes as electrolyte. The complete CHP system comprises the following components: ● ●

PEM FC stacks Cell voltage monitor system (CVMS)

The use of Nafion® as electrolyte in fuel cells

A. Front view of the 59*63* 167 cm3 CHP cabinet. B. View into the cabinet. The generator consists of the following components: 1. PEM FC stack 2. CVMS 3. A fuel supply (not shown) 4. Air supply (compressor), frequency transformer for electronic flow control (4b) 5. Humidifier 6. FC cooling loop with a heat exchanger 7. A grid connected DC→AC inverter 8. An electronic control unit (not shown) 9. An independent hydrogen safety system (not shown)

Fig. 12. Micro-CHP unit developed by IRD Fuel Cell A/S.

● ● ● ● ● ●

Hydrogen fuel supply system Air supply system (compressor) Humidifier Cooling system Grid connected DC→AC inverter FC control unit system

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Larger power output can be obtained by connecting several systems in parallel giving 4, 6, 8 kWe, etc. The electrical efficiency is approximately 45% a nominal load, but can be varied according to the load. The electrical efficiency of the total system is around 35% due to the parasitic power used by the balanceof-plant (BoP) components, in particular the air supply. Stationary FC systems incorporating PEM technology is expected to evolve into three different markets: large-scale power generation, residential and smallscale stationary power generation. The advantage of using PEMFC technology beside environmental benefit is their low maintenance. PEMFC for large power generation (100 kW) have been demonstrated; however the efficiency of PEMFC seems to be insufficient for large-scale baseload applications, which probably will be the domain of other fuel cell types such as SOFC, PAFC and MCFC technology. PEMFCs are generally believed to cover the lower end of the power output scale, where fuel cells are expected to be economically viable. The progress of using PEMFC in small-scale stationary power units between 1 and 10 kW has grown substantial. The exact size of the fuel cell system needed differs within geographical areas and their respective needs. Unlike North America, Europe and Japan will show the highest demand for FC systems in the segment of microCHP [124]. Here, PEM fuel cells are particularly attractive and will gain a lead over the more efficient SOFC technology that is expected to dominate the US market, where the main application will be at locations not attached to the grid. In Japan, where the electrical grid is generally reliable and extends over the entire country, units of ⬃1 kW is of interest [125]. 4.3.3. Portable application

Application of small PEMFC systems (ranging from a few W up to ⬃1 kW) for portable devices provide a potential alternative to batteries and diesel or gasoline generators. The use of fuel cells is motivated by several factors including high energy to weight ratio and long run time for portable electronic equipment [126,127]. The fuel could be hydrogen from a metal hydride storage system, sized to meet the desired run time between fuelling or a liquid methanol/water mixture [128]. The hydrogen fuel cells give high power density and would be excellent for portable power, but complicated because of hydrogen storage and the fact that hydrogen distribution to the public today remains unsolved. DMFC is a simpler and close to future possibility. 4.4. DMFC and their applications

DMFC technology has reached a stage in its development where commercial exploitation is now seen as a reality. However, one major challenging problem to be solved for the practical application of direct methanol fuel cells is the fact that it suffers from severe fuel crossover through the polymer electrolyte,

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Nafion®. The crossover of methanol from anode to cathode causes a partial chemical short circuit in the fuel cell and decreases the power and energy density. Even at low methanol concentrations at the anode, sufficient methanol and water diffuse through the electrolyte to significantly impact the performance of the cathode. Any methanol reaching the cathode chemisorb on the electrode surface with subsequent oxidation to CO2 at the high potentials reducing the cell voltage. The overall increased demand for electrical power and equipment with convenient rechargeability for high-technology electronic applications leads to applications of DMFC, where the energy density must be high. Fuel cells can substitute batteries entirely or partly as in a fuel cell/battery hybrid, where the fuel cell acts as the battery charger [129,130]. Despite energy-conversion efficiencies in the range 30–40%, fuel cells still has an overall advantage over batteries in energy density terms. The specific energy of methanol is 4.28 Wh/L, making it viable to have DMFC system 1000 Wh/L or Wh/kg and giving considerable savings in weight and volume compared with 350–400 Wh/L or 150–200 Wh/kg, for the best battery systems. 4.4.1. The direct methanol fuel cell (DMFC)

The structures and materials of the MEAs used in PEMFC and DMFC are very similar (Section 2.1). In DMFC the electrode reactions can be represented by the following overall cell reaction as shown in Fig. 13. The practical cell voltages obtained using liquid methanol are considerably lower, and the losses are higher than for the hydrogen feed PEMFC. Reviewing the literature [109,131,132], the relative low performance compared with the hydrogen PEMFC is thus caused by ● ●

poor kinetics of the methanol reaction methanol permeation through the Nafion® membrane.

The anode reaction shows that water is needed with methanol to enable the oxidation reaction. This implies that pure methanol cannot be used, but a mixture of methanol/water is necessary. The methanol oxidation reaction is quite complicated, in which surface intermediates play a key role [131,133,134]. The cathode reaction in DMFC is similar as that for the hydrogen feed fuel cell; the protons and electrons recombine with oxygen. Studies examining the extent and impact of methanol transport, in DMFCs using Nafion® membrane, as a function of various operating conditions have frequently been published [97,102,135–138]. Undesired methanol transport through the membrane takes place by diffusion and electro-osmotic effects. Methanol arriving at the cathode catalyst will react with the oxygen present to form CO2. Methanol reaching the cathode therefore causes a depolarisation due to the competing electrochemical processes of oxygen reduction and methanol oxidation.

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In the DMFC the electrode reactions can be represented by the following two half cell reactions: Anode: CH3OH + H2O → CO2 + 6H+ + 6e−,

E 0a = 0.046 V

Cathode: 6H+ + 6e− + 3/2O2 → 3H2O,

E 0c = 1.23 V

and the following overall cell reaction is represented by the equation Cell: CH3OH + 3/2O2 + H2O → CO2 + 3H2O,

E 0cell = 1.18 V

DMFC SINGLE CELL PERFORMANCE: 1.0 M MeOH, λMeOH/λair = 6/3, Tcell = 70°C, ΔpAir = ΔpCH3OH = max 40 mbar (g) 0.9 0.180 Cell voltage [V] 0.8 0.160 Cell power density [W] 0.7 0.140 ηel 0.6 0.120 0.100 0.5 0.080 0.4 0.060 0.3 0.040 0.2 0.020 0.1 0.000 0.0 0.00 0.10 0.20 0.30 0.40 0.50 0.60

Power Density [W/cm2]

Cell Voltage [V]

The obtained performance, i.e. voltage, power density and efficiency vs. current density for a DMFC depends on operating conditions such as temperature, fuel concentration, fuel/air stoichiometry and operating pressure. Fig. below shows the performance obtained for a singlecell DMFC using a MEA based on Nafion® 117 membrane; measured at 70°C using a 1.0 M CH3OH liquid anode fuel.

Current Density [A/cm2]

Fig. 13. DMFC electrode reactions and single-cell performance.

The rate of methanol permeation decreases with increase in current density. At higher currents, the methanol concentration is reduced due to consumption at the anode/electrolyte interface in the porous electrode structure. This reduces the driving force for diffusion through the membrane. In practice, the effects of methanol crossover is counteracted by careful MEA and system design. The state-of-art facilities use the thicker Nafion® 117 (180 μm), in contrast to the Nafion® 112 (50 μm) used in the hydrogen feed PEMFC. Although the thicker membrane increases the cell resistance, the gain from reduced methanol crossover results in improved performance. The fuel must be supplied as a diluted methanol solution (typically 0.5–2.0 M) as a consequence of the high methanol crossover, adding the size and complexity of the system [131]. 4.4.2. Applications of DMFC

The use of portable devices worldwide has grown dramatically in recent years and continues to do so. Also, the overall increased demand for electrical

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power, estimated at 10%/year in the US (similar increases have been foreseen in Europe), has resulted in a corresponding increase in demand for uninterrupted power supply (UPS) systems. This market is currently dominated by batteries combined with gasoline- or diesel-fuelled generators, which pollute excessively through their high emission levels and noise. Current trends towards higher fuel prices and greater environmental awareness strongly supports a shift in consumer demands. Fig. 14 shows a 1 kW DMFC stack developed and designed for UPS application. The stack consists of 51 cells connected in series. The MEA is based on Nafion® and gives a maximum power density of 120 mW/cm2, operating at a temperature of 70°C. The flow distribution/interconnect plate is graphite-based for long lifetime. The cells and stacks are designed with a low-pressure loss reducing the power used for the auxiliary pumps, etc. The electrical efficiency is 35% at current of 36 A but can be varied according to the load (Fig. 15). Unlike a hydrogen feed PEMFC, the direct methanol fuel cell does not require components such as fuel processor, humidifier or cooling system. The advantage of methanol fuel cell compared with conventional batteries is the prospect of longer run time and the potential for instantaneous refuelling. This makes them attractive and well suited for use in electronic equipment such as portable computers, mobile phones and other handheld electronic equipment. Other potential applications are powering remote telecommunication

Fig. 14. 700 Wel DMFC stack developed by IRD Fuel Cells A/S.

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I/V Performance Curve: DMFC_ 51 Cells Tstack ~ 70°C, λAir = 4, λMeOH = 6, 1.0M MeOH 1250

50 45

1000

35 750

30 25

500

20 15

Stack Voltage at ambient Air pressure [V] Stack Power at ambient Air pressure [W] Stack Voltage at 0.5 Bar(g) Air pressure Stack Power at 0.5 Bar(g) Air pressure

10 5 0

0

10

20

30 Current [A]

40

50

Stack Power [W]

Stack Voltage [V]

40

250 0 60

Fig. 15. I/V performance graph of a 51-cell DMFC stack developed by IRD Fuel Cells A/S.

transmission equipment and remote scientific investigation equipment, emergency AC power for hospitals and other consumer leisure applications such as camping, sailing, etc. The challenge, especially for FC products in the mW range, is the miniaturisation, as compact design are essential in power sources for portable application. The DMFC has been considered as a possible technology for automotive application but is not as developed as the PEM and concerns related to implementation of a methanol infrastructure makes the development slow [117]. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12]

A.J. Appelby and R.L. Foulkes (Eds.), Fuel Cell Handbook, Van Nostrand, New York, 1989. K. Kordesch and G. Simader (Eds.), Fuel Cells and their Applications, VCH Verlaggsgesellschaft GmbH, Weinheim, Germany, 1996. J. Larminie and A. Dicks (Eds.), Fuel Cell Systems Explained, 2nd., Wiley, New York, 2003. K. Kordesch, J. Electrochem. Soc., 125 (2004) 77C–91C. A.J. Appelby, J. Power Sources, 29 (1990) 3–11. W.T. Grubb and L.W. Niedrach, J. Electrochem. Soc., 107 (1960) 131–135. W.T. Grubb, General Electric Company, US Patent 2,913,511, November 17,1959. W. Vielstich, Fuel Cells: Modern Processes for the Electrochemical Production of Energy, Wiley, New York, 1970. O.Savadogo, J. New: Mat. Electrochem. Systems, 1 (1998) 47–66. U.Bossel, The Birth of the Fuel Cell, European Fuel Cell Forum, Göttingen, Germany, 2000. H.H. Gibbs, V.W. Va, and R.N. Giffin, US Patent 3,041,317, E.I. du Pont de Nemours and Company, June 26, 1962. D.J. Connolly and W.F. Gresham, US Patent 3,282,875, E.I. du Pont de Nemours and Company, November 1, 1966.

The use of Nafion® as electrolyte in fuel cells [13] [14] [15] [16] [17] [18] [19] [20] [21] [22] [23] [24] [25] [26] [27] [28] [29] [30] [31] [32] [33] [34] [35] [36] [37] [38] [39] [40] [41] [42] [43] [44] [45] [46]

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 21

Functional fluoropolymers for fuel cell membranes Renaud Souzy and Bruno Ameduri Laboratory of Macromolecular Chemistry, UMR (CNRS) 5076, Ecole Nationale Supérieure de Chimie de Montpellier, 8 Rue Ecole Normale, 34296 Montpellier Cedex 5 - France 1. INTRODUCTION Nowadays, the production of energy is mainly linked to nuclear, coal and combustion of fossils. However, these sources are not environment-friendly since they emit nuclear wastes or CO and CO2. Of the less polluted sources of energy arising, for instance, from wind, water, sun and others reported in this book, fuel cells are of growing interest. Fuel cells are original energy sources in which an electrochemical generator which directly converts the chemical energy of a fuel (hydrogen, methanol, ethanol, ethylene glycol, etc.) along with oxygen (from air, for example), into electricity, heat and water. Fuel cells [1–4] have already been involved in the production of stationary electrical energy and in the composition of materials for energy used in various fields such as transportation, space, telecommunications, portable electronic systems (portables, cellular phones), domotics (coproduction of electrical energy and heat, auxiliaries of power (APU) for automotives (board-computer, electrical commands, air conditioning) and computer security. These systems are required to show similar performances and comparable costs to vehicles using a thermal engine. Many investigations have been carried out on various characteristics of proton-exchange membrane fuel cell (PEMFC) in connection with: (1) the synthesis of membranes (the objective is to find original membranes with low cost and good performance and to increase their conductivity, thermal and mechanical stability low permeability to methanol, etc. (2) reduction of the overvoltage of electrochemical reactions involved such as the oxidation of hydrogen (or other fuels including methanol or ethanol) and the reduction of oxygen;

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(3) Optimisation of the transportation of reactants and the evacuation of heat at the membrane-electrode-assembly (MEA) system; (4) the electrodes and catalysts (or electrocatalysis); (5) diffusion layer and active layer; (6) process engineering; (7) tests in the heart of the cell; (8) bipolar plates; (9) stack. Many investigations have been undertaken on the number of patents and publications has greatly increased since 2000. The objective of this chapter concerns the synthesis and the properties of organic polymers involved in the expansion of the membrane, regarded as the ‘‘heart’’ of the fuel cell. Strict requirements are made of the membranes so that they must show a good thermostability (up to 80–100°C for 5000 h), high ionic conductivities mainly under high-humidity conditions, good mechanical strength, good chemical stability (especially the oxidative stability) and low methanol crossover. After a brief summary of the preparation of PEMFC from hydrogenated polymers, this chapter focusses on strategies of synthesising fluorinated polymers and the properties of the resulting membranes. 2. PEMFC BASED ON NONFLUORINATED POLYMERS Many hydrogenated polymers [5–26] have already been used in the preparation of PEMFC, and Table 1 provides a nonexhaustive list, and their performances in fuel cells have already been reviewed [1c,5–11]. These aromatic or heterocyclic polymers can be sulphonated polystyrenes (crosslinked or not) [12–14], sulphonated polyimides (PI) [15], sulphonated poly(aryl ether sulphones) [16,17], sulphonated poly(aryl ether ketones) [18], sulphonated phenol formol resins [19], sulphonated poly(phenylene oxides) [20,21], sulphonated poly(pphenoxybenzoyl-1,4-phenylenes) [22,23], phosphonic poly(phenylene oxides) [24], sulphonated silicates [25], sulphonated poly(benzimidazoles) [25] and sulphonated organic–inorganic hybrids [26]. Most of these nonfluorinated ionomer membranes, although of attractive price, are characterised by poor resistance to oxidation and thermal degradation. 3. SYNTHESIS OF FLUOROPOLYMERS FOR PEMFC In contrast to hydrogenated polymers, fluorinated polymers, regarded as high value-added materials, are potential candidates for PEMFCs due to their outstanding properties, which make them available for various applications [27–31].

Functional fluoropolymers for fuel cell membranes

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Table 1 Non-fluorinated polymers used in PEMFC Polymers

Structure

Sulphonated polystyrenes

( CH2

Reference

CH CH2

SO3H

SO3H

Sulphonated polyimides

SO3 HNEt3

[

O

O

O

-

Sulphonated poly(arylether sulphone)s

O3S HNEt3

O

N ] [R x

N O

[12–14]

) n

CH

N ]

O

O

[1c,16,17]

O

S

S q

O

y

O

SO3H O

[15]

O

O

N

O

O n

m

O

p

HO3S

Sulphonated poly(arylether ketone)s or SPEEK Sulphonated phenol formol resins

[1c,18]

O O

O

C m

n

HO3S OH

[19]

OH CH2

CH2 n SO3H

R

Sulfonated poly(phenylene oxide)s

[20,21] O n HO3S

Sulphonated poly(pphenoxybenzoyl1,4-phenylene)s

SO3H

O

O

O

O

n

Phosphonic poly(phenylene oxide)s

[22,23]

n'

[24]

R O n CH2P(O)(OH)2 R : CH3 or CH2P(O)(OH)2

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Table 1 (Continued) Polymers Sulphonated poly(benzimidazole)s

Structure

Reference [1c,25]

H N

N N

N

H

Sulphonated silicates

[26]

O O

n SO3H

Si O CH2

n

SO3H

The small size and the high electronegativity of the fluorine atom confers on the polymers a strong C–F bond and a low polarisability. Fluorinated polymers also exhibit high thermostability and chemical inertness, low refractive index and friction coefficient, good hydrophobicity and lipophobicity, valuable electrical properties, low relative permittivity and low surface energy. In addition, they are nonsticky and resistant to UV, ageing and concentrated mineral acids and alkalies. The high value of the unique characteristics of fluorinated polymers in the development of modern industries has ensured an increasing technological interest in them since the discovery of the first fluoropolymer, poly(chlorotrifluoroethylene), in 1934. Their fields of applications are numerous: they are used in paints and coatings [32] (for metals [33], wood [34], leather [35], stone [36], optical fibres [37], antifouling [38]), textile finishings [39], novel elastomers [31], high-performance resins, membranes [40], surfactants and fire-fighting agents [41], functional materials (for photoresists or microlithography [42], optical fibres, and conductive polymers [43]), biomaterials and thermostable polymers for aerospace [27–31]. Owing to the good properties of thermostability, chemical inertness and enhanced acidity of the sulphonic acid group in–CF2SO3H, various fluorinated polymers have already been involved as proton-exchange membranes for fuel cell applications. This chapter aims at describing different methods of synthesising fluoropolymers of use to PEMFCs and is divided into two major parts that report these syntheses and the properties of the resulting fluoropolymers for original membrane applications:the first one concerns the direct copolymerisation of functional fluorinated (especially acid function) monomers with commercially available

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monomers (and also fluoroalkenes); the second one deals with the chemical change of oligomers or polymers. This chemical modification can be carried out by the direct reaction of an oligomeric species onto a reactive group or by the irradiation of fluoropolymers, followed by grafting. 3.1. Fluorocopolymers of functional fluorinated monomers with fluoroalkenes from direct copolymerisation

These fluorinated (co)polymers bearing acidic side groups are particularly interesting materials for the preparation of membranes because of their efficient protonic conduction, their chemical and thermal stabilities and their resistance to ageing [30,40,44,45]. Three main groups can be distinguished, dependending on their acidic function: sulphonic, carboxylic or phosphonic. The synthesis of the corresponding monomers is discussed separately below, taking into account the aliphatic and then the aromatic monomers. 3.1.1. From fluorinated aliphatic monomers 3.1.1.1. Strategies of synthesis of functional fluorinated aliphatic monomers

The most pertinent studies were conducted by the Du Pont de Nemours company as early as 1962. They reported the synthesis of trifluorovinylsulphonyl fluoride, F2C  CFSO2F, by pyrolitic dehydrofluorination of 2,2,2,1-tetrafluoroethane sulphonyl fluoride [46,47]. Trifluorovinyloxy monomers of higher molar masses bearing a SO2F end group have been patented by Du Pont [46,47], Dow Chemical [48] and, more recently, by Solvay–Solexis [45]. These processes involve sulphone as a key reactant while the last step consists of a pyrolysis, as depicted hereafter. 3.1.1.1.1 Monomers bearing sulphonic acid function

This sultone allows the further introduction of a sulphonyl fluoride end group and an acid fluoride. The latter reacts to hexafluoropropylene oxide (HFPO) and, after pyrolysis, leads to the formation of sulphonyl fluoride perfluorovinyl ether, as follows [46,47]:

(A) Du Pont technology.

CF2

SO3

CF2

NR3 CF2

CF2

CF2

O

SO2

FOCCF2SO2F

CFCF3 O

FOCCF CF3 Δ

COF2

CF2

CF

(OCF2CF)

OCF2CF2SO2F

n

CF3 (OCF2CF) CF3

n

OCF2CF2SO2F (n = 0,1)

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This trifluorovinyloxy monomer easily reacts with tetrafluoroethylene (TFE) and leads to Nafion® membrane initially well used in the chlore/alkaly process, because of good stability to oxidation, to reduction and to corrosion. However, the high cost, medium thermal stability in fuel cell conditions (and hence poor performances under low humidification and at elevated temperatures (above 90°C) because of water loss) and the methanol crossover (when Nafion® is used in direct methanol fuel cell (DMFC) applications) are limitations, although recent investigations have improved these characteristics [49]. (B) Dow chemical technology. This company uses chloropentafluoropropylene oxide (CPFPO) instead of HFPO to promote an easier formation of the F2C  CFO end-group [48] as compared with the previous monomers. The presence of a shorter space between the trifluorovinyloxy and the sulphonyl end-group is observed to lead to F2C  CFCF2CF2SO2F. The radical copolymerisation of this fluorinated monomer bearing a sulphonyl fluoride function with TFE was successfully achieved and led to several patents [48], but neither industrial development nor commercialisation of these promising experimental membranes was achieved. (C) Solvay–Solexis technology. This company has recently revisited potential functional fluoropolymers for membranes and has proposed an alternative synthesis, especially of the above monomer. The process requires a hypofluorite, FOCF2CF2SO2F, which is added to 1,2-difluoro-1,2-dichloroethylene at low temperatures, the last step consisting of a simple dehalogenation [45]. (D) Asahi glass technology. The Asahi Glass Company develops the membranes in two ways. First, the electrochemical fluorination of a hydrogenated cyclic sulphone leads to FOC(CF2)2SO2F (in high yields), which reacts with HFPO and undergoes a pyrolysis to yield F2C  CFO(CF2)3SO2F [50]. The second procedure generates a fluorinated thioether as intermediate, which can then be oxidised into ClCO(CF2)2SO2Cl. Three additional steps are required to produce F2C  CFO(CF2)3SO2F. (E) Other ways [51–54]. Several synthetic methods have been achieved on a research scale. First, Krespan and England [51] synthesised perfluoroallyl fluorosulphate from hexafluoropropylene (HFP) and sulphur trioxide catalysed by BF3, and then isolated F2C  CFCF2OC2F4SO2F according to the following process: FSO2 CF2 COF+ KF

FSO2 CF2 CF2 OK

FSO2 C2 F4 OCF2 CF=CF 2

In addition, Kostov et al. [52] prepared FOCCF2SO2F (from the isomerisation of the tetrafluoroethane-β-sultone), which was reacted then with F2C  CFCF2OSO2F (prepared from the addition of HFP onto SO3, catalysed by B(OCH3)3 as the catalytic complex), leading to F2C  CFCF2OC2F4SO2F. Nguyen and Wakselman [53] used another method of synthesis and introduced

Functional fluoropolymers for fuel cell membranes

475

various HFPO units. More recently, DesMarteau [54] achieved the synthesis of original F2C  CFOCF2CF(CF3)OC2F4-G, where G represents SO2NHSO2CF3, N(Na)SO2CF3 or N(Na)SO2C4F8SO2N(Na)SO2CF3. 3.1.1.1.2. Carboxylic perfluoroalkoxyvinyl monomers [55–61] The most simple

key monomer F2C  CFCO2H can be prepared according to several routes [58,59]. In contrast to the preparation of sulphonated perfluorovinyl ethers, the synthesis of carboxylated perfluorovinyl ether is difficult. Less acidic perfluorovinyl ether monomers bearing a carboxylic acid have been commercialised by various companies: ● ● ●

F2C  CFOCF2CF(CF3)OCF2CF2CO2H from Du Pont [60]; F2C  CFOCF2CF2CO2H from Dow Chemical [61]; F2C  CF[OCF2CF(CF3)]nO(CF2)3CO2CH3 (n  0,1) CFO(CF2)3CO2CH3 from Asahi Glass Co.[68,69].

and

F2C 

3.1.1.1.3. Phosphonic acid perfluoroalkoxyvinyl monomers Besides most important carboxylic and sulphonic groups, novel perfluorovinyl ethers bearing an ω-phosphonic acid group were synthesised. The first one, F2C  CFO(CF2)3P(O)(OH)2, was achieved from the chemical change of carboxylic acid into phosphonic acid involving an iodinated intermediate ClCF2CFClOC3F6I [62–64]. Then, the synthesis of dimethyl perfluoro(3-vinyloxypropyl) phosphonate was achieved and co- or terpolymerised with TFE and perfluoro (propyl vinyl ether) [64]. Another alternative was proposed by Petersen et al. [65,66]; that is, the conversion of iodides into the corresponding phosphonites.Other exotic functional perfluorovinyl ethers were reported in a very interesting review published by Ukihashi and Yamabe [63]. Tatemoto and Nakamura achieved the preparation of F2CCFP(O)(OH)2 according to the following scheme [67]: ICl + F2C=CFCl

ClCF2CFClI

1) P(OMe)3 2) Zn

CF2=CFP(O)(OH)2

Most applications [68–70] from copolymers containing TFE and perfluorinated vinyl oxy-ω-sulphonyl fluoride or ω-carboxy (or derivatives) were searched from membranes for chlorine-alkali electrolysis, perfluorinated ion-exchange materials, hydrogen–oxygen solid polymer electrolyte fuel cells, or applications in space, ground (for vehicles) and undersea power sources. 3.1.2. Aliphatic fluorofunctional copolymers for fuel cell membranes

The example that has drawn the interest of many academics and industries concerns the copolymerisation of TFE with a perfluorovinyl ether with or

476

Renaud Souzy and Bruno Ameduri

without HFPO and bearing carboxylic or sulphonyl fluoride end group, is as follows: n F2C=CF2

+

m F2C=CFO RF

SO2F

(CF2 CF2)x CF2CF

y O(CF2CFO)n(CF2)p SO2F

CF3 x = 3.6 −13.5, n = 0, 1, 2, p = 1−5

The hydrolysis of the sulphonyl fluoride group produces the corresponding sulphonic acid derivative used for membranes in chlor-alkaly, ion-exchange membranes for electrolysis or for fuel cell applications. The most known copolymers commercially available since 1962 are Nafion® and Flemion® [1,40,71–80] (from Du Pont and Asahi Glass Co., respectively) when p  2 and n  1, although Tosflex® (n  0, 1 and p  1 or 5), for which the functional end group is an anion-exchange unit [71], is produced by Tosoh Co. Ltd. In addition, Dow® and Hyflon® Ion H (from Dow Chemical and Solvay–Solexis, respectively) when n  0 [45,72] and Aciplex® (from Asahi Chemicals Co.) when n  2 [1,45,63,73–80] have also led to potential membranes for these above applications [1,71–74] (Table 2). Three very interesting reviews on such copolymers for obtaining protonexchange membranes (PEMs) devoted to fuel cells have recently been published by Li et al. [4], Doyle and Rajendran [40] and by Arcella et al. [45]. In the current research on fuel cell membranes, efforts have been concentrated on PEMs with lower methanol permeabilities [81] and lower cost than Nafion® (US$ 780 m2) [5,82] and longer durability. However, it has been recently reported [83] that new solution-cast Nafion® membranes (NR-111) would allow the cost to drop to US$ 50 m2 with a volume production of ⬃ 2 million m2 year1. For similar strategic fuel cell applications, other monomers with longer chain lengths [84–86] were also used in the copolymerisations with TFE, and original studies have dealt with the use of phosphonate or phosphonic acid end groups [87–89] from fluoromonomers such as F2C  CFOC3F6P(O)(OH)2. In addition, functional perfluoroalkoxyalkyl vinyl ethers, mainly those containing a sulphonyl fluoride or their derivatives, were also successfully copolymerised with vinylidene fluoride (VDF) in various processes [47,85,90–95]. Other strategies starting from the radical copolymerisation of fluoroalkenes with sulphonic [96,97] or carboxylic acid [98] were also reported. 3.1.3. PEMFC based on aromatic fluorinated polymers

This subsection reports the preparation and the characterisation of PEMFC based on aromatic (per)fluorinated polymers. In the past few decades, attention has been focused on the preparation of new fluorinated monomers and aromatic fluoropolymers. This topic was

Functional fluoropolymers for fuel cell membranes

477

Table 2 Fuel cell membranes arising from copolymers of tetrafluoroethylene (TFE) and perfluorovinyl ether alkyl sulphonylfluoride Structural parameters (and monomer contents)

Supplier and trademark

Equivalent weight (IEC, meq g1)

Thickness (m)

Nafion® 120

1200 (0.83)

250

Nafion® 117

1100 (0.91)

175

Nafion® 115

1100 (0.91)

125

Nafion® 112

1100 (0.91)

50

Flemion® T

1000 (1.00)

120

Flemion® S

1000 (1.00)

80

Flemion® R

1000 (1.00)

50

1000–1200 (0.83–1.00)

25–100

800 (1.25)

125

n  1, x  5–13.5, p  2

n  0 –1, p  1–5

n  0, p  2–5, x  1.5–14

DuPont

Asahi Glass

Asahi Chemicals Aciplex® S

n  0, p  2, x  3.6–10

Dow Chemical Dow® Solvay Hyflon® Ion

900 (1.11)

recently reviewed [99] because of the characteristic effects of the aromatic group on physicochemical properties (e.g. Tg and the thermostability of the obtained polymers). To the best of our knowledge, it can be observed that aromatic fluorinated macromolecules for PEMFCs by direct radical (co)polymerisation can be prepared from two groups of functionalised aromatic perfluorinated monomers: (i) α,β,β-trifluorostyrene (TFS) and (ii) [(α,β,β-trifluorovinyl)oxy] benzene (TFVOB) as sketched below: F

F

F

Q

G

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Renaud Souzy and Bruno Ameduri

In TFS, Q stands for “-“; in TFVOB, Q is equivalent to O. G stands for sulphonic or phosphonic acid function. The first part covers the PEMFC obtained by direct (co)polymerisation of functionalised TFS (3.1.3.1). Various routes dealing with the synthesis and characterisation of (co)polymers containing functionalised TFVOB have been investigated and are presented in the second part (3.1.3.2). 3.1.3.1. PEMFC based on functionalised TFS 3.1.3.1.1. Synthesis and polymerisation of TFS The synthesis of TFS and its (co)polymerisation with different comonomers have been reported by various authors [100–105]. Quasi-exhaustive ways of synthesising TFS have been reviewed by Heinze and Burton [106]. TFS was prepared after a coupling reaction of perfluoroalkenylzinc reagents [F2C  CFZnX, (Z) F3C-CF  CF-ZnX, (E) F3C-CF  CF-ZnX with X equivalent of bromide or iodide] with aryl iodides in the presence of Pd(PPh3)4 as catalyst to yield the corresponding fluoroalkenes (Scheme 1). The synthesis of p-sulphonic acid-TFS was patented by Ballard Power System Company [107] (Scheme 2), inspired by Polish researchers [108]. The corresponding sulphonic acid monomer was obtained by hydrolysis of the phalogenosulphonate-TFS. In addition, materials can be prepared by a cyclodimerisation reaction of TFS [109–111].

F2C=CF

ZnBr

Pd(PPh3)4

I R

F2C=CF

1-10H ours 60-80˚ C

R

Scheme 1: Synthesis of TFS by Heinze and Burton [106].

I

ClSO3H + HCl

I

SO 2Cl

KF

I

yield : 90 %

yield : 100 %

F2C=CFZnBr

F2C=CFZnBr

F2C=CF

SO2Cl yield : 69 %

Scheme 2: Synthesis of 4-fluorosulphonate–TFS [107].

SO 2F

F2C=CF

SO2F yield : 69 %

Functional fluoropolymers for fuel cell membranes

479

The bulk polymerisation of TFS was achieved for the first time by Prober [101] in 1953. Then, in 1981, Tevlina et al. [112] copolymerised TFS (I) with vinyl fluoromonomers such as N-vinylpyrrolidone (II), H2C  CF–CN (III), FHC  CF–COOMe (IV) and F2C  C(CF3)COOMe in the presence of azobisisobutyronitride (AIBN). 3.1.3.1.2. Poly TFS incorporated in PEMFC p-Chloro or fluorosulphonate TFS synthesised by Stone et al. [107] and patented by Ballard Power System was copolymerised in emulsion (in the presence of dodecylamine hydrochloride) with TFS functionalised or not functionalised (Scheme 3). In 1999, Stone et al. [113] proposed a PEMFC based on phosphonic acid TFS. Polymers were prepared from two basic steps (Scheme 4): (i) Synthesis of 4-iodo-benzene phosphonic acid dimethyl ester (4-1), and (ii) synthesis of the pdimethyl phosphonate-TFS (4-2). This monomer either homopolymerised or copolymerised (Scheme 4). Although it is known that TFS does not homopolymerise under radical conditions, these authors claimed that a homopolymer of a monomer (4-2) was first prepared by them. The best yields were obtained in bulk polymerisation initiated by AIBN. The prepared membranes were characterised by a low intrinsic viscosity and very poor mechanical properties. Nevertheless, the homopolymer (43) was hydrolysed to afford an ionomer mixture, which was soluble in an aqueous base. As a consequence, the physical properties of ionomer did not fulfill the requirements for using these polymers as a PEMFC. In a second method, Stone et al. [113] copolymerised monomer (4-2) with TFS (Scheme 4) by emulsion polymerisation in 21% isolated yield. The optimised ratio between TFS and dimethylphosphonate–substituted-TFS monomer in the copolymer (4-5) was 2.4:1. The molecular weights of the resulting copolymer were Mn  38,100 and M w  105,900 g mol1. Furthermore, homopolymer (4-3) (membrane A) was hydrolysed under acidic conditions (hydrochloric acid in dioxane, 100°C, 20 h). The yield and the equivalent weight of acid functions were 95% and 130 g mol1, respectively.

( CF2

CF ) m

( CF2

SO2F

CF ) n

( CF2

CF ) p

A2

A1

( CF2

CF ) n

A3

m, n, p, q > 0 A1, A2, A3 : CF=CF 2, CN, NO2, OH, OR, SO3H, PO2H2, PO3H2, COOH, OSO3H, OPO2H2, OPO3H2, +NR3, CH2NR3+

Scheme 3: Copolymerization of fluorosulphonate-TFS [107].

480

Renaud Souzy and Bruno Ameduri

O 3 PCl3

I

AlCl 3

I

I

PCl 2

70 ˚C

Br2

I

CH3OH

P (4-1)

OCH3 OCH3 Yield = 75 %

F2C=CF ZnBr (27-2) Pd(PPh3)4 Bulk Polymerization CF2 CF

O P OCH3 OCH3

F2C=CF

AIBN

n

(4-2)

Yield = 50 % Yield =34 % O

water K2S2O8

P OCH3 OCH3

CF=CF2

(4-3)

Hydrolysis (CF2 CF) (CF2 CF) n m

(CF2 CF) (CF2 CF) m n

(4-5)

O

P OH OH

(4-4)

O

P OCH 3 OCH3

Scheme 4: Synthesis and homopolymerization of dimethylphosphonate-4-substituted-TFS according to Stone et al. [113].

Copolymer (4-4) [113] was hydrolysed by the authors using two processes: (i) basic conditions (potassium hydroxide, 84°C, 64 h), membrane C1 and (ii) acid conditions with a dimethylformamide (DMF) pretreatment, membrane C3. The electrochemical properties of these membranes are presented in Table 3 and compared with other copolymers bearing phosphonic acid groups [66,114]. Finally, the authors concluded that the best results were obtained with an acidic hydrolysis and they explained that the membranes based on a sulphonic acid–TFS gave better results than those obtained from a phosphonic acid homologue. 3.1.3.2. PEMFC based on functionalised TFVOB In this subsection, two main kinds of aromatic PEMFCs incorporating functionalised TFVOB are presented: (i) polymers prepared by thermocyclodimerisation, and (ii) macromolecules obtained by direct (co)polymerisation of TFVOB with commercially available fluoroalkenes [115]. One of the most interesting properties of perfluoroalkyl TFVOB is that it undergoes thermal cyclopolymerisation [2π  2π] with temperature [116–125] (up to 150°C) (Scheme 5). The formed perfluoroalkylpolymer is a thermoplastic and thermoset perfluorocyclobutane (PFCB) [116(c),116(g),126,127].

Functional fluoropolymers for fuel cell membranes

481

Table 3 Electrochemical characteristics of different ionomer membranes achieved by Stone et al. [113], Xu and Cabasso [114], and Kotov et al. [66], compared with those of Nafion® Ionomer

EW (g mol1)

Effective EW (g mol1)

A (membrane based on homopolymer 27-4)

130

215

–

–

[113]

C1 (membrane based on copolymer 28-2)

350

670

0.01–0.1

15

[113]

C3 (membrane based) on copolymer 28-2

200

380

–

77

[113]

D (membrane based on 3.86 : 1 ratio n : m [CF2-CF2]n[CF2CFX]m with XO (CF2)3P(O)(OH)2)

357

357

76

22

[66]

E (membrane based on 4.63:1 ratio n:m [CF2-CF2]n [CF2-CFX]m with XO (CF2)3P(O) (OH)2)

370

370

69

21

[66]

F (membrane based on poly (dimethylphenylene oxide phosphonic acid))

111

202

–

–

[114]

G (membrane based on poly(dime thylphenylene oxide diphosphonic acid))

70

107

–

–

[114]

Nafion®

Transverse Water proton absorption at conductivity 100°C (%) (mS cm1)

1100–1500 1100–1500

50–110

Reference

[40,49,73,82]

EW, equivalent weight.

F

ArO

F

F

F

F

F

OAr

F F >150 ˚C

Ar

O

F

..

F

F F

F F O Ar

Scheme 5: Formation of aromatic PFCB [116(c),116(g),126,127].

Ar

O

F

F F F O Ar

482

Renaud Souzy and Bruno Ameduri

As a matter of fact, Babb et al. [127(c),128(a)] developed a series of TFVOBs prepared from bis- and trisphenols, such as tris(hydroxyphenyl)ethane. These different perfluorinated aryl ethers were thermocyclodimerised and led to thermoset polymers (Tg  180°C) with good thermal stability (they are stable up to 434°C), thermal/oxidative stability and mechanical properties [123(b),127(d)]. Furthermore, these authors [116(c)] prepared PFCB aromatic polyethers containing a siloxane group. Their syntheses involved an aryl Grignard reagent from 4[(trifluorovinyl)oxy]bromobenzene that led to a high-yield (87%) synthesis of 4-[(trifluorovinyl)oxy]phenyldimethylsilane. The latter was finally hydrolysed in situ and then condensed to yield bis[1,3-[4-[(trifluorovinyl)oxy]phenyl]]-1,1,3,3tetramethyldisiloxane in 43% yield. Such a monomer was thermocyclodimerised (at 210°C) to yield original siloxane PFCB. In 2000, Smith et al. [123(b)] reported the synthesis of different PFCB polyarylethers (Scheme 6). Interestingly, the reactive Grignard [131] or lithium [128(b),132,133] compound of 4-[(trifluorovinyl)oxy]bromobenzene [116(c)] gave rise to an increasing number of organic/inorganic fluorinated compounds [116(c),128(b),132–134]. The current intensification of interest in the preparation of PEMFCs based on electrolyte polymers has prompted us to synthesise aromatic monomers such as trifluorovinyl ethers functionalised by acid groups. In particular, we reported the preparation of 4-TFVOB phosphonic acid [129] (Scheme 7). According to various methods of phosphonation such as Michaelis–Arbuzov or Michaelis–Becker reactions, it was shown that the best

HO

Ar

OH

1) KOH / BrCF2CF2Br

F

F

F

F

F

O Ar

O

F

F F

2) Zn / CH3CN

F O

F F F O Ar n

CH3 Ar :

OCF=CF 2 F 3C

CF3

H3C

Scheme 6: Poly aryl vinyl ether synthesized by Smith et al. [123(b)].

HO

NiCl2

Br 1)KOH / Br C2F4Br

HP(O)(OEt)2

2) Zn / CH3CN

NaH

F2C=CFO

F2C=CFO HP(O)(OEt)2

(8-1)

Pd(PPh3)4

tBuLi Et2O -80˚C F2C=CFO

Ö

Br

Li

P OR OR (8-2)

BrSiMe3

Ö F2C=CFO

(R: Et or Me)

ClP(O)(OR)2

P OH OH (8-3)

Functional fluoropolymers for fuel cell membranes

P(OR)3

Scheme7: Synthesis of new aromatic perfluorovinyl ether monomers containing phosphonic acid functionality [129].

483

484

Renaud Souzy and Bruno Ameduri

yield was achieved when the reaction involved a palladium triphenyl phosphine complex as the catalyst. In 2000, Ford et al. [133] reported the synthesis of aromatic perfluorovinyl ether monomers containing the sulphonamide and the sulphonic acid functionality for different applications such as the preparation of PEMFC [135]. As in the previous example, Ford et al. used a [p-((trifluorovinyl)oxy)phenyl]lithium, which was cross-coupled with FSO2Cl to give 4-[(trifluorovinyl)oxy]benzenesulphonyl chloride in 65% yield ((9-1), Scheme 8). Monomers (9-2), (9-3) and (9-4) were prepared in 91, 85 and 80% yield, respectively. The thermal behaviour of monomer (9-3) was studied by differential scanning calorimetry (DSC). The exothermic polymerisation started from 175°C (Tonset) and 214°C (Tmax). These different monomers having a structure similar to that of monomer (9-1) (Scheme 9) and polymers prepared by thermocyclodimerisation of these monomers were patented by 3M Innovative Properties Company [117(d)] in 2001 (Scheme 9) for PEMFC applications. 3.1.3.3. PEMFC synthesised from direct radical terpolymerisation of α,β,β-Trifluorovinyl benzyl ethers More recently, Souzy et al. [136] studied the radical homo-, co-,

OCF=CF2

OCF=CF2 t-BuLi

OCF=CF2 NH3

FSO2Cl

Li

Br

SO 2NH2

SO 2Cl (9-1)

OCF=CF2 OCF=CF2

OCF=CF2

OCF=CF2

(9-2)

SO2Cl

1-NR3, FSO2(CF2)4SO2F + 2- H O O

S

N H

S

OCF=CF2

OCF=CF 2

OO

(9-3)

SO 2NHSO2(CF2)4SO 2NHSO 2 (9-4)

Scheme 8: Synthesis of aromatic perfluorovinyl ether monomers containing the sulphonamide and the sulphonic acid functionality according to Ford et al. [133].

Functional fluoropolymers for fuel cell membranes

485

and terpolymerisation of 4-[(α,β,β-trifluorovinyl)oxy] bromobenzene with commercially available fluoroalkenes such as VDF and CTFE, HFP, and perfluorovinyl methyl ether (PMVE) (Scheme 10). The authors optimised the conditions of co- and terpolymerisation and even tetrapolymerisation in terms of the nature of the radical initiators, the nature of solvents (fluorinated or nonhalogenated)

XO2S

OCF=CF2 F2C=CFO

F2C=CFO

OCF=CF2

OCF=CF2

SO2X

SO2X

XO2S

F2C=CFO

F2C=CFO

OCF=CF2 OCF=CF2

XO2S

SO2X

Scheme 9: Functionalised polyarylene vinyl ethers according to 3M Innovative Properties company [117(d)].

n H2C=CF2

+

m F2C=CFZ

+

p F2C=CF O

HFP : Z=CF3 PMVE : Z=OCF3 CTFE : Z=Cl Br radical

(CH2CF2)( CF2 CF) t

Z

u

CF2 CF O

v

Br

Scheme 10: Terpolymerisation of 4-[(α,β,β-trifluorovinyl)oxy]bromobenzene with Fluoroalkenes (VDF, HFP, PMVE and CTFE) [136].

486

Renaud Souzy and Bruno Ameduri

and the relative reactivities of the comonomers. Furthermore, they synthesised terpolymers incorporating 4-[TFVOB] bromobenzene and showed that the VDF/PMVE mixture enhanced the incorporation of the α,β,β-trifluorovinyl benzyl ethers in the terpolymers. From these optimal experimental conditions [115,137], these authors have reported the synthesis of new polymer electrolyte membranes based on fluoropolymers incorporating 4-TFVOB sulphonic acid (TFVOBSA), synthesised in a 72% overall yield. Further, the radical (co) and terpolymerisation of this monomer or of the sulphonyl chloride precursor (TFVOBSC) with VDF, HFP, and PMVE (Scheme 11) were also investigated. The incorporation of the aromatic functional monomer was yet limited to 9% and it was noted that the higher its feed content, the lower the yield. Poly(VDF-ter-HFP-ter-TFVOBSC) and poly(VDFter-PMVE-ter-TFVOBSC) were hydrolysed by a slight alkali treatment without affecting the content of fluoromonomers in the terpolymers and without any dehydrofluorination of VDF units [137]. The membranes, formulated from a commercially available fluoropolymer, led to good film-forming properties. However, their electrochemical properties were disappointing (the conductivity values ranged from 0.01 to 0.1 mS cm1, and were far from those of Nafion® (Table 4)). 3.1.3.4. Conclusion Functional TFVOBs are interesting precursors that can undergo

thermal cyclopolymerisation yielding low-molecular-weight perfluoroalkylpolymers and provide high-Tg thermoset polymers with good thermal stability. These perfluorinated aryl ethers are currently used for the preparation of ion-exchange resins and ionomer membranes. Recently, investigations dealing with radical terpolymerisation of aromatic fluoromonomers functionalised by sulphonic acid groups [137] or phosphonic

H

F

H

F VDF

F

F

F2C=CFO F R R: CF3( HFP) or OCF3(PMVE) (I)

[

CH2 CF2

x

SO 2Cl

Rad.

CF2 CF R

y

CF2 CF O

] z w

SO2Cl

Hydrolysis (II)

[

CH2 CF2

x

CF2 CF R

y

CF2 CF O

]

z w

SO3H

Scheme 11: Radical terpolymerization (I) of TFVOBSC with VDF or HFP or PMVE, and hydrolysis (II) of the corresponding terpolymers [115,137].

Functional fluoropolymers for fuel cell membranes

487

acid [129] with VDF, CTFE, HFP and PMVE led to fluoropolymers bearing acid functionality with low yields and low conductivity values. 3.2. Fluorinated graft copolymers for PEMFC by chemical modifications 3.2.1. Introduction

Various fluorinated graft copolymers have been synthesised by the chemical modification of hydrogenated or fluorinated copolymers [30(b)]. This subsection gives nonexhaustive examples of grafted copolymers involved in fuel cell applications from either hydrogenated (grafted with fluorofunctional synthons) or fluorinated polymers (grafted with hydrogenated functional reactants). 3.2.2. Chemical modifications of polyparaphenylenes

An original approach was suggested by Le Ninivin [22], who synthesised novel polyparaphenylenes (PPPs) bearing 2-sulphonic acid-1,1,2,2-tetrafluoroethyl side groups. They were obtained by nucleophilic substitution of p(1-sulphonic acid-1,1,2,2-tetrafluoro)phenate onto poly(p-fluorobenzoyl-1,4-phenylene) as follows: F

HO3S

O n

CF2CF2O

KO 3SCF2CF2O K2CO3 , DMAc, 145˚C

F

O

OH

O m

O p

These amorphous polymers (Tg  156°C) exhibited moderate to high molecular weights (with PS standards): M n 26,000 and M w 65,000 g/mol. These fluorinated PPPs led to the formation of original membranes endowed with a good thermal stability (these films were stable in air up to 310°C), interesting electrochemical properties, very low methanol crossover (its intrinsic permeability to methanol was lower than that of Nafion® for the same thickness), high-ionic exchange capacity (IEC) (1.3 meq H g1) and satisfactory conductivity (8.5 mS cm1) owing to the electron-withdrawing C2F4 adjacent group that enabled the sulphonic acid to show an enhanced acidic character for a thickness of 40 μm. All these relevant characteristics show that such original membranes are potential candidates for direct methanol fuel cells [23]. 3.2.3. Chemical modification of fluoropolymers from irradiation followed by chemical grafting

The synthesis of original fluorinated graft copolymers achieved from the introduction of grafts brought by macroinitiators has recently been reported [30(b)]. They can be obtained via two different routes: (i) from

3.2.3.1. Introduction

488

Table 4 Ion-exchange capacities and proton conductivities of different PEMFCs Membrane

(CF2

CF)n

Proton conductivity (mS cm1)

Reference

175

0.67–1.25

50–110

[46,47]

75

1.1–2.6

2–10

[107]

90

1.5

0.10

[113]

110

1.3

2.0

[22,23]

(CF2 CF)n'

R

(CF2

IEC (mEq g1)

SO3H

CF)n

BAM3G

(CF2 CF)n'

R

P(O)(OH)2

SO3H

O

O

O

n

O

n'

Sulphonated poly(p-phenoxybenzoyl-1,4-phenylene) (sPPBP)

Renaud Souzy and Bruno Ameduri

[-(CF2CF2)x-[CF2-CF-[OCF2CF(CF3)]pOCF2CF2SO2F]]n (p: 0 or 1) Nafion® 117

Film thickness (μm)

OCF2CF2SO3H

F

O

O

40

1.3

8.5

71

0.4

0.037

95

0.5

0.059

51

0.6

0.082

[23]

O

n

n'

[

CH2 CF2

x

CF2 CF R

y

CF2 CF O

]

z w

R : CF3 (HFP) or OCF3 (PMVE)

SO3H

Blend of commercially VDF/HFP (80/20 wt%) with poly(VDFter-HFP or PMVE-ter-[trifluorovinyloxy]benzene sulphonic acid)

[115,137]

Functional fluoropolymers for fuel cell membranes

Sulphonated poly(p-fluorobenzoyl-1,4-phenylene)-co-poly(pethoxytetrafluoro-p-phenoxybenzoyl-1,4-phenylene) (GPS)

489

490

Renaud Souzy and Bruno Ameduri

the copolymerisation of fluoromonomers, one of them containing an initiating species that does not participate in the (co)polymerisation; and (ii) from the activation of the fluoropolymers under thermal initiation, ozone, plasma, swift heavy ions, X-rays or electron beam [138–145], named the ‘‘grafting from’’ technique. One example arising from the thermal activation of hydrogenated polymer was proposed by the Du Pont Company [146], which grafted H2C  CHC2F4OC2F4SO2F onto polyethylene. The copolymer had a Tm  115°C and a 10% weight. loss by thermal gravimetric analysis (TGA) at 380°C under nitrogen, and could be used in fuel cell applications. Regarding the synthesis of graft copolymers from the activation process, three different methods may be used [138,144] : (a) If the (pre)irradiation is carried out in air, radicals react with oxygen leading to the formation of peroxides and hydroperoxides (hydroperoxide method). When in contact with monomer, the irradiated polymer initiates grafting by thermal decomposition of hydroperoxides. (b) In the absence of air, these macromolecular radicals remain trapped in the polymer matrix and initiate the grafting in the presence of a monomer (trapped radicals method). (c) Simultaneous radiation grafting is, therefore, a single-step process whereas the preirradiation method involves a two-step process [138,144,145]. These three ways are briefly sketched as follows:

radiation vacuum polymer

trapped radicals M simultaneous irradiation and grafting of M Δ

radiation air O

O

O

OH

O

OH

Z M

Z

M

Z

M n

n

M n

with Z : " " or O

macroinitiator

In each part, the fluorinated comb-like copolymers can be prepared from different kinds of F-homopolymers (e.g. PTFE, PVDF and PCTFE) and also those arising from different families of F-copolymers such as poly(TFE-co-HFP) (or Poly(tetrafluoroethylene-co-hexafluoropropene)copolymer FEP), poly(ethylene-alt-tetrafluoroethylene)copolymer(ETFE), and poly(TFE-co-PAVE) (or poly(terafluoroethylene-co-perfluoropropyl vinyl ether)copolymer (PFA). Grafting the required monomer to introduce the function properly has also been extensively investigated by many authors, offering a wide range of

Functional fluoropolymers for fuel cell membranes

491

well-architectured copolymers involved in many applications such as biomaterials (e.g. artificial hearts, cardiovascular prostheses), compatibilising agents, protection of substrates (e.g. metals), pH-sensitive membranes, membranes for purification of water and fuel cells. 3.2.3.2. Synthesis of graft copolymers from the ozone activation of the polymer

Ozone, commonly written O3, is a cheap gas and is nowadays well known as a topic of environmental concern. The ozonisation (or ozonation) of polymers has been investigated by many authors and was recently reviewed [147]. Polyvinylidene fluoride (PVDF) containing peroxide initiated the polymerisation of PEOMA at 100°C in NMP from [PEOMA]/[PVDF] weight ratio ranging between 1:1 and 6:1 [148]. Various electrochemical properties of these membranes were studied, such as (i) a very high liquid electrolyte uptake capacity (these membranes could absorb 74 wt% of liquid electrolyte – higher than the best ones that absorb 65% in the presence of inorganic fillers [149]; (ii) an ionic conductivity of 1.6 mS cm1 at 30°C; (iii) a satisfactory transference number of 0.15, characteristic of that in polymer electrolytes or concentrated solutions [150]; and (iv) electrochemical stability enabling these membranes to find applications in Lithium-ion rechargeable batteries. 3.2.3.3. Activation by electron beam and γ-ray 3.2.3.3.1. Introduction Radiation-induced grafting (from

60

Co) can also be used for the synthesis of original copolymers [138,143–145]. Different types of high-energy radiation are available to be used for the grafting process [138–141,144] although cross-linking may occur, as remarkably reported by Forsythe and Hill [145(a)]. 3.2.3.3.2. Activation of fluorinated homopolymers 3.2.3.3.2.1. Synthesis of PVDF-g-PM graft copolymers from the irradiation of PVDF followed by grafting Machi et al. [151] synthesised PVDF-g-PAA graft

copolymers after irradiation of PVDF films by 20 Mrad electron beam followed by immersion in acrylic acid (AAc). These resulting membranes exhibit an electric resistance of 6.2 Ω cm2, compared with, e.g. a film that does not contain any AAc (10–15 Ω cm2). However, the most pertinent surveys of the activation of PVDF were conducted by a Scandinavian team, a Swiss Institute and an English Laboratory. These Finnish authors [161–163,181–187] synthesised and then fully characterised original sulphonated PVDF-g-PS (polystyrene) (or PVDF-g-PSSA) poly(styrene sulphonic acid) copolymers in a threestep procedure: these authors irradiated porous films of PVDF with electron beams at various doses (25–200 kGy), which was followed by the grafting of

(A) Results from Sundholm’s team.

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styrene, and, in a final step, the sulphonation of the aromatic ring was achieved in the presence of chlorosulphonic acid, summarised as follows PVDF

γ -rays 25−200 kGy

(CH2CF2)x

CH

CF2

1) H2C=CH 2) ClSO3H

(CH2CF2)x

CHCF2 O

O

O

H

Act-PVDF

(VDF)y

O

Act-PVDF

R

CHCF2 (VDF) y

CHCF2

Z

Z

CH2 CH p

CH2 SO3H

CH q

SO3H

PVDF-g-PSSA with Z =O or " " .

Surprisingly, the authors claimed that no homopolystyrene was formed. The degree of sulphonation was 11–71% [161]. However, to achieve 95–100% sulphonation, a double ClSO3H concentration was required for a 2-h reaction time. Under these conditions, the authors did not observe any trace of chlorine atom in the PVDF-g-PSSA membrane. These authors observed a high degree of grafting (50–86%) and they also noted that the grafts were formed from both C–H and C–F branch sites of PVDF. Nevertheless, the presence or absence of a C–O bond (that could arise from the formation of peroxides) was not mentioned. Relevant evidence of the structure [161,162], thermal behaviour [161–163,181] and conductivities (up to 120 mS cm1 at room temperature) [161–163,181,182] slightly higher than those of Nafion® were provided. Surprisingly, these authors claimed that the PVDF-g-PSSA membranes were stable up to 370°C under air atmosphere, and upto 270°C in a highly oxidising atmosphere [162,181] (from 340°C, PS grafts started to decompose [181]). These Finnish authors assumed that the grafting which occured in the para position of the phenyl ring, took place in the amorphous regions of PVDF [162]. The original films were characterised by Raman [161] and NMR [183] spectroscopies, wide-angle X-ray scattering (WAXS), and small-angle X-ray scattering (SAXS) [161,162,184]. Swelling tests in various solvents and in water and tests of conductivity (by impedance spectroscopy) indicated that these films were potential PEMs for fuel cell applications [163,185]. Indeed, the higher the content of sulphonic acid functions, the higher the conductivity, and these values were enhanced when the crystallinity rate decreased [162].

Functional fluoropolymers for fuel cell membranes

493

By confocal Raman spectroscopy, this Finnish team [187] characterised the fuel cell-tested PVDF-g-PSSA membranes and also noted that the cross-linked membranes in the presence of divinylbenzene did not undergo any degradation as observed on noncross-linking ones. In addition, this same group performed the synthesis of controlled PVDFg-[PVBC-g-PSSA] graft copolymers for designing proton-exchange membranes for fuel cell applications (where VBC stands for vinyl benzyl chloride) according to the same strategy [188]. These PVDF-g-PVBC copolymers act as suitable macroinitiators via their chloromethyl side groups in the atom transfer radical polymerisation [189] of styrene, as follows: PVDF

1) γ rays 2) VBC

PVDF Z

Z

CH2

CH2

CH n

CH2Cl

CH2Cl

CH m

PVDF-g-PVBC with Z=O or " PVDF

[

CuBr / bipy 100-130˚C

[

PVDF-g-PVBC

"

CH2

CH2 CH2

(CH2CH)x

Cl

CH m ]

]

CH n

CH2(CH2 CH)y

Cl

PVDF-g-[PVBC-g-PS]

The highest conductivity measured for these membranes was 70 mS cm1, which is assumed to be similar to that of Nafion®. Scanning electron microscopy/ energy-dispersive X-ray results showed that the membranes had to be grafted through the matrix with both PVBC and PS to become proton conducting after sulphonation [188]. (B) Surveys carried out by Scherer. Scherer et al. [190,191] investigated the synthesis of PVDF-g-PS graft copolymers, and then sulphonated into PVDF-gPSSA as a useful ion-exchange membrane. More surveys by these authors were devoted to the synthesis, properties and applications of ETFE-g-PM and FEP-gPM graft copolymers reported later.

Grafting acrylate (Acr) onto PVDF for original lithium batteries, was also accomplished by Kronfli’s team [192] the PVDF-g-P(Acr) graft copolymer leading to an improvement in the adhesion of composite electrode to current collectors and to an increase in electrolyte solvent uptake. Graphite–LiCoO2 cells containing such modified PVDF-g-PM showed good rate performance and stable cycle life. An improved process of preparing (C) Preparation of PVDF-g-P(Acr) copolymers.

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microporous membranes for solid electrolytes for Li cells by casting the graft copolymer solution was also achieved by this group. By the same strategy, Flint and Slade [164] synthesised PVDF-g-PSSA, and the physical and electrochemical properties of the resulting membranes were investigated. These membranes were used in MEA for fuel cell applications. The conductivity of the membrane containing 30% graft was 30 mS cm1 with an overall resistance of 6.5 Ω cm2. Interestingly, the authors compared the physical and electrochemical properties of their PVDF-g-PSSA with those of FEP-gPSSA and Nafion®. (D) Synthesis and applications of PVDF-g-PVBC. By grafting VBC onto two grades of PVDF activated by γ -rays 60Co, Danks et al. [165] synthesised PVDF-g-PVBC graft copolymers. Then, the chloromethyl endgroups of the grafts underwent an amination reaction into benzyl trimethylammonium (BTMA) hydroxide [165] or BTMA chloride (BTMAc) [166] for anion-exchange membrane applications: CH2CF2

CHCF2 Z CH2 CH n

+ CH2N(CH3)3OH

with Z : O or " "

Indeed, these quaternary ammonium side groups allowed these novel membranes to be used as alkaline fuel cells for portable applications owing to the conductivity of hydroxide ions. Unfortunately, this last step led to the formation of very brittle PVDF-g-PBTMAc graft copolymers because of the degradation of the polymer backbone by the expected dehydrofluorination of PVDF, making them unsuitable for use as membranes for fuel cells or electrochemical devices. Hence, these British authors decided to use the same strategy with a perfluorinated copolymer, a poly(TFE-co-HFP) or FEP copolymer that is not base-sensitive so that they obtained a membrane endowed with an IEC of 1.0 meq g1. 3.2.3.3.2.2. Synthesis, properties, and applications of PTFE-g-PM graft copolymers Radiolysis of PTFE has been extensively studied in various books and

reviews [138–141,193,194] (irradiated PTFE can still be used for grafting polymerisation several years after it has been exposed [70,193]). Many monomers have been grafted onto irradiated PTFE ranging from acrylic acid [195], methyl trifluoroacrylate (MTFA) [196], styrene [151–156], vinyl pyrridine [157,158,197], and, to a lesser extent, N-vinyl pyrrolidone (NVP) [198]. Usually, decreases in the crystallinity content and melting point of PTFE are noticed. Kostov et al. [157,158] extensively developed researches on the grafting of fluoropolymers (mainly PTFE) and reproduced Chapiro’s results [138]. This

Functional fluoropolymers for fuel cell membranes

495

Bulgarian team observed that the resulting graft copolymers were used not only for ionomers and enzyme biosensors [195(b)], but also for IEMs. Interesting surveys were also carried out by Nasef et al. [152–155] on the synthesis of PTFE-g-PSSA copolymers for PEMs, for which striking structural investigations and thermal [152], chemical and mechanical stabilities [152–155] have been carried out. Kinetics approaches were also attempted [158,195(b)]. More recently, Liang et al. [152] used a similar strategy (60Co beam, under 20 kGy at 110 Gy min1) for preparing PTFE-g-PSSA graft copolymers for fuel cell membranes. This Chinese group carried out the grafting directly on films having Tg of 117°C and a thicknesses of 70 μm, with degrees of grafting ranging between 0.9 and 31.2%. They obtained grafted copolymers with Tg, thickness and resistances in the range 117.5–125°C, 73–140 μm, and 11.9–48.6 Ω cm2, respectively. However, these membranes undergo oxidative degradation above 60°C. These three teams studied the effect of grafting and the experimental conditions on the degree of grafting (d.o.g), the kinetics of grafting and the properties of the grafted films (IEC, water uptake, dimension charge percentage and the specific resistance). 3.2.3.3.2.3. Activation of fluorinated copolymers followed by the grafting of various monomers Various TFE containing grafted copolymers have been involved

in fuel cell applications. This subsection considers three main examples: poly(Eco-TFE), poly (TFE-co-HFP) and poly(TFE-co-PPVE) copolymers, where E, HFP and PPVE stand for ethylene, hexafluoropropylene and perfluoropropyl vinyl ether, respectively. Copolymers containing ethylene and TFE, called ETFE copolymers, have been successfully activated and grafted leading to ETFE-g-poly(M) graft copolymers, where M represents acrylic acid [167,168,200], ethyl acrylate [201], methyl acrylate [201], dimethylaminoethyl methacrylate [202], N,N-methylene-bis-acrylamide [203], NVP [202], 1-vinyl imidazole [202], styrene [167,174,175,204–206] and trifluorostyrene [172]. The last two monomers were used for the preparation of IEMs (Table 5). To improve the IEC up to 1.75 meq g1, some of ETFE-g-PAAc membranes were sulphonated [200(a)], leading to both sulpho- and carboxyl groups that brought about an increase in the IEC values. Some of the IEMs were suitable for fuel cell applications after sulphonation of ETFE-g-PS graft copolymers [167,174,175,204–206] (Table 5). Various relationships [167] between the membranes’ properties and composition, their IECs and their respective fuel cell performances were reported mainly by Horsfall and Lowell [167,168,204,205] or by the Paul Scherer Institute [191,206,207]. These laboratories also noted that such new types of membranes perform well by comparison to Nafion® standard [204]. Some membranes were found to have stable resistivities at high current density and high power density, greater than 1 A cm2 [167]. (A) Activation of poly(E-co-TFE) copolymers followed by grafting.

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Renaud Souzy and Bruno Ameduri

Table 5 Fluorinated graft copolymers and their properties compared with those of Nafion® membrane, achieved from the irradiation of fluoropolymers (FP) followed by grafting of M monomer, for fuel cell applications. AAc, AEM, 4-AS, DVB, d.o.g., e-Beam, EWC, FP, IEC, S, SFT, STFS, 4-VP and σ stand for acrylic acid, anion-exchange membrane, 4acetoxy styrene, divinyl benzene, degree of grafting, electron beam, equivalent weight capacity, fluoropolymer, ionic-exchange capacity, styrene, swollen film thickness, substituted trifluorostyrene, 4-vinyl pyridine and conductivity, respectively FP PTFE (film)

PTFE

PTFE

PTFE

Way of activation

M

Total dose, S, then 20 kGy sulphonation Dose rate, (ClSO3H) 110 Gy min1 60

Co (1.3– S, then 15.0 Gy hr1 sulphonation

γ-rays

S, then sulphonation (ClSO3H)

d.o.g. (%)

IEC (meq g1)

Properties

Reference

1–31

0.7–0.9

σ up to 11.8 mS cm1, SFT  73–140 μm

[151]

Variable

0.4–1.2

Water uptake, good [152–155] thermal, chemical and mechanical stabilities

Variable

Variable

60 Co 4-VP, then 0.7–13.0 (1–50 kGy) quaternisation

Variable σ

[156]

Low High transport number [157,158] specific (0.93) and diffusion electrical coefficient, good resistance mechnical properties, low specific resistance, AEM

PVDF

e-Beam (20 Mrad)

AAc and 4-AS

Variable

Variable

Electrical Resistance 4.2 Ω cm1, IEM

[159]

PVDF

γ−rays (3 Mrad)

S, then sulphonation (H2SO4)

18–30

0.7–1.7

σ  30 mS cm1, Electrical Resistance 1.4 Ω cm1, H2O uptake 37–60, SFT  100–111 μm

[160]

PVDF

PVDF (80 μm film)

PVDF

e-Beam S, then Variable Variable (20–200 kGy) sulphonation up to 100% 175 kV (ClSO3H) 15 Mrad

60

Co (6.3 Mrad)

σ up to 120 mS cm1, [161–163] 57% swelling rate at 95°C, good thermostability

S and DVB, then sulphonation (H2SO4)

18–30

0.68–1.70

37–60% H2O uptake (at 95°C), σ  20–30 mS cm1, SFT 100–111 μm, Overall cell resistance, 6.5 Ω cm2

[164]

VBC, then amination

26

0.68–1.70

8–30 mS cm1, 37–60% H2O uptake, SFT  100–111 μm, Alkaline AEM

[165,166]

Functional fluoropolymers for fuel cell membranes

FP

Way of activation

M

d.o.g. (%)

IEC (meq g1)

Properties

497

Reference

FEP

60 Co (10–30 kGy)

AAc

variable

IEC vs. d.o.g.

Resistance vs. d.o.g., EWC  50–65%, variable σ

[167,168]

FEP

60 Co (3.5 kGy)

AAc, then sulphonation

5–32

0.5–2.8

4–40% H2O uptake, stable performances after 100 h at 50°C

[169]

S, then sulphonation (ClSO3H)

19

1.39

68% H2O uptake, SFT  78 μm

[160]

Variable

[170,171]

FEP

FEP

60

Co (6 Mrad) 60

Co

S, then Variable Variable sulphonation

FEP

60 kGy

TFS, then sulphonation

19

1.27

Cross-linking rate 12% Swelling rate 16% in H2O at 100°C; σ, 30 mS cm1, specific resistance, 35 Ω cm1

[172]

FEP

60 Co (10–30 kGy)

VBC, then amination

3–30

Variable

σ = 20 mS cm1 Thermal stability, AEM

[165,166,173]

ETFE ETFE ETFE

60

Co

60

Co

e Beam (1.5 MeV) 80 kGy

ETFE

e Beam

Nafion

Variable Variable Mechanical and thermal stabilities, IEM

60 Co S, then Variable Variable (12.7 kGy h1) Sulphonation

ETFE

PFA

AAc

e Beam (100–1200 kGy)

None

[167]

variable σ

[167]

S, then Variable sulphonation

1.9–2.0

Variable σ, low methanol crossover, good power performance, promising for DMFC

[174]

S, then Variable sulphonation (ClSO3H)

1.39

Cross-linking by DVB, Reduction of MeOH crossover

[175]

IEM for fuel cell

[172]

Variable mechanical properties and electrochemical properties vs. d.o.g., PEMFC

[176–180]

STFS

Variable Variable

S, then Variable Variable sulphonation

None

None

0.67–1.25

37% H2O uptake, [40,49,73,82] SFT, 209 μm, specific Resistance, 12.8 Ω cm2, σ, 50–110 mS cm1

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Renaud Souzy and Bruno Ameduri

Interestingly, this second group compared the performances of the membranes obtained from PVDF-g-PSSA, ETFE-g-PSSA and FEP-g-PSSA and found that the one obtained from the ETFE polymer exhibited better mechanical properties, because of (i) the higher molecular weight of ETFE; (ii) the better compatibility of ETFE with the graft component: and (iii) the reduced extent of radiation – induced chain scission occurring on ETFE [191,195,206,207]. (B) Activation of poly(TFE-co-HFP) or FEP, followed by grafting. FEP (or fluorinated ethylene-propylene) resin is a linear, semicrystalline fluoroplastic containing TFE and HFP base units [30(a),208]. Early investigations on the radiation chemistry of FEP were achieved by Florin and Wall [209]. It is well known [30(a),145,208] that FEP undergoes a degradation when it is irradiated in vacuum below its Tg (80°C). Forsythe and Hill [145] reviewed well the influence and the intensity of the dose on the activation, cross-linking or degradation of FEP. FEP has also been an interesting raw material for grafting various monomers such as acrylic acid [167–169,212], MMA [171], styrene [160,170,171,173,213,214], divinyl benzene [177–180], substituted trifluorostyrene [172] and vinylbenzyl chloride [165,166,173]. In addition, useful cross-linking agents such as divinyl benzene [175,214] and N,N-methylene-bisacrylamide [203] were grafted to enhance the mechanical properties of the membranes. As above, most investigations were developed on the synthesis of FEP-gPS by Horsfal and Lowell [167,204,205], Nasef et al. [171,213], Slade’s team [165,166,173], and by Scherer’s group [170,206,207,214], as precursors of original FEP-g-PSSA (after sulphonation of the aromatic rings) (Table 5). Interestingly, these well-architectured FEP-g-PSSA copolymers were rather stable (the desulphonation occurring from 200°C) and were used as original PEMs. The most interesting results were achieved by Horsfall and Lovell [167,168], who made some correlations between the d.o.g., the IEC and the equilibrium water content, starting from various fluorinated (co)polymers such as FEP, ETFE, and PFA. Danks et al. [165,166] compared the chemical and thermal behaviours of PVDF-g-PVBC and FEP-g-PVBC graft copolymers. They showed that the former underwent a dehydrofluorination during the functionalisation step (amination). But this step was successfully achieved from the FEP-g-PVBC graft copolymer on which ammonium sites were introduced for potential membranes in alkaline fuel cell applications [165,173]. Although the IEC decreased more rapidly at 100°C, amino FEP-g-PVBC copolymers were stable at 60°C, for at least 120 days and their conductivity values were satisfactory (0.01–0.02 S cm1 at room temperature).

Copolymers of TFE and PAVE (CF2  CFOCnF2n1, mainly n  3) are called PFA [30(a),215].

(C) Poly(TFE-co-perfluoropropylvinylether), poly(TFE-co-PPVE) or PFA.

Functional fluoropolymers for fuel cell membranes

499

In contrast to the vast work concerning the synthesis of graft copolymers starting from the irradiation of ETFE or FEP, only a little work has been conducted on the activation of PFA. Except for a few studies dealing with the crosslinking of PFA by electron beam [216], to the best of our knowledge, Nacef’s [176–180] and Cardona’s [217] groups worked on the synthesis of PFA-g-PM comb-like copolymers by irradiation of PFA followed by grafting. Indeed, this former team [176–180] investigated the irradiation of PFA followed by the grafting of styrene to achieve PFA-g-PS graft copolymers, which were then sulphonated to yield PFA-g-PSSA as a precursor of fuel cell membranes (Table 5). These authors optimised the conditions of activation/grafting in terms of monomer concentrations, irradiation doses and dose rates and the choice of solvents for the grafting process [176]. As in the above cases, the higher the d.o.g., the greater the monomer concentrations, until the styrene concentration reached as high as 60 vol%. Nasef et al. [177–180] extensively studied the morphology and various electrochemical, thermal and physicochemical properties of these membranes: swelling behaviour, IEC hydration number, ionic conductivity, and water uptake, which increased when the d.o.g. was increased, just like the thermal and chemical stabilities [178]. These properties were dependent upon the d.o.g. They also noted that the degree of crystallinity decreased with an increase in grafting and that both of the above mechanical properties decreased when the d.o.g. increased [177]. In addition, XPS [179,180] was used by the authors to monitor the membrane degradation after the fuel cell test. Indeed, an oxidation degradation took place in the PFA-g-PSSA membrane during fuel cell tests, due mainly to the chemical attack at the tertiary H of α-carbon in PS side chains [177,179,180]. 3.2.3.2.4. Conclusion A wide range of different fluorinated polymer-gpoly(M) comb-like copolymers was synthesised from the activation of fluoropolymers (mainly by γ -rays, or electronbeams) followed by grafting, and then used as original fuel cell membranes. Their mechanical, physicochemical, thermal and electrochemical properties were studied. The quality of the well-architectured fluorinated copolymers depends upon the nature of the activation of the fluoropolymer and of the good compromise to generate as many radicals as possible without affecting the properties of the fluoropolymeric backbone. Interestingly, different states (powders and films) and nature of fluoropolymers were activated, and most starting polymers are PVDF, PTFE, ETFE, FEP, and PFA. A number of papers have been reported on the comparison of membranes prepared under similar conditions from the above (co)polymers. For example, considering two isomers, Danks et al. [165,166] noted that the performance of PVDF was drastically affected after the amination reaction of PVDF-g-PVBC, which led to dehydrofluorination. Conversely, ETFE was a better candidate that

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offered potential membranes for alkaline fuel cells [165,166]. The comparison of grafting monomers to various polymeric substrates and the performances of the resulting graft copolymers were achieved by different teams [168,184,204]. Correlations between the d.o.g., IEC and equilibrium water content were shown to be independent of the chemical nature of the fluorinated starting polymer [168]. 4. CONCLUSION Fluoropolymers are known to be used in many high-tech fields and their characteristics also enable them to be involved in fuel cell membranes. Indeed, the fluorine atoms have positive effects such as: (1) the improvement of the thermal, chemical and oxidizing stabilities of the resulting (co)polymers and of their mechanical properties to some extent; and (2) enhancement of the acid behaviour of a sulphonic acid function when adjacent to a fluorinated group, hence enabling the resulting membrane to exhibit good protonic conductivities. This chapter provides an up-to-date review of the syntheses, properties, and applications of membranes for fuel cell applications. Various main routes have been successfully used to produce fluoropolymers of controlled architecture, which have been basically separated into two main families: (a) Those achieved from the direct radical copolymerisation of functional (i.e. SO2X, with X  F, ONa, OK, OH, CO2H, or P(O)(OH)2) monomers with commercially available fluoroalkenes (TFE, VDF, HFP, CTFE, etc.) Most works have been achieved on aliphatic copolymers and some of them are already produced on pilot or industrial scales, in contrast to fluorinated aromatic copolymers, for which few research attempts have been carried out (except for BAM3G membrane). (b) The chemical modification by irradiation of fluoropolymers (PTFE, PVDF, ETFE, FEP, and PFA) followed by grafting of proton-exchange monomers or PS (post-functionalisable into sulphonic acid PS). The obtained copolymers seem to suffer from oxidising decomposition, concomitant with thermal degradation. Although some conductivity values seem to be interesting, the presence or absence of the oxygen atom as a link between the polymeric backbone and the grafts is sometimes not mentioned. In addition, the chemical modification needs special, heavy and expensive equipment to enable irradiation, and industrial production is probably difficult to achieve. The development of fluorinated materials for electrolytes for fuel cells requires the synthesis of original functional fluorinated monomers for obtaining new polymers as well as the optimisation of the characteristics of existing polymers with regard to molecular weight, cross-linking, new functions (e.g. too few polymers bear phosphonic acid function) or may be other alternatives as the formation of hybrid or composite membranes or those made of multilayers.

Functional fluoropolymers for fuel cell membranes

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In spite of its cost, its permeability to methanol and its recyclability [49], Nafion® seems to be the only fluorinated copolymer to exhibit mechanical, chemical and proton-exchange properties useful for use as membranes for fuel cells, and much research is still necessary to find new alternatives to Nafion®. Hence, further investigations need to be conducted, and scientists should be motivated to pursue such exciting researches. ACKNOWLEDGEMENTS The authors acknowledge the Centre National de la Recherche Scientifique (CNRS), the French consorption GDR PACEM 2479 (CNRS) and the Commissariat à l’Energie Atomique (Grenoble – France) for financial support and stimulative discussions. REFERENCES [1]

[2] [3]

[4] [5] [6] [7] [8]

[9] [10] [11]

[12] [13] [14] [15] [16] [17] [18] [19] [20]

(a) W.G. Grot, Macromol. Symp., 82 (1994) 161–172; (b) S. Baneerjie and D.E. Curtin, J. Fluorine Chem., 125 (2004) 1211; (c) M.A. Hickner, H. Ghassemi, Y.S. Kim, B.R. Einsla, and J.E. McGrath, Chem. Rev., 104 (2004) 4587. K.D. Kreuer, Solid State Ionics, 97 (1997) 1. W. Vielstich, A. Hubert, M. Gasteiger, and A. Lamm (Eds.), Handbook of Fuel Cells— Fundamentals, Technology and Applications, Fuel Cell Technology and Applications, Vol.3, Wiley, New York, 2003. Q. Li, R. He, J.O. Jensen, and N.J. Bjerrum, Chem. Mat., 15 (2003) 4896. J.A. Kerres, J. Membr. Sci., 185 (2001) 3. O. Savagado, J. New Mat. Electrochem. Syst., 1 (1998) 47. K.D. Kreuer, J. Membr. Sci., 185 (2001) 29. D. Jones and J. Roziere, J. Membr. Sci., 185 (2001) 41 and W. Vielstich, A. Hubert, M. Gasteiger, and A. Lamm (Eds.), Handbook of Fuel Cells—Fundamentals, Technology and Applications, Fuel Cell Technology and Applications, Vol.3, Wiley, New York, 2003. pp. 447–463. M. Rikukawa and K. Sanui, Prog. Polym. Sci., 25 (2000) 1463. G. Inzelt, M. Pineri, J.W. Schultze and M.A. Vorontyntsev, Electrochim. Acta, 45 (2000) 2403. K.D. Kreuer, Handbook of Fuel Cells—Fundamentals, Technology and Applications, Fuel Cell Technology and Applications, W. Vielstich, A. Hubert, M. Gasteiger, and A. Lamm (Eds.), Vol.3, Wiley, New York, 2003, pp. 430–446. G.D’Alelio, General Electric, US Patent 2,366,007, 1944. J.M. Abrams, Ind. Eng. Chem., 48 (1956) 1469. K. Prater, J. Power Sources, 29 (1990) 239. S. Faure, R. Mercier, P. Aldebert, M. Pineri, and B. Sillion, CEA, Fr. Patent 9,605,707, 1996. R. Nolte, K. Ledjeff, M. Bauer, and R. Mülhaupt, J. Memb. Sci., 83 (1993) 211. R. Nolte, K. Ledjeff, M. Bauer, and R. Mülhaupt, BHR Group Conf. Ser. Publ., 3 (1993) 381. F. Helmer-Metzman, F. Osan, A. Schneller, H. Ritter, K. Ledjeff, R. Nolte, and R. Thorwirth, European Patent 574,791,A2, 1993. B. Adams and E. Holmes, J. Soc. Chem. Ind., 54 (1935) 17. A.S. Hay, General Electric, US Patent 3,432,466, 1969.

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List of Symbols and Abbreviations AAc AEM AIBN CTFE DMF DVB DPn DSC DMFC d.o.g. ETFE FEP HFP HFPO IEC IEM MEA Mn Mw NMP PAA PCTFE PEOMA PEMFC PFA PFCB PM PMVE PSEPVE PSSA PTFE PVDF RT S SFT TFE

acrylic acid anion-exchange membrane azobisisobutyronitrile chlorotrifluoroethylene dimethylformamide divinyl benzene average degree of polymerisation in number differential scanning calorimetry direct methanol fuel cell degree of grafting poly(ethylene-alt-tetrafluoroethylene) copolymer poly(tetrafluoroethylene-co-hexafluoropropene) copolymer hexafluoropropylene hexafluoropropylene oxide ion-exchange capacity ion-exchange membrane membrane electrode assembly average molecular weight in number average molecular weight in weight N-methyl pyrolidinone poly(acrylic acid) polychlorotrifluoroethylene (CF2 CFCl) poly(ethylene oxide methacrylate) proton-exchange membrane fuel cell poly(tetrafluoroethylene-co-perfluoropropyl vinyl ether) copolymer perfluorocyclobutane poly(monomer) perfluoromethylvinyl ether perfluorosulphonyl fluoride ethoxy propyl vinyl ether [or perfluoro(4methyl-3,6-dioxaoct-7-ene)-1-sulphonyl fluoride] poly(styrene sulphonic acid) polytetrafluoroethylene polyvinylidene fluoride room temperature styrene swollen film thickness tetrafluoroethylene

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TFS Td Tg TGA TVFOB VBC VDF XPS UV 4-VP

trifluorostyrene decomposition temperature glass transition temperature thermal gravimetric analysis α,β,β -trifluoroethenyl oxybenzene vinyl benzyl chloride (or chloromethyl styrene) vinylidene fluoride X-ray photoelectron spectroscopy ultraviolet 4-vinyl pyrridine

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 22

Films and powders of fluorine-doped tin dioxide Hubert Cachet Université P. & M. Curie, Case courrier 133 Laboratoire LISE – UPR 15 CNRS - 4, Place Jussieu – 75252 Paris Cedex 05 - France 1. INTRODUCTION This chapter reviews the up-to-date developments and applications of fluorinedoped tin dioxide materials that have attracted the attention of a large number of research groups from several disciplines in the last two decades. Undoped tin dioxide (SnO2 and TO) is a wide bandgap (theoretical calculated value 3.6 eV [1]) semiconductor, with an n-type character arising from oxygen vacancies. The electrical conductivity can be largely improved and stabilized by doping with foreign impurities. The most common doping elements are either antimony by substituting Sn4 tin cations as Sn1x4Sbx5O22  x  e or halogens (X  Cl, F and Br) by substituting oxygen as Sn4O2x2X  x  e. Chlorine is often an unintentional dopant because of the precursor chemistry, but it is much less efficient than fluorine. With the latter dopant, low electrical resistivity values have been achieved with polycrystalline F-SnO2 films (FTO), approaching 1  104 Ω cm with a carrier density in the range (1–20)  1020 cm3 [2]. This high electrical conductivity and high optical transparency in the visible range, associated to a satisfying mechanical, chemical and electrochemical stability, make SnO2 films well suited for the practical applications requiring transparent conductive electrodes (TCE) or near-infrared reflective coatings. Unlike antimony, fluorine is incorporated into the SnO2 lattice without generating large densities of electronic gap states [3]. Eventhough highly doped, the semiconducting behaviour of tin dioxide is maintained: the consequence is a large overvoltage pointed out for oxygen evolution, which allows the use of tin oxide films as anode for the electrochemical oxidation of organic pollutants and wastewater treatment based on generation of hypochlorous species [4]. Transparent conducting oxide (TCO) materials are involved in a large variety of applications including flat-panel displays, windows of buildings, smart windows, thin-film photovoltaic solar cells, photo- and spectroelectrochemistry.

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Among the available TCOs, FTO is of great technological and economical importance, justifying particularly the large number of research and industrial works on thin film and powder synthesis. FTO materials have been produced by various techniques: spray pyrolysis deposition (SPD) [5–29], rf or dc reactive magnetron sputtering [30–34], chemical vapour deposition (CVD) [35–52] spincoating [23,53] and dip-coating techniques [54,55] and via soft chemistry route [23,56–61]. Among the above techniques, SPD and CVD under various methods appear to be the most investigated deposition techniques. The description of the characteristics and specificities of each preparation techniques will be discussed in Section 2. This will be followed by a survey of the effects of fluorine on the electrical and structural properties of FTO in Section 3. Some aspects of the optical properties of FTO will be presented in Section 4, with regard to applications such as spectrally selective coatings or UV detectors. Fluorine incorporation into SnO2 matrix has been studied by many authors on the basis of analytical techniques, especially by resonant nuclear analysis (RNA). The most significant results obtained in this domain will be presented and discussed in Section 5. Practical applications of FTO in (photo)electrochemistry will be illustrated in Sections 6, excluding the case of FTO in lithium batteries treated in Chapter 5. In Section 7, various applications of FTO films and powders are briefly described, including some limited aspects of photovoltaic devices (treated in detail in Chapter 23) and in the field of electrical engineering. 2. SYNTHESIS OF FTO FILMS AND POWDERS 2.1. Spray pyrolysis deposition

Spray pyrolysis synthesis of tin oxide thin films is a low-cost technique, largely used because of its simplicity and its ability to vary the chemical composition of the solution containing the tin and fluoride precursors very easily. An aerosol of liquid droplets is produced either by using a nozzle and a carrier gas (nitrogen, oxygen, or compressed air) or by atomizing the solution by means of ultrasonic waves (Pyrosol process). The aerosol is formed of micron-sized liquid droplets. It is thermally decomposed on a heated substrate, in the temperature range 350–600°C. The oxidizing agent may be oxygen or chemical species added to the spray solution as H2O2 [11]. An Arrhenius behaviour is generally observed for the deposition rate for temperatures lower than ⬇ 500°C, indicating a kinetic control of the deposition process. At higher temperatures, a levelling-off often occurs, indicating a diffusion control. Note that in the case of a carrier gas spray set-up, the actual deposition temperature is poorly defined because of the gas expansion. The dispersion in deposition temperatures may reach a few tens of degrees. Then it is not proper to compare in detail the deposition temperatures given in the literature, which depend on the experimental arrangements. For compensating this effect, some authors used intermittent [62] or X–Y scanned spray

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systems [3]. This problem is largely minimized when an ultrasonic atomizer is employed. The SPD process was successfully applied on an industrial scale, especially by glass manufacturers to elaborate window coating reflecting nearinfrared radiations from the 1980s. For instance, a 3-m-wide glass, heated at 600°C, was passed at 4.5 m/min on an annealing platform. The FTO coating was obtained by separately spraying an aqueous solution of SnCl2 and NH4F–HF [63]. Another industrial process was based on pyrolysis of a powder, for instance, dibutyl-tin(IV) difluoride, which was applied as a suspension on the moving hot glass surface [64,65]. As the tin feedstock, a variety of organometallic compounds have been considered, most of them containing chloride. The most commonly used are tin(IV) tetrachloride (SnCl4) and butyl-tin(IV) trichloride (SnBuCl3) dissolved in alcoholic or hydroalcoholic solutions. Other precursors used were dimethyl-tin(IV) dichloride, dibutyl-tin(IV) diacetate (DBTDA) and tin(II) dichloride. Chlorine atoms act as a co-dopant with fluorine. The latter is commonly introduced into the spray solution either as NH4F or HF. An original approach was to synthesize FTO films from HCl-acidified stannous fluoride methanolic solutions [66]. The effect of halogen doping (F, Cl and Br) on growth kinetics was studied by Agashe and Major [67]. The strength of the halo acids and hence the electronegativities of the halogens played an important role in governing the reaction process, which controlled the growth kinetics of the films. Growth rate (⬇1 nm s1 for precursor concentration in the decimolar range) was found to be considerably affected by F doping, especially at higher doping levels, whereas Cl and Br doping reduced the growth rate to a large extent [68]. A possible explanation was the production of highly halogenated gases that can etch the growing SnO2 layer [14]. The effect of F doping on growth rate seems to depend on the experimental conditions. In another experimental work, no difference was found between undoped and F-doped SnO2 for deposition temperatures  500°C [3]. Above this temperature, a net increase in deposition rate was measured (Fig. 1). This was tentatively attributed to an easier incorporation of F atoms; however, at lower temperatures F and Cl atoms are incorporated simultaneously [69]. Spray pyrolysis has been largely used for nanoparticle synthesis [70]. An evaporation – reaction type aerosol generator for ultrafine tin dioxide particles was developed to promote homogeneous nucleation before collecting oxide particles [71]. To our knowledge, however, this technique has not yet been applied to fabricate F-doped tin oxide nanoparticles. 2.2. Sputtering deposition

A few attempts were made to deposit fluorinated SnO2 from solid targets. A group of Japanese researchers prepared F–SnO2 by rf magnetron sputtering from a target made with SnO2 powder mixed with either SnF2 or SnF4 powder [30]. Conductive films were obtained but the crystallinity was inferior to that of

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Deposition rate (nm/s)

3.0

2.5 ATO 2.0

FTO TO

1.5

1.0 360

400

440

480

520

560

substrate temperature (˚C)

Fig. 1. Effect of doping and substrate temperature on the deposition rate of tin oxide (TO) films sprayed from a 0.2 M SnCl4/CH3OH solution; undoped TO films (䊐); F-doped (F/Sn0.7) FTO films (o); Sb doped (Sb/Sn0.04) ATO films (䉱) (from A. Messad, Ph.D. Thesis, University Paris VII, December 1993).

the CVD prepared films, irrespective of the sputtering parameters (substrate temperature, dc power, O2 gas pressure or F precursor gas pressure). The more recent works deal with dc reactive magnetron sputtering using a metallic tin target and various plasma atmospheres: argon/oxygen/CF4 [31], argon/oxygen/SF6 [32] and argon/oxygen/freon [33,34]. The main interest in the sputtering technique is to allow film deposition at relatively low temperatures, i.e. 180°C in Reference [31]. The drawbacks are: (i) a mixed composition with the presence of Sn2 and Sn4 species, implying the presence of different phases such as SnO, Sn3O4, Sn2O3 and SnO2, as proved from XRD and XPS analyses [34] and (ii) a low electrical activation of incorporated fluorine attributed to the formation of tin fluorides, which yield film resistivities above 103 Ω cm. As recently demonstrated, the sputtered films tend to be crystallized when the stoichiometry approaches that of SnO or SnO2, being amorphous in between [34]. 2.3. Chemical vapour deposition

As the spray pyrolysis technique, CVD has been extensively investigated by researchers [35–38,41,47,49–51,72–74] and especially by industrial companies for large-scale production or solar cell applications [75–82]. Tin precursors were either highly toxic tetramethyl tin (Me4Sn), chlorine substituted compounds (Me2SnCl2, BuSnCl3), DBTD or other organo-metallic compounds with fluorinated ligands to circumvent the need for a separate fluorine dopant (Sn(II)(O2CCF3)2, (n-Bu)2Sn(O2CCF3)2). The vector gases were O2, O2/O3 or O2/N2 mixtures. In some cases, trifluoroacetic (TFA) and fluorohydric (HF) acids

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were used as fluorine sources. Halogen substituted methane and ethane were the most commonly used fluorine precursors in CVD processes (CF3Br, CHF3, CHClF2, CCl3–CF3 and CF3–CH2F). The overall thermal CVD process is composed of two steps, the evaporation step (in the range 100–300°C) and the deposition step itself (in the range 400–700°C). As for SPD, a relatively high deposition temperature is required for thermal CVD. This is the ability of the plasma-assisted CVD to reduce the deposition temperature to room temperature, allowing a less conductive tin oxide film to be deposited on plastic substrates [51,75]. For instance, SF6 was used as the doping precursor to be introduced in an O2/Ar/ Me4Sn gas mixture in very small quantities [51]. Other promising routes for preparing FTO films reported in the literature were “electrostatic spray pyrolysis” (ESP) [83] and “electrostatic-assisted vapour deposition” (ESAVD) techniques [84]. In both the cases, a liquid containing the precursors (SnCl4/HF/C2H5OH [83]; Sn(CHOO)2/HF/CH3OH [84]) was forced through a capillary that was subjected to an electric field. The emitted and charged droplets were decomposed into vapour and reacted when they contacted the heated substrate. Interestingly, 500-nm-thick FTO films with a resistivity of 2  104 Ω cm were produced at a deposition temperature of 600°C [84]. 2.4. Soft chemistry route

Thin films prepared by the soft chemistry route are usually deposited by the dip-or the spin-coating techniques, both of them leading to uniform deposits. The dip-coating technique is well suited for coating large-area substrates or having non-flat morphology. Undoped SnO2 was synthesized by reacting citric acid and tin citrate in ethylene glycol in the presence of small amounts of nitric acid at 60°C [85]. Further heating at 110°C eliminates water and nitric acid traces and an esterification reaction takes place producing a more viscous solution. A tindoped indium oxide (ITO)-coated glass was dipped in this solution at a low speed, dried in an oven at 140°C and then calcinated at 500°C. The result was a 1-μm-thick transparent and homogeneous SnO2 film. This strategy was successfully extended to Sb doping of tin oxide, yielding conductive films with resistivity in the range 103 Ω cm for a [Sb]/[Sn] ratio over 6% [86]. FTO films were synthesized by a dip-coating technique, based on the hydrolysis of a SnCl2·2H2O/NH4F methanolic solution [55]. When the substrate is withdrawn vertically at 1.3 mm s1, a 0.6-μm-thick film is formed in each dipping cycle. Continuous films are obtained after several cycles, which appear to be a mixed phase consisting of crystalline tin oxide over an amorphous background, probably Sn(OH)Cl. It was found that heat treatment at 400–500°C was necessary to obtain films of useful quality, with a resistivity of 0.024 Ω cm and an average transmission of ⬇85% at a thickness of ⬇ 3 μm [55]. Recently, the sol–gel dipcoating technique was applied to SnCl2·2H2O/HF solution in isopropyl alcohol. Highly transparent films with a nanocrystalline structure (6 nm grain size) and a

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1 Ω cm resistivity were found, useful, for instance, as antistatic coating [87]. Another original sol-gel route was initiated on the basis of a new F-containing Sn(IV) molecular precursor (tert-amyloxyfluorodipentan-2,4-dionatotin(IV) abbreviated as SnF(OtAm)(acac)2) with the specificity of pre-existing Sn–O and Sn–F bonds, as they exist in the final FTO material [23]. In the process, the organometallic precursor undergoes hydrolysis and condensation reactions to form a preformed oxide network in solution. A stable xerosol was obtained with the composition SnF(OtAm)(acac)2/CH3CN/2H2O, from which either highly conductive nanometer-sized FTO powders [60] or FTO films by SPD or spincoating techniques [23] were effectively synthesized. 3. STRUCTURAL AND ELECTRICAL PROPERTIES At a given deposition temperature, the structure of SPD tin oxide films using SnCl4 as tin precursor was found to be governed by two factors: (i) the tin concentration and (ii) the doping by fluorine [88]. Transmission electron microscopy (TEM) observations have shown that the grain size (i) decreases with the decrease in tin concentration in the spray solution and (ii) increases with the increase in the deposition temperature. The grain sizes were typically in the range 30–80 nm for deposition temperatures between 400 and 500°C in the case of a concentrated SnCl4 (0.2 M) solution [89]. Doping by fluorine had no influence on the mode of growth or on the grain size. FTO films contain planar defects like stacking faults and twins. Fig. 2(a) gives a high-resolution TEM image obtained for an undoped sample. This image shows the presence of {011} cassiterite twins, which are parallel to the growth direction. The doping by fluorine induces a significant increase in the density of planar defects during the growth, independent of the effect of the SnCl4 concentration [89,90]. This is visualized in the TEM image in Fig. 2(b) and the electron diffraction pattern in Fig. 2(c), both relative to the same highly doped FTO film. Fig. 2(c) shows the presence of streaks, which are characteristic of the high density of planar defects. For FTO films, the density of twins reaches 5  1012 cm2 at a variance with density much lower than that observed for undoped tin oxide films. X-ray and electron diffraction studies reveal a 100  preferred orientation with the (200) planes parallel to the substrate (Fig. 2(c)). Such a 100  preferred orientation was also found for undoped TO films grown by SPD from DBTDA solutions [91]. The introduction of an F doping with NH4F drastically decreased the peak intensity of the (200) plane, yielding non-preferentially oriented FTO films. A preferential orientation for FTO is recovered if the FTO films are deposited on highly oriented TO films [91]. The morphology of FTO films grown from SnCl2 ((110) preferred orientation [92,93]) or SnCl4 ((200) preferred orientation [88,91]) solutions was discussed in detail, taking into account the conditions of growth and the presence of the cassiterite twins [93]. In terms of optical and electrical characteristics of the

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Fig. 2. Effect of F doping on the microstructure of spray-deposited tin dioxide films. (a) HRTEM image of an undoped TO films showing the (110) planes separated by 0.335 nm and the lower density of defects. (b) Low-magnification TEM image visualizing the high density of planar defects (from Bruneaux [88], with permission of Elsevier). (c) Diffraction pattern revealing a 100  preferential orientation. The streaks are induced by a high density of planar defects arising from the F incorporation (from Bruneaux [88], with permission of Elsevier).

deposits, the best properties (electrical resistivity as low as 2  104 Ω cm or optical transmittance in the visible range  90%) are obtained when the development of (101) twin planes during the film growth is inhibited. It is concluded that the density of twins is a determining factor for the intragranular electrical resistivity of FTO films with respect to an application as a transparent electrode. Electrical conductivity σ is defined as σ  qNμ, where q is the elementary charge, N the carrier (free electrons for FTO) concentration and μ the carrier mobility. In single-crystal materials, μ is defined by the effective mass of free carriers, the dominating scattering mechanism and the type and density of scatterers. In polycrystalline films, grain boundaries may additionally contribute to limit the carrier mobility. In a recent work, electronic properties of degenerated

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FTO films prepared by atmospheric pressure CVD were studied by combining electrical resistivity, Hall and Seebeck effects and plasma and collision frequency measurements [94]. From the variations of the plasma frequency with the carrier density, the reduced effective mass value m* was found independent of carrier concentration (for 2  1020  N  6  1020 cm3), which indicates a parabolic conduction band. Its value was m*  0.28 0.02 for the FTO films, in agreement with previous literature data [95]. The Hall mobility was found to be very close to the optical mobility, showing a very small contribution of grain boundaries to the total resistivity of films. Scattering by grain boundaries can be neglected to the first approximation because the grain size (30–100 nm) is much greater than the electron mean free path (⯝10 nm). At variance, the high density of twins generated during the film growth may also cause electron scattering, because the mean distance between neighbouring crystallographic defects may become close to the mean free path [88]. Thermopower measurements along with the 300–500 K temperature independent mobility (30 cm2 V1 s1) indicated that electron scattering by impurity ions screened by free carriers would be the dominating scattering mechanism [94]. The effect of grain boundaries has been evidenced by analysing the dependence of resistivity ρ on the carrier density measured at room temperature [88]. Hall and resistivity measurements were performed on a series of fluorine (FTO) and chlorine-doped (TO) films. SPD was formed at 500°C, which covered the range 7  1017  N  4  1020 cm3 of carrier concentration. A unique curve ρ (N) was obtained for all the samples as depicted in Fig. 3(a). Such a behaviour has been accounted for within the frame of the one-dimensional grain boundary model developed for polycrystalline silicon films [96]. Grain boundaries act as carrier traps with an estimated density of 4  1012 cm2 creating more or less extended electron depleted regions. The overall resistivity ρ can be expressed as the weighed sum of two contributions, one from the space charge regions and the other from the neutral bulk regions. For carrier concentrations N  51018 cm3, the resistivity is thermally activated and mainly governed by the barrier effect of grain boundaries, which act as carrier traps. The mobility drops for N ⬇ 5  1017 cm3, which would correspond to a total depletion of each grain (Fig. 3(b)). For N  5  1018 cm3, the film resistivity is solely governed by the properties of the bulk material. For degenerated tin oxide films, the influence of grain boundaries is fully negligible, with a quasi-mobility constant (Fig. 3(b)) and a value partly limited by the high density of structural defects. The effect of the atomic fluorine/tin ratio in the starting spray solution on the resistivity and the carrier concentration has been investigated for the SnCl2/HF system [97]. The results at two substrate temperatures (300 and 400°C) are illustrated in Figs. 4(a) and (b), taken from Reference [97]. They show a strong influence of the deposition temperature with the resistivity (carrier density) difference of about six (four) orders of magnitude, respectively. The carrier density increases

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resistivity (Ω.cm)

100 10 1 0.1 0.01 1E-3 1E-4 1E17

1E18

1E20

1E21

Hall mobility (cm2V−1s−1)

10 1

0.1 0.01 1E17

(b)

1E19

carrier density (cm−3)

(a)

1E18

1E19

carrier density

E120

1E21

(cm−3)

Fig. 3. Dependence of the resistivity (a) and the mobility (b) on the carrier concentration for undoped (o) and F-doped (䊏) TO polycrystalline films. The influence of grain boundaries is only sensitive for N  5  1018 cm3.

and the resistivity decreases with the increase in the atomic F/Sn ratio with a saturation visible for both quantities for F/Sn  60%. The lowest value of resistivity (3  104 Ω cm) was obtained at 400°C for an F/Sn ratio between 60 and 70%. The Hall mobility, calculated from the data in Figs. 4(a) and (b), is found to be strongly reduced with an increase in the F concentration in the solution, indicating that at room temperature, the interaction of free carriers with ionized impurities significantly influences the electrical transport in FTO thin films [97]. Increase in the F/Sn atomic ratio above 60–70% yields a marked increase in resistivity and a lowering of the efficiency of the F doping [98]. Recently, spray-deposited FTO-coated ITO films were developed as new transparent highly conductive film, capable of supporting annealing temperatures in the range 400–600°C [99]. Such a problem was encountered in the fabrication process of dye-sensitized TiO2 solar cells. When ITO films are exposed to such high temperatures, their electrical resistivity increases more than three times due

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Fig. 4. Dependence of the resistivity (a) and the carrier concentration (b) on the value of the F/Sn atomic ratio in the spray solutions for two deposition temperatures (see text) (from Gordillo et al. [97], with permission of Elsevier).

to the interaction with oxygen from the atmosphere. The FTO/ITO resulting films achieved a very low resistivity (1.4  104 Ω cm), keeping an optical transmittance above 80% in the visible range. The change in the film resistivity was less than 10% even when exposed to high temperatures of 300–600°C for 1 h in the air. 4. OPTICAL PROPERTIES Optical properties of FTO films have been intensively investigated, characterized by a direct bandgap of ⬇4.1 eV and an indirect allowed transition of 2.6 eV along with an assisting phonon of 0.08 eV [100,101]. For degenerated highly doped FTO, the fundamental absorption threshold depends on the carrier concentration owing to the occupancy of states at the bottom of the conduction band (Bürstein–Moss effect). Taking advantage of the large bandgap, UV-enhanced and solar blind photodetectors were developed based on a FTO/ZnS junction [102]. The sensors fabricated from these SPD FTO films had an unbiased internal quantum efficiency of nearly 100% in the spectral range 250–320 nm and were insensitive to solar radiation, allowing precise UV measurements under direct solar illumination. FTO films are highly transparent in the visible spectrum range, with a transmission of the order of 80–85%, possibly modulated by interferences in the film as long as the film thickness is  100 nm [103]. The refractive index n(λ) is close to 2.0 with a

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low-frequency dispersion in all the visible spectrum. For wavelengths λ 0.58 μm, an empirical formula n(λ)  {1  λ2/(0.370  λ2  0.0105)}1/2, has been established which is valuable for FTO films [101]. The near-infrared (NIR) behaviour was found to obey the Drude oscillator model, accounting for the change in the carrier density [101,104]. The main parameters of the model are the relaxation time for the conductivity of free electrons τ (or collision time ⬇1015 s) and the plasma frequency ωp/2π (⬇102 THz), which reads ωp2  4π 2Nq2/ε0ε m*, where ε0  8.84  1012 F/m and ε is the relative dielectric permittivity arising from non-conducting electron polarizability effect. The plasma frequency defines the frontier required for a total reflection of electromagnetic radiations. This is the basis for developing IR reflecting coatings for greenhouses [105] or automotive windshields [106], solar absorbers [107], and spectrally selective reflectors for silicon solar cells [108]. 5. FLUORINE INCORPORATION INTO FTO MATERIALS A few studies have been devoted to determine the amount of fluorine and chlorine actually incorporated during the deposition process. Two analytical methods have been used, the secondary ion mass spectroscopy (SIMS) [109] and the resonant nuclear reaction (RNR) technique [98,110]. The latter technique is based on the 19 F(p,αγ )16O reaction with γ -ray detection. The signal can be advantageously calibrated in absolute F concentrations by using a standard with a known amount of fluorine (e.g. LiF or CaF2). SIMS experiments for SPD tin oxide films have shown that the behaviour of the chlorine and fluorine contents vs. the F/Sn ratio in solution at different substrate temperatures are quite different. F content presents a nonmonotonic variation with a maximum around 350°C, with a small temperature dependence. At variance, Cl content strongly varies with temperature with an exponential dependence on the inverse of the temperature [109,111]. RNR technique allows determining the total F content and also the concentration profile through the film in a non-destructive way. Fig. 5 shows the total fluorine concentration in spray-deposited FTO films as a function of the fluorine concentration in the spray solution [98]. It is worthwhile to point out that even for the lowest F doping, the fluorine content in FTO films is larger than the free electron concentration (a few 1020 cm3), suggesting that all the F atoms are not active as the dopants. Fig. 6 gives an example of the depth F distribution determined by RNR in a FTO film having an atomic F/Sn ratio of 5.1%. The width of the F distribution was found to be equal to the film thickness determined by Rutherford backscattering analysis of tin, indicating a constant F concentration through the film. In Table 1, experimental data are gathered on the F content of sprayed FTO films, elaborated from solutions with different F/Sn atomic ratios [112]. RNR measurements give an average composition of the entire beam area. On the contrary, using the electron probe microanalysis technique (EPMA) with a 1 μm

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Fig. 5. Incorporation of fluorine in tin oxide films measured by resonant nuclear reaction (RNR). Atomic F concentration in the films against the F/Sn ratio in wt% (from Acosta et al. [98], with pemission of Elsevier).

Fig. 6. Incorporation of fluorine in tin oxide films measured by resonant nuclear reaction (RNR). F depth concentration profile showing the regular distribution through the FTO film (Unpublished results of the author).

spatial resolution, it is possible to obtain distinct information on the bulk and the relatively spread intergranular regions. It can be seen that a high F/Sn ratio in solution not only yields a large F bulk concentration, but also a strong F enrichment of the grain boundaries. The latter result suggests the possible formation of bad conducting tin fluoride domains in the intergranular region, contributing to an increase in the overall resistivity.

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Table 1 F content measured for various FTO films elaborated by spray pyrolysis from starting solutions with different F/Sn atomic ratios. RNR stands for Resonant Nuclear Reaction and EPMA for Electron Probe MicroAnalysis (Unpublished results of the author) FTO samples

Deposition temperature (°C)

F/Sn in solution (At.%)

Square resistance (Ω)

Resistivity (Ω cm)

F/Sn in films (At.%) RNR

EPMA Bulk intergrain

SnCl4/NH4F/MeOH

500

70

180

0.0029

2.5

SnCl4/NH4F/MeOH

380

140

60

0.0016

™TEGO (Goldschmidt)

500

–

2.2

0.0005

0.8

SnF2/HCl/MeOH

500

200

25

0.0008

13

SnF2/HCl/MeOH

400

200

150

0.003

SnF2(acac)2/ CF3CH2OH

500

200

230

0.07

5.6

2.7

9.5

5.1

Note: RNR, resonant nuclear reaction; EPMA, electron probe microanalysis.

6. ELECTROCHEMICAL APPLICATIONS The electrochemical stability of FTO is limited in the cathodic domain by the reductive decomposition of the oxide. According to Bard et al. [113], the formal potential for cathodic decomposition Edec in aqueous media is pH-dependent and very close to the hydrogen evolution potential. For the overall SnO2/Sn reaction, Edec  0.094  0.06 pH (V vs. NHE), where NHE stands for the normal hydrogen reference electrode. It means that tin dioxide decomposes by electroreduction or by nascent hydrogen, chemically or electrochemically produced. At potentials more positive than Edec, highly doped FTO can be used as an electrode material. In spite of its degenerated state, some limitations depending on the applied polarization potential, arise due to the semiconducting character of TO [114]. It is commonly assumed that electrons are transferred by direct tunnelling from the conduction band in the bulk into the redox electroactive species in solution. Tunnelling is possible because of the high carrier density maintaining the width of the interfacial space charge region 2 nm. Indeed the tunnelling rate is governed by the interfacial potential barrier Eb  Eredox  Efb, where Eredox is the redox potential in solution and Efb the so-called pH-dependent flatband potential. For oxides, Efb is pH-dependent [115]. The control of the electron

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transfer rate by the potential barrier was exemplified by measurements of the exchange current of the ferri – ferrocyanide redox reaction at different pH values [114]. The exchange current was found two orders of magnitude higher at pH 6 than that at pH 13. As a practical conclusion, a metallic character will be observed for small Eb, with the risk to enter in conflict with the FTO stability. On the contrary, for large Eb, a dissymmetry will exist between the anodic and cathodic branches of the current – voltage characteristics. Electroreduction of dissolved gases (CO2 [116] and O2 [117]) at FTO electrodes are good examples of reaction-instability coupling. In the absence of surface modifications, FTO undergoes several reduction stages during CO2 reduction by cyclic voltammetry between 0.1 and 1.4 V vs. SCE (saturated Calomel reference electrode). At the first stage, formyl –CHO– and methoxy –CH3O– groups are formed on the surface. Dissolution of FTO accompanies CO evolution with a competition between CO and H2 evolutions. The surface is covered with elemental Sn, and the oxide electrode acts as a metal electrode [116]. Such a cathodic instability can be overcome by surface modification. For instance, an electrochemical sensor for characterizing biofilm formation in seawater [118] or scale deposit [119] needs to use a stable and efficient transparent electrode for the reduction of dissolved oxygen. This goal was achieved by depositing a very small gold charge (equivalent to a few monolayers) at the surface by a simple chemical treatment [117]. The electrochemical response of the Au-modified FTO was found to be as fast as that of a massive gold electrode. Electrochemical impedance spectroscopy (EIS) showed that SnO2 operates under weak depletion conditions with a surface-state-mediated charge transfer. A similar improvement in the electrochemical response was also achieved in the past by radiolytically grafting iridium nanoaggregates (1016 cm2 surface density) covering 10–20% of the FTO surface [120]. Because of a high overvoltage for oxygen evolution, tin oxide was largely used as anode, in particular for the electrochemical wastewater treatment. Oxidative destruction of biorefractory organic pollutants in industrial waters, as phenol or its derivatives, was successfully achieved [121]. The rate of phenol removal was found to be better on tin oxide electrode than for more conventional anodes as PbO2 or Pt. However, examination of the recent literature in this field shows that antimony is preferred to fluorine for doping (lower overvoltages, better long-term stability); this type of application is out of scope of this chapter. Another important application was the electrochemical generation of chlorine or hypochlorous species for wastewater treatment [122], or surface cleaning of windows immersed in seawater [123] through the electrooxidation of chloride ions present in brine or in seawater. A systematic study of the stability of tin oxide anodes in the chloride oxidation regime was undertaken [124,125]. Localized corrosion phenomena, studied by in situ quartz microbalancce observations, were

Films and powders of fluorine-doped tin dioxide

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observed and related to the interaction with OH° and Cl° electrogenerated radicals. This phenomenon was characterized by a peak-shaped corrosion rate curve, the amplitude of which depended on the pH, the chloride concentration and the type of dopant (F or Sb). The main conclusions from data on chloride solutions at pH 8 and in natural seawater were: (i) the corrosion rate is higher for F than for Sbdoped samples and (ii) the corrosion rate strongly depends on the Cl content, becoming almost negligible at [Cl]  0.3 M, in favour of applications in seawater. Corrosion was attributed to the electron capture of Sn–O surface bonds by OH° and Cl° radicals in competition with oxygen and chlorine evolution reactions. From EIS experiments, it was observed that the corrosion process is also responsible for a positive band-edge shift due to positively charged corrosion intermediates, allowing the FTO electrode to become active with respect to anodic charge transfer reactions. FTO is also commonly used as transparent and conductive substrate in TiO2 dye-sensitized nanocrystalline solar cells (DSSC) [126]. It has been observed that the FTO/TiO2 contact is non-ohmic but has to be considered as rectifying [127]. Owing to the electron affinity difference between TiO2 and FTO, photogenerated electrons are efficiently and irreversibly collected at the FTO electrode [128]. The consequence is a severe limitation of the direct current opposite to the photocurrent, allowing open-circuit potentials as high as 800 mV to be reached. Classically, efficient TiO2-based DSSC operate in non-aqueous media. Performances and lifetime are partly reduced in aqueous systems due to the weakness of the sensitizeroxide bond leading to unstable cells in neutral or basic aqueous media. A new strategy was recently proposed that consists of exploiting a more stable Sn(oxide)–O–Sn–C(alkyl) linkage in order to strongly bind a dye onto nanocrystalline TO or FTO powders [129,131]. FTO powders were prepared by the controlled hydrolysis of the tert-amyloxyfluorodipentan-2,4-dionatotin(IV) complex followed by a thermal treatment in air at 550°C. They consist of aggregated cassiterite, FTO nanoparticles 10–15 nm in size, containing 3 at.% of fluorine. These powders react with perylene-substituted organotrialkynyltins to give coloured dye-sensitized FTO powders (DSFTO), stable in air and at any pH in aqueous solutions. Their photoelectrochemical properties have been investigated by cavity microelectrode technique (CME), which performs electrochemical tests without using any binding additive [131]. Under blue laser illumination through an optical microscope, a 600 mV open-circuit voltage has been obtained with DSFTO, in contact with a 0.2 M NaBr aqueous solution [129–131]. The photocurrent – voltage curves for the same system under various blue light (488.0 nm) intensities are depicted in Fig. 7, showing an appreciable photocurrent due to the oxidation of Br ions. The photocurrent response is practically the same for undoped or F-doped-dye-sensitized tin oxide nanoparticles, the only difference being the presence of a small cathodic current wave for the doped material [129].

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500

400

dark 1mW

I (nA)

300

2 mW 5 mW

200

10 mW

100

0

−100 0.0

0.2

0.4

0.6

0.8

1.0

1.2

E (V/SCE)

Fig. 7. Photoelectrochemical response of a dye-sensitized F-doped tin oxide powder enclosed in a cavity microelectrode in contact with a 0.2 M NaBr aqueous electrolyte under various blue light (488.0 nm) intensities, recorded at a potential scan rate of 10 mVs1. (from Toupance et al. [129], with permission of Elsevier.)

7. OTHER APPLICATIONS A common method of fabricating photovoltaic (PV) modules begins with a substrate of soda-lime glass that has been coated on one surface with a thin FTO layer. The transparent conductive oxide forms the top contact for the Si [132] or CdTe [133–136] based solar cell layers that are, in turn, deposited on the tin oxide. FTO films used as electrodes for amorphous silicon solar cells were found to be sensitive to the reducing atmosphere of hydrogen plasma, with an effect on both transparency and conductivity [137]. The plasma resistance was found to be greatly improved by coating FTO with a thin layer of F-doped ZnO or Nb-doped TiO2 [138]. Recent works on thin film CdS/CdTe solar cells have dealt with the contact between the transparent conducting oxide (TCO) and the CdS underlayer. FTO was overtaken by other TCOs, the best efficiency being obtained with an F-doped In2O3 layer [136]. Electrochemical corrosion of FTO induced by water vapour in Si or CdTe thin-film PV modules has been recently reported [139]. PV cells have been built using the Schottky-type junction formed between FTO and n-Si [140,141] in a planar configuration or by using photoelectrochemically etched macroporous silicon [141,142].

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Another recently proposed application of FTO coatings with a pre-adjusted electrical conductance was to passivate electrical circuits in electronic devices, particularly against corrosion and, if necessary, as protection against pyroelectric short circuits [143]. In the same lines, an important industrial aspect was the manufacture of electrically conductive pigments based on F-doped tin dioxide materials [144–146]. Various preparation techniques were described leading to FTO powders with small particle size ( 200 nm) and electrical resistivity in the range 50–1000 Ω cm. These conductive powders are suitable as pigments or fillers for plastics, adhesives, paints, printing inks, dyes, lacquers, paper and textile fibres, as charge control additives for toners, and finally as electrophotographic photoreceptor [147]. The last applications to be mentioned concern the field of electrical engineering. FTO has been considered as an attempt to decrease the weight of plates in lead batteries [148]. Some beneficial effects have been obtained on discharge capacity and cycling life by mixing the active mass with a few wt% FTO powders [149]. FTO was also used to protect the titanium grid supporting the PbO2 electrode against electrolyte aggression [150]. Since the plate incorporates a material that is self-passivating under the electrical potential and highly acidic conditions found in the Pb-acid battery, any pinholes, gaps or flaws in the Pb coatings are naturally resealed. Joule effect in FTO thin films can serve to construct resistive heaters. When plugged into the mains, transparent domestic electric radiators have been designed. A recent laboratory application was to investigate thermally activated water scaling by in situ electrogravimetry. For that, piezoelectric quartz were equipped with a FTO heating element for promoting calcium carbonate scaling onto the heated surface and real-time measurements of the deposited mass [151]. 8. CONCLUSIONS In this chapter, we have tried to point out the particular interest of FTO within the class of transparent conductive materials, as attested by the variety of its applications and the number of industrial patents. FTO can be routinely synthesized by a number of well-established techniques, as spray pyrolysis and CVD, yielding well-defined physical characteristics. Fluorine incorporation is very efficient for improving the tin oxide conductivity, but in most cases the amount of incorporated fluorine atoms is larger than the quantity required for doping. The presence of fluorine is found to be correlated with that of planar defects, identified as cassiterite twins, which tend to limit the free electron mobility. FTO can be produced either as polycrystalline thin films or as nanosized powders. Within the present state of the art, FTO needs a relatively high crystallization temperature, excluding for the moment to deposit highly conductive films onto plastic substrates. In the near future, we can expect some progress in this field by new developments in thin film synthesis methods, for instance by plasma-assisted CVD or by pulse laser deposition.

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ACKNOWLEDGEMENTS The author thanks Dr Michel Froment for helpful discussions and comments. REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21]

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Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 23

Doped transparent conducting oxides suitable for the fabrication of high efficiency thin film solar cells A. Bosioa, N. Romeoa, and V. Canevarib a

Department of Physics, University of Parma, 43100 Parma, Italy

b

IMEM-CNR Institute, 43100 Parma, Italy

1. INTRODUCTION In recent years there has been a great interest in metallic oxides thin films due to their various industrial applications [1–3]. Thin films of these materials are produced by several techniques and are called transparent conducting oxides (TCOs). The most studied and popular TCOs are: SnO2:F(FTO), ZnO:Al(AZO), In2O3:Sn(ITO) and Cd2SnO4(CTO) [4]. These metallic oxides exhibit very good optical transparency  90% for the visible light and near-infrared radiation, and very high n-type conductivity. For these reasons TCOs are generally near to be degenerated semiconducting materials with free carrier concentrations between 1018 and 1020 cm3. The high transparency and also the high electrical conductivity make the TCOs suitable for a great variety of applications. In fact, they are used in optoelectronic devices and as transparent electrode in photovoltaic modules. Also, they have been employed in glass coatings, for example, as transparent heating elements for planes and car windows. Furthermore, they could be used as transparent heat-mirror coatings for buildings, cars and energy-saving light bulbs due to their high reflectivity in the IR part of the spectrum. Since it is not possible to obtain both the high electrical conductivity and optical transparency in any intrinsic material, one way to reach acheive this is to create electron degeneracy in a wide bandgap oxide. This could be made in two different ways: (a) Introducing donor elements into the oxide matrix; (b) Exploiting the deviation from correct stoichiometry like the structural defects and/or oxygen vacancies. Point (a) is explained by considering that the substitution of a higher valence cation by a donor impurity in the oxide, e.g. tin or antimony in indium oxide or fluorine in tin oxide, increases the electron concentration and hence the n-type

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conductivity. On the contrary, the replacement of a lower valence cation by an acceptor impurity generates a hole (broken bond) that works like a trap (deep level in the energy gap) in the n-type metallic oxide, decreasing its n-conductivity. Since the mean grain size of TCOs thin films are in the range 100–1000 Å depending on the deposition method, the high electrical conductivity of doped and undoped films depends only on carrier (electron) concentration and not on their mobility. This is due to the fact that the mobility in these films is considerably lower than that in the bulk materials, because it is limited by grain boundaries. In the past years, a number of new TCOs have been developed starting from multicomponent oxides such as GaInO3, ZnSnO3, Cd2Sb2O6:Y, Zn2SnO4, MgIn2O4 and In4SnO12 [5]. All these metallic oxides reach high n-type conductivity following the behavior described above. In addition, the new p-type TCOs have been intensively studied in recent years in order to make a p–n junction. In 1997, it was reported for the first time that CuAlO2 thin film exhibits p-type conductivity. Later, a new series of materials based on Cu was discovered such as CuGaO2 and SrCu2O2 [6,7]. In 2000, UV-emitting diode based on p–n heterojunction composed of p-SrCu2O2 and n-ZnO was successfully fabricated by heteroepitaxial thin film growth. The major area of interest, however, is in n-type TCOs due to their utilization in industrial applications. One of these applications is in photovoltaic (PV) modules fabrication, where it is necessary to seek very low resistivity. This direction is strongly accelerated by rapidly rising demand for enlargement of module size. In PV module production, the role of TCOs does not concern only the very high electrical conductivity and very high optical transparency, but also the chemical and physical stabilities with respect to all the materials constituting the device. In fact, one of the most promising materials suitable for large-scale PV modules production is CdTe. Solar cells based on CdTe technology have been studied for several years and they seem to be ready for use in industrial production [8]. This PV device is commonly fabricated in the front-wall configuration (Fig. 1). This means that starting from the soda-lime glass it is necessary to cover it by a TCO thin film, which is the front contact. On top of the TCO film are deposited the active layers and, finally, the backcontact. The active layers are CdS and CdTe films. They form the p–n junction where the PV effect-takes place. Since CdS film is always n-type, as a consequence the CdTe film must be p-type. In efficient solar cells, CdS films are prepared by RF sputtering, closespaced sublimation (CSS) and chemical-bath deposition (CBD); and CdTe films are deposited by CSS [9,10], electrodeposition and metal–organic chemical vapor deposition (MOCVD). Each technique has its own merits but in particular, in order to realize a device easily scalable for industrial needs (i.e. the costs of the equipment, the deposition rate and the final cell performance), the CSS and

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537

Light Soda Lime Glass Active layers

{

TCO (Sputt.)

Front contact

CdS (Sputt.) CdTe (CSS) Sb2Te3 (Sputt.) Mo (Sputt.)

}

Back contact

Fig. 1. Schematic of the CdTe/CdS solar cell structure.

sputtering techniques seem to be best suited. Despite these remarkable advantages, the preparation of thin film solar cells based on CdTe/CdS heterojunction still exhibits quite a few open problems and it is, therefore, subjected to a great margin of progress. One of the open questions, resolved in our laboratory, is certainly the back contact, which is crucial for the temporal stability of the solar cell. In fact, in order to realize a low resistance, possibly ohmic contact with p-type CdTe film, various metals like Cu, Hg, Pb, Ag or Au, which due to their ability of diffusing into the different layers may deteriorate the device, are made use of. We solved the backcontact problem by depositing a thin layer of Sb2Te3, which is a stable compound that exhibits a forbidden energy gap of 0.3 eV and is a degenerate p-type semiconducting material with a resistivity of 104 Ω cm [11]. The Sb2Te3 layer is deposited by sputtering at a substrate temperature of 300–350°C. This high substrate temperature allows the formation of a p Sb2Te3 layer on top of CdTe film, which assures, together with its low resistivity, the ohmic behavior of the backcontact. Despite this, the efficiency of the solar cells are in the range 10–14% because the I–V characteristics are affected by a series resistance. As a consequence, the fill factor is quite low (55–65%). We demonstrated that this problem in the I–V characteristic cannot be attributed to a rectifying or non-ohmic backcontact, but this strongly depends on the way the front contact is made. Having resolved the problem of the stability of the backcontact, we found out that instability in the cell can rise from the front contact, namely the TCO-CdS. In order to solve this problem we studied several types of TCOs from the point of view of the final behavior of the device. 2. TRANSPARENT CONDUCTING OXIDE LAYER (TCO) 2.1. Interaction with CdS

All the TCO thin films are deposited on soda-lime glass by RF magnetron sputtering at 500°C substrate temperature. The deposition rate is typically in the

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range 5–10 Å/s, the average power density is in the range 8–12 W/cm2 and the argon pressure is 103 mbar. All the different TCO’s sputtering targets were from Cerac. They are 3 in. in diameter prepared by the hot-pressing technique. All the TCO thin films deposited under these conditions have been characterized by measuring their electrical conductivity and optical absorbance. The TCOs films with the right resistivity and transparency were used to fabricate CdTe/CdS solar cells. The structure of this device has been described above and depicted in Fig. 1. The device is in the front-wall configuration, i.e. films of CdS, p-type CdTe and an ohmic backcontact are subsequently deposited onto a transparent conducting oxide coated soda-lime glass. The backcontact on CdTe film is made by depositing 1500 Å of Sb2Te3 and 1500 Å of Mo films in a sequence at a substrate temperature of 300–350°C. Both the materials are deposited by RF sputtering with a deposition rate greater than 10 Å/sec. CdTe films are deposited by CSS at substrate temperatures in the range 480–540°C, while the CdTe source is kept at about 650°C. During the deposition, the argon partial pressure is 1 mbar and the distance between the source and the substrate is around 4 mm. The typical deposition rate is 2 μm/min and the time required for the film preparation is 3–4 min [12]. CdS films are prepared with a typical deposition rate of 10 Å/sec by means of a partial argon pressure in the sputtering chamber of 103 mbar. The substrate temperature is 300°C and the final thickness of the CdS films is 1000 Å. During our investigations we found that independent of the type of transparent conducting oxide, the system formed by soda-lime glass, covered by TCO and CdS film, needs to be annealed in air in order to obtain high efficient devices. The annealing temperature and time are, respectively, about 500–520°C and 20 min. During annealing in air, there is an interaction among the soda-lime glass, the TCO and CdS films, which is seen by the presence of a thin layer of a new material on top of CdS, i.e. the reaction product. This material is soluble in hot water and must be removed before the deposition of the CdTe film. The results of X-ray photoelectron spectroscopy carried out on this material are shown in Fig. 2. The oxygen 1s signals at binding energies of 532.5, 530 and 529 eV correspond to CdSO4, In2O3 and CdO, respectively. From these results we deduce that the material, which is formed on top of CdS is constituted by two layers, namely CdSO4 with a thickness in the range within 20–60 Å and CdO with a thickness in the range 300–500 Å. These materials are soluble in hot water and have to be removed before depositing CdTe. Presumably, the thickness of CdSO4 and CdO layers depends on both the annealing temperature and time. The amount of these layers also depends on the presence of Na, which comes from the soda-lime glass, and the thickness of both TCO and CdS films. The annealing of CdSO4 and CdO layers in air, at the same temperature of a CdS film deposited onto alkali-free glass, did not show the formation of any material on top of the CdS film, indicating that TCO and Na are

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Intensity (counts × 1000)

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required for their formation. After annealing in air, we observed the formation of these layers, with different thicknesses, on top of the CdS film with all the TCOs checked in our laboratory. From this point of view, the choice of the TCO is very important in the preparation of the CdTe/CdS solar cells not only for its electrooptical characteristics, but also, for its interaction with CdS. 2.2. In2O3 (IO) family

In2O3 (IO) films are normally polycrystalline with cubic structure and a typical grain size of about 100–500 Å depending on the deposition technique used. The most commonly used techniques for the deposition of IO thin films are reactive RF sputtering, chemical vapor deposition, spray pyrolysis, glow discharge and activated reactive evaporation. IO films often exhibit superior electrical and optical properties with respect to the other transparent conductors; this fact is inprinciple due to the higher mobility in IO. In fact, In2O3 films prepared with various methods of deposition have mobilities in the range 10–75 cm2/V/s, which is very high considering that IO films are polycrystalline thin films. IO films exhibit a direct optical bandgap, which lies between 3.55 and 3.75 eV and has been shown to increase with an increase in carrier concentration owing to the Burstein – Moss shift. For films with thickness under 1 μm, the optical transparency in visible and near IR regions is about 80%–90%. For this reason, IO is a more widely used transparent conductor. 2.2.1. Tin doped In2O3 (ITO) films

ITO films prepared in our laboratory were obtained by RF sputtering from different targets with different stoichiometries. In particular we studied four kinds of ITO: In2O3 containing, respectively, 1, 2, 4 and 10% weight of SnO2. Some ITO films were deposited in a sputtering gas mixture containing Ar  O2; the O2 partial

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pressure was varied between 2 and 20% with respect to total of Ar  O2 pressure. All these films exhibit a very low resistivity of 2  104 Ω cm; the resistivity of the ITO films does not depend on the stoichiometry of the target and sputtering deposition parameters such as the substrate temperature, deposition rate, Ar pressure and power density. This probably means that Sn is an effective dopant and the doping level does not depend on the deposition parameters. Since the optical transparency of the ITO films is more than 80%, the films are transparent enough to be used as front contact for CdTe/CdS solar cells. However, we found out that the efficiency of the solar cells, by using an ITO film thicker than 1.2 μm, could be quite high (above 14%), but was not much reproducible. We explain this considering that all the ITO targets modify their surface after several runs forming some In-rich nodules, which can cause some occasional discharges during the sputtering deposition [13]. This discharge instability produces a non-stoichiometrically uniform ITO film. In order to improve the reproducibility of the ITO film we introduce a small amount of H2 and trifluoromethane (CHF3) during the film deposition in the sputtering chamber. The H2 partial pressure was changed in the range 1–10% and the CHF3 in the range 1–10% with respect to the total Ar  H2  CHF3 pressure. With the maximum quantity of CHF3, we did not observe any indium-rich nodule formation on the surface of the ITO target. This is probably due to the fact that the free indium atoms can react with fluorine on the surface of the target forming a stable compound. Therefore, this reaction prevents the segregation of In into superficial nodules and, as a consequence, the sputtering discharge is more stable producing stoichiometric and uniform ITO films. From the point of view of the stability of the whole device, the quality of the ITO layers is better for films prepared in Ar  4% H2  10% CHF3 gas. In fact, after the annealing the soda-lime glass/ITO/CdS system in air at a temperature above 520°C, only a small amount of CdSO4 and CdO was observed, while, under the same annealing conditions, ITO films deposited in the Ar  O2 gas mixture presented a very thick CdSO4/CdO layer on top of CdS film. The amount of this reaction product depends also on the thickness of the TCO film, and is less for the thicker ITO films. This fact is probably due to a better shield against Na diffusion from the soda-lime glass. 2.2.2. Germanium-doped indium oxide (IGO) films

Germanium-doped indium oxide (IGO) films were prepared by RF sputtering from a target of In2O3 containing 4% weight of GeO2 and the typical deposition rate was larger than 10 Å/sec. The surface of this target did not present any formation of In-rich nodules and the film deposition was much stable and reproducible. The preparation of IGO film was carried out by using pure argon or a mixture of Ar, H2 and CHF3 as sputtering reactive gas. In the first case, we obtained an IGO film with a resistivity of the order of 8.5  104 Ω cm and a very good transparency in the visible region of the spectrum. This implies that

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germanium has the same behavior as tin in a In2O3 matrix, and seems to be an effective dopant for In2O3. The IGO film deposited by reactive sputtering by using a mixture of Ar  H2  CHF3 as sputtering gas presented a resistivity of 2  104 Ω cm (Fig. 3). This result showed that hydrogen could create some oxygen vacancies into the In2O3 matrix giving more possibility for the fluorine to substitute oxygen and, therefore, can dope the IGO film better during the deposition. The H2 and CHF3 gas pressures were both 5% with respect to the total Ar  H2  CHF3 gas pressure. The IGO films showed a good stability with regard to annealing of the TCO/CdS system in air at a temperature of 520°C since we observed a very small amount of CdSO4/CdO on top of CdS film after annealing in air. The IGO film stability and its very low resistivity allowed us to use, as a TCO layer, only a 4000- Å-thick film. Using this TCO we fabricated solar cells with efficiency up to 14% with a very good reproducibility. 2.2.3. Fluorine-doped indium oxide (IFO) films

The last TCO of the IO family studied in our laboratory was fluorine-doped indium oxide layer (IFO). In2O3 can be sputtered at a relatively high deposition rate ( 10 Å/s) without any change in the target surface. When In2O3 is not intentionally doped during the sputtering deposition, it grows with a resistivity of the order of 1  102 Ω cm. In this case, the conductivity is due to native defects, such as oxygen vacancies (Fig. 4a). We were able to prepare IFO films with a resistivity of ⬇ 2  104 Ω cm by introducing Ar containing 5% of H2 and 5% of CHF3 in the chamber during the sputtering deposition (Fig. 4b). The IFO films obtained by this process are very smooth and transparent. Besides, we found out that 1000 Å of this

Fig. 3. Variation of the IGO:F film resistivity as a function of the percentage of the trifluoromethane (CHF3) with respect to the total Ar  CHF3 pressure used in the sputtering chamber during the film deposition. The hydrogen was fixed at 2% with respect to the total Ar  H2 pressure and the substrate temperature was kept at 450°C for all the deposited IGO films.

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Fig. 4. (a) Variation of the INO film resistivity as a function of the percentage of H2 with respect to the total Ar  H2 pressure used in the sputtering chamber during the film deposition. (b) Variation of the INO:F (IFO) film resistivity as a function of the percentage of CHF3 with respect to the total Ar  CHF3  H2 pressure used in the sputtering. The substrate temperature was kept at 450°C for all the deposited films.

material are sufficient to passivate sodium atoms, which can diffuse into the film from the soda-lime glass. For this reason and for its intrinsic stability, IFO is perhaps the best-suited material for the production of solar cells. In fact, we always observed a very small amount of CdSO4/CdO when we annealed this TCO covered by CdS at 520°C in air. Depositing, in sequence, 4000 Å of IFO and 1000 Å of CdS onto a 1 in2 soda-lime glass, after the usual stabilization in air of the system, we obtained a CdTe/CdS solar cell with 14% efficiency. 2.3. Fluorine-doped tin oxide (FTO) films

SnO2 (TO) is the first transparent oxide to have received relevant commercialization. Under an optimum condition of deposition, undoped TO films are generally polycrystalline with a tetragonal structure.

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The n-type conductivity is primarily due to its non-stoichiometry. Often there is a mixture of two phases: SnO and SnO2. Indicating by x the fraction of SnO in the mixture, the Sn1x4 Snx2O2x2 oxide could be formed. This oxide contains x oxygen vacancies and xSn2 atoms that are free to donate 2x electrons for conduction. Pure TO thin films, prepared by sputtering in Ar atmosphere in our laboratory, starting from an oxide target, exhibit a resistivity of 11 Ω cm and an optical transparency of 90% in the region of interest of the light spectrum. The films were doped by mixing the Ar sputtering gas with CHF3 in the range 1–10% in regard to the total Ar  CHF3 gas pressure. The minimum resistivity that we were able to obtain by making use of the maximum quantity of CHF3 in the sputtering chamber during the SnO2 deposition is about 8  104 Ω cm. A disadvantage of pure IO thin films in comparison with the other IO family TCOs is the resistivity of the SnO2:F, which is 3–4 times greater. This implies that in order to obtain the same sheet resistance as that of the IO family TCOs films, we are forced to deposit a SnO2:F film, 3–4 times thicker, losing the transparency. As a consequence we did not prepare any solar cell with SnO2:F alone, but we used this material as a buffer layer against Na diffusion from soda-lime glass. For example, we deposited a SnO2:F film, 1000–5000 Å thick, on top of an ITO film prepared as described above. The final performance of the solar cell did not change in comparison with the device fabricated by making use of ITO alone. 2.4. Fluorine-doped zinc Oxide (FZO)

ZnO thin films have been prepared by a variety of thin film deposition techniques such as reactive DC and RF magnetron sputtering, chemical vapor deposition (CVD), reactive thermal evaporation and activated reactive evaporation (ARE). The sputtered films are normally polycrystalline with an average grain size of about 50–300 Å and a wurzite-type structure with a strong c-axis preferred orientation perpendicular to the substrate. Pure and stoichiometric zinc oxide films show a very high resistivity and a direct bandgap of about 3.2 eV. Furthermore, ZnO has one of the largest electromechanical coupling coefficient. Due to this property, zinc oxide is a well-known piezoelectric material, which has been used as a transducer for surface acoustic wave device. In the last 10 years, ZnO has been intensively used as TCO for PV applications and as a gas sensor device. In the first case, a ZnO film with high transparency in the visible region and high conductivity is needed. In order to reach a good conductivity, ZnO is often doped with trivalent cations such as indium and aluminum. In our laboratory, we obtained ZnO:Al (AZO) thin film, deposited by RF sputtering, with a resistivity of the order of 8  104 Ω cm and a very good transparency, over 90% in the range 4500–8500 Å of the light spectrum. The starting target was a hot-pressed powder mixture of 98% ZnO and 2% Al2O3, supplied by Cerac Inc. Since the ZnO:Al thin films are not stable at high temperature with regard to the Al diffusion, as we observed by making use of

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Fig. 5. Variation of the ZnO:F film resistivity as a function of the percentage of the trifluoromethane (CHF3) with respect to the total Ar  CHF3 pressure used in the sputtering chamber during the deposition. The hydrogen was fixed at 5% with respect to the total Ar  H2 pressure and the substrate temperature was kept at 300°C for all the deposited ZnO films.

these films in our CdTe/CdS solar cell, we tried to dope ZnO films with fluorine. For this purpose, we used reactive RF sputtering technique starting from a pure Cerac ZnO target and introduced it into the sputtering chamber a gas mixture containing Ar, H2 and CHF3. The H2 partial pressure was fixed to 5% with respect to the total Ar  H2 pressure, while the CHF3 pressure varied in the range 1–10% of the total Ar  CHF3 pressure in order to change the fluorine doping level in the ZnO:F film. The presence of H2 in the sputtering chamber assures the creation of some oxygen vacancies into the ZnO matrix, giving more possibility to fluorine to substitute oxygen atoms and, as a consequence, to dope ZnO film better during its deposition. We observed that the presence of hydrogen into the sputtering chamber made the deposition of ZnO films almost independent from the substrate temperature in the range 200–350°C. In fact, low resistivity of the order of 8  104 Ω cm and high transparency has been achieved by making use of the maximum quantity of CHF3, independent from the substrate temperature in the above-mentioned range (Fig. 5). We also demonstrated that this doping level did not change with heat treatment in air at a temperature higher than 500°C. This means that fluorine substitutes oxygen makes a stable chemical bond with Zn and so it is an effective dopant for ZnO. 3. CdS LAYERS The CdS film in the CdTe/CdS solar cell is the so-called window layer. It allows, being n-type, the formation of the p–n junction with p-type CdTe. With its energy

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gap of 2.42 eV, CdS results to be transparent in the visible part of the solar light spectrum. Therefore solar light can penetrate into the CdTe layer, where the photovoltaic effect takes place. In order to have a high efficiency solar cell, we need an excellent p–n junction and a very good contact on the p-type CdTe. In our laboratory, we have solved the second problem regarding deposition by sputtering a Sb2Te3 thin film on top of the CdTe layer. Since the p–n junction, which is formed between n-type CdS and p-type CdTe, is strongly dependent on the proper interaction between these two layers, it in the deposition technique used to prepare these materials becomes very important. In particular, for CdS, the best-suited deposition techniques are RF sputtering, close-spaced sublimation (CSS) and chemical bath deposition (CBD). Although the highest efficiency was obtained by using a CdS layer prepared by CBD, we prefer to use the sputtering method of deposition, since CBD is not suitable for large-scale production. The CdS layers were grown by sputtering at a substrate temperature of about 220°C, with a deposition rate up to 10 Å/sec in an argon atmosphere of 103 mbar. After the deposition, the soda-lime glass covered by the TCO-CdS bilayer was heat-treated in a vacuum chamber at 500°C for 30 min. Later, the CdTe/CdS was completed in the usual way, which is described in Ref. [12]. The efficiencies of solar cells prepared with this type of CdS layer are quite poor, i.e. they are in the range 8–10%. This result is due to the fact that these devices have a very high diode reverse saturation current in order to give high-efficiency devices. One possible explanation of this is that the interaction between CdS and CdTe is not correct and, as a consequence, the grain boundaries in the CdS film are active to canalize the diode reverse current. We solved this problem by introducing argon containing 3% of CHF3 into the sputtering chamber during the CdS deposition. This gas is decomposed and ionized in the sputtering discharge freeing F ions that, being strongly electronegative, are directed to the substrate that is the positive electrode and two different events can occur: (i) The presence of energetic F ions near the substrate favors the formation of fluorine compound such as CdF2 during the growth of the CdS film [14]. (ii) The F ions, accelerated by the electric field present in the discharge, hit the film surface during the deposition with energy sufficient to sputter back the Cd or S atoms that are not well bonded. This effect leaves a high-quality CdS film, concerning both the optical and the structural properties. We can see in Fig. 6 that the CdS film deposited in argon  CHF3 has an energy gap higher than that of the film deposited in argon alone. The role of fluorine has been stated as an ineffective dopant in CdS as we did not observe any change in the resistivity of the CdS layers. CdTe/CdS solar cells fabricated by using 80-nm-thick CdS films deposited by sputtering in presence of

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Fig. 6. Transmission spectrum of the sputtered CdS layer, deposited in pure argon (a) and deposited in argon  CHF3 (b). The shift toward shorter wavelengths of the absorption edge proves the beneficial effects of deposition in the presence of CHF3.

Fig .7. J vs. V characteristics of the CdTe/CdS solar cells taken in standard conditions in our laboratory.(a) Behavior of a device obtained by using a CdS layer deposited by sputtering in pure argon and then annealed in vacuum at 500°C substrate temperature for 30 min. Typical photovoltaic parameters are: Jsc25.3 mA/cm2, Voc738 mV, ff0.53 and efficiency is about 10%.(b) Behavior of a device obtained by using a CdS layer deposited by sputtering in argon  CHF3. Typical photovoltaic parameters are Jsc25.8 mA/cm2, Voc848 mV, ff0.72 and efficiency is about 15.8%.

fluorine showed very high efficiency in the order of 15–15.8%. We explain this fact by considering that CdS(F) may contain CdF2 segregated in the grain boundaries. While the CdF2 segregated in the grain boundaries can be useful to passivate them, the CdF2 layer grown on the CdS surface may adjust the interaction among CdS and CdTe during the CdTe deposition by CSS. These remarkable results are shown in Fig. 7, which demonstrate that CdS(F) can be used as

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deposited without any other treatment. Probably, this fact makes the CdS sputtering deposition in the presence of fluorine the best candidate for the industrial production of PV modules based on the CdTe/CdS technology. 4. CONCLUSIONS The cells fabricated without antireflecting coating, with an area of about 1 cm2 by making use of 1 in2 soda-lime glass covered with different types of TCOs previously described, exhibit a very good device efficiency of up to 15.8%. The J vs. V characteristics were tested in the dark and under illumination and the photovoltaic parameters were measured under the standard condition of 300 K, 100 mW cm2 and AM1.5 by using a solar light simulator supplied by Oriel Inc. The best results were obtained by using an IFO thin film, 400-nm-thick, as a front-contact covered by a 80-nm-thick CdS:F film as a window layer. In conclusion, we put in the evidence that CdS films, in order to be used in high-efficiency solar cells, need to be made in the presence of a reactive gas containing fluorine-like CHF3. This reactive gas could form an insulating material, namely CdF2, that can passivate the CdS grain boundaries and adjust the proper interaction of the CdS/CdTe system during the CSS deposition of CdTe. With this kind of CdS films, we are able to obtain solar cells that improve their efficiency by making a light-soaking of 20 h under 10 suns at 110°C or of 20 min under 20 suns at 280°C (in the open-circuit conditions). In several cases the efficiency is higher after ageing. We believe that this stability high solar cells is due to: (a) the use of Sb2Te3 as backcontact, which does not contain any copper or any doping elements that can diffuse into CdTe. (b) the use of a stable TCO, namely a fluorine-doped In2O3. (c) the use of a CdS layer with possible grain boundaries passivation due to CdF2. This is a very satisfactory result, especially with regard to the reproducibility and time stability of solar cell, which makes the solar cell technology mature for large-scale production. REFERENCES [1] T.J. Coutts, D.L. Young, X. Li, W.P. Mulligan, and W. Wu, J. Vac. Sci. Technol. A Vac. Surf. Films, 18 (2000) 2646. [2] M.H. Sohn, D. Kim, S.J. Kim, N.W. Paik, and S. Gupta, J. Vac. Sci. Technol. A, 21 (2003) 1347. [3] N. Biyikli, T. Kartgoglu, O. Artur, I. Kimukin, and E. Ozbay, Appl. Phys. Lett., 79 (2001) 2838. [4] K.L. Chopra, S. Major, and D.K. Pandya, Thin Solid Films, 102 (1983) 1. [5] T. Minami, Y. Takeda, S. Takata, and T. Kakumu, Thin Solid Films, 308–309 (1997) 13. [6] K. Ueda, T. Hase, H. Yanagi, H. Kawazoe, H. Hosono, H. Ohta, M. Orita, and M.J. Hirano, Appl. Phys., 89 (2001) 1790.

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[10] [11] [12] [13] [14]

A. Bosio et al. A. Kudo, H. Hanagi, H. Hosono, and H. Kawazoe,Appl. Phys. Lett., 73 (1998) 220. N. Romeo, A. Bosio, A. Romeo, M. Bianucci, L. Bonci, and C. Lenti, Proc. Int. Conf., PV in Europe-From PV Technology to Energy Solutions, Rome, Italy, 7–11 October 2002, p. 433. X. Wu, J.C. Keane, C. DeHart, R.G. Dhere, D.S. Albin, A. Duda, and T.A. Gessert, Proc. 17th Eur. Photovoltaic Sol. Energy Conf., Vol. II, 22–26 October 2001, Munich, Germany, pp. 995–1000. C.S. Ferekides, D. Marinskiy, V. Viswanathan, B. Tetali, V. Palekis, P. Selvaraj, and D.L. Morel, Thin Solid Films, 361–362 (2000) 520. N. Romeo, A. Bosio, R. Tedeschi, and V. Canevari, Thin Solid Films, 361–362 (2000) 327. N. Romeo, A. Bosio, R. Tedeschi, and V. Canevari, Mater. Chem. Phys., 66 (2000) 201. P. Lippens, A. Segers, J. Haemers, and R. De Gryse, Thin Solid Films, 317 (1998) 405. N. Romeo, A. Bosio, and V. Canevari, Proc. 3rd World Conf. Photovoltaic Energy Conversion, 11–18, May 2003, Osaka, Japan, Vol. I, pp. 469–470.

Fluorinated Materials for Energy Conversion T Nakajima and H Groult (Editors) © 2005 Elsevier Ltd. All rights reserved.

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Chapter 24

Fluoride technologies application within the Molten-Salt Reactors fuel cycle Jan Uhlir Nuclear Research Institute Rez plc, CZ-250 68 Rez, Czech Republic This chapter deals with fluoride technologies that are now under technological development and not used so far in nuclear industry, with the exception of the production of thorium tetrafluoride. However, the utilization of these fluoride technologies is considered to be very prospective in connection with the development of future advanced nuclear reactor types, which are likely to be deployed mainly in the second half of this century. The biggest chance for the fulfillment of fluoride technologies discussed hereafter is within the molten-salt reactor fuel cycle. 1. MOLTEN-SALT REACTORS Molten-salt reactors (MSRs) represent one of the promising high-temperature nuclear reactor types for future generation of electricity and of heat for hydrogen production. They could also be used as transmuters to burn plutonium and other transuranium elements occurring in the spent nuclear fuel of nowadays-existing nuclear reactor types [1]. This seems to be of great importance because the spent fuel management strategy represents one of the most serious problems, which should be appropriately solved for further sustainable development of nuclear power. Therefore, a new generation of nuclear reactors, planned for deployment mainly in the second half of this century, have, in addition to a high-temperature character, the property to minimize production of nuclear waste; moreover, some of them could be used for transmutation of the most dangerous long-lived radioactive isotopes included in the spent nuclear fuel. MSRs are usually characterized as non-classical nuclear reactor types due to a specific character of their fuel, which is liquid-constituted by a molten fluoride salt mixture circulating between a reactor core and a heat exchanger. The fission materials (uranium and transuranium elements) are dissolved in carrier molten salt, which is also a heat transferring agent. MSRs also exhibit the high assumptions for the applications as power generating transmuters. The typical fuel of the MSR working as the

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nuclear transmuter is a mixture of fluorides of plutonium and other transuranium elements (Np, Am, Cm), usually called as minor actinides dissolved in the carrier fluoride salt. The other very promising mode of MSR operation is based on the use of 232Th–233U fuel cycle with minimized production of long-lived nuclear waste with comparison with the 238U–239Pu fuel cycle currently used in present reactor types. In this mode, the MSR works as a reactor breeder producing its own fissile material 233U from fertile 232Th. Essentially the main advantages of MSRs emerge from the prerequisite that these reactor types should directly be connected with the “on-line” reprocessing of circulating liquid (molten salt) fuel. This fuel cleanup is necessary within a long run to keep the reactor in operation. As a matter of principle, it permits to clear away typical reactor poisons like xenon, krypton, lanthanides, etc. and also the products of burned plutonium and transmuted minor actinides. On the other hand, the technologies of liquid transuranium molten-salt fuel processing from the current spent fuel and the online reprocessing of MSR fuel represent two killing points of the whole MSR technology, which have to be successfully solved before MSRs deployment in the future. The history of MSR technology dates back to the end of the 1940s when Ed Betis and Ray Bryant from Oak Ridge National Laboratory (ORNL) began to study the possibility of using nuclear reactors with liquid-fuel based on molten fluoride salts. Later on, in the 1950s, the Aircraft Nuclear Propulsion Program (ANP) was carried out in ORNL. Originally, the program was intended to develop a strategic bomber propulsion. The first small MSR was realized in 1953. It was in operation only for a short time; however, in 1954, the second smaller MSR that reached the power of 2.5 MWt was capable of demonstrating the MSR technology more successfully. Its primary fuel circuit was cooled by helium gas and the circulating fuel comprised NaF–ZrF4–UF4 mixture with molar composition of 53:41:6. The fissile material was uranium-235. The reactor was operated about 100 MWh and with the maximal operating temperature of circulating fuel being 882°C. At the end of 1950s, after the launching of first space rockets, the main motivation of using the MSRs for military purposes ceased. However, based on the promising results obtained during ANP, the development of the technology was consecutively oriented toward peaceful ends and the work continued during the programs called molten salt reactor experiment (MSRE) in the 1960s and molten salt breeder reactor (MSBR) in the end of the 1960s and beginning of the 1970s. In those days, the main effort was focused to the demonstration of viability of the MSR and to the verification of the reactor’s operation. In the frame of MSRE, the experimental MSR was realized and operated during the period 1965–1969. Initially, the fuel for MSRE was constituted by 7LiF–BeF2–ZrF4–UF4 and the mixture of LiF–BeF2 served simultaneously as the heat-transferring agent. In 1968, the original 235U was extracted from the fuel salt and later replaced by 233 U (this was the first time 233U had been used as a reactor fuel). Also the

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composition of carrier fuel salt was changed a bit; the new fuel salt was composed only of 7LiF–BeF2–UF4 (the reason of using 7Li isotope was to prevent the undesirable tritium production from 6Li). The moderator of the reactor was graphite; structural material of MSRE was nickel alloy INOR-8 (later called Hastelloy-N); the coolant of secondary circuit was LiF–BeF2 eutectics. The power of MSRE was nearly 8 MWt and the working temperatures here in the range 600°C [2–5]. Also, MSR fuel cycle chemistry was studied in depth during MSRE and the liquid fuels for MSR were processed; however, the MSR spent fuel reprocessing was never fully realized in a pilot scale. Inspite of this, considerable effort was carried out in radiochemical laboratories to develop separation processes for uranium, protactinium and rare earth elements from the carrier molten salt. Also the basic flow-sheet was done during the MSBR program to design the main principles of MSBR spent fuel on-line reprocessing. Despite the good results obtained during MSRE and expectations of future realization of a 1000 MWe demonstration MSBR unit, in the beginning of the 1970s, the US – Atomic Energy Commission decided to stop further development of MSRs and to support hereafter only the development of nuclear reactors with solid fuel. Nowadays, based on the new requirements of sustainable development of nuclear power, MSR technology is under revival of interest in the frame of the development of advanced nuclear reactor types. “Generation IV International Forum” chiefly constituted by most of the several industrialized countries supports this renascent research and development. However, it is necessary to realize that the knowledge and experience of MSR technology is not well proportioned. Whereas the knowledge of the MSR performance is quite comprehensive, the MSR fuel cycle technology, including the “on-line” reprocessing, represents one of the most poorly developed and verified areas. On the other hand, this area is experimentally studied at several workplaces at present. 2. FUEL CYCLE TECHNOLOGIES OF MSRS The fuel cycle technologies of MSRs can be partially different depending on the type of fuel used in the reactor. While the fuel processing for the MSR working as a transmuter (molten-salt transmutation reactor, MSTR) represents several processes and technologies of spent nuclear fuel partitioning, the preparation of liquid fuel for MSR working under the 232Th–233U cycle is different and much more easier. One of the specific category is the on-line reprocessing of MSR circulating fuel salt, which has to be done continuously and is in principle similar for both types of MSRs. Based on the point of view of the emplacement of the fuel processing/reprocessing technologies, they are often, particularly for MSTR, divided as front-end (MSTR fuel processing) and back-end (on-line reprocessing) technologies. There

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is also another fundamental difference between the front-end and back-end fuel cycle technology of the MSR: whereas the main parts of the front-end fuel processing technology could be placed separately from the MSR site, the back-end on-line reprocessing has to be directly connected with the reactor primary fuel circuit [6,7]. The final part of fuel processing technology should be also placed in the reactor site and tightly attached to the on-line reprocessing, because the refilling of the fresh fuel into the reactor has to be carried out in connection with the removal of the burnt out fuel and fission products from the primary fuel circuit. Based on a top-rated principle applicable in chemistry and chemical technology that should also be implemented in the decision-making, chemical technologies chosen as preferable for the MSR fuel cycle are choosen. It seems to be most advisable to keep one chemical form (fluorides) of fuel for the entire fuel stream system in developing a compact fuel cycle of the liquid-fuel reactor system. This principle results in a fundamental requirement for the choice of separation (partitioning) methods that could be applied in the reactor fuel processing and especially in the on-line fuel reprocessing. Based on this principle and on the strict requirement of a radiation resistant technology, predominantly fluoride pyrochemical and pyrometallurgical procedures appear in the foreground of interest for MSR (MSTR) fuel processing and on-line reprocessing. The chemical separation (partitioning) technologies, which are often discussed in the frame of MSR (MSTR) fuel cycle and can also be called “fluoride technologies,” are mainly ● ● ●

Fluoride volatilization techniques. Electrochemical separation processes from molten (fluoride) salt medium. Molten (fluoride) salt/liquid metal reductive extraction.

In addition to these, the production of thorium tetrafluoride for the fresh MSR fuel processing also ranks as a fluoride technology. The positioning of technologies mentioned above within the MSR (MSTR) fuel cycles is evident from Figs. 1 and 2. 2.1. Production of Thorium Tetrafluoride

The preparation of thorium tetrafluorides is one of the intermediate reactions in the production of thorium metal. The final ThF4 used for the subsequent production of thorium metal and in the fuel mixture for MSR, must be anhydrous and oxide-free. There is only one way to produce this ThF4; it is a gas-phase hydrofluorination of thorium oxide with anhydrous HF [8] according to the following exothermic reaction: ThO2 (s)  4HF (g) → ThF4 (s)  2H2O (g)

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Fig. 1. Proposed fuel cycle of MSTR (MA, minor actinides - Np, Am, Cm; FP, fission products, mainly lanthanides).

Fig. 2. Fuel cycle of MSR working under 232Th–233U cycle.

The technology has been mastered well and used industrially in the United States, India and other countries. The process is usually performed in a series of hydrofluorination screw-fed horizontal reactors by a stepwise increase in working temperatures from 260 to 566°C. The exploitable by-product of the thorium tetrafluoride production is aqueous hydrofluoric acid. Although the production of thorium tetrafluoride comes under the MSR fuel cycle only marginally, similarly to the production of uranium tetrafluoride or

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lithium fluoride and beryllium difluoride, the technology is not so widely known and therefore is briefly mentioned. 2.2. Pyroprocesses Based on Fluoride Volatilization Techniques

The separation techniques known as “volatilization” or “fluoride volatility” are typical “dry” processes that involve the fluorination of spent fuel with fluorine gas and subsequent separation of resultant volatile compounds, represented mainly by uranium hexafluoride, from non-volatile fluorides, in the first instance from PuF4. There are two main volatilization techniques: a reaction between the fluorine gas and a fused salt and the reaction between the fluorine gas and solid powdered material [9,10,3]. The volatilization technique based on bubbling of fluorine gas into fused fluoride salt containing elements that form volatile fluorides was studied in ORNL mainly during the MSRE program in 1960s and 1970s. The main objective of the program was the development and verification of a possibility to remove uranium (in the form of volatile UF6) from the molten fluoride salt carrier (based on LiF–BeF2 or LiF–BeF2–ZrF4 mixtures). The processes were quite successfully verified with complete separation of uranium from the salt in 1968 [3]. The volatilization technique based on heterogenous reaction between fluorine gas and powdered spent-fuel oxides is known as “fluoride volatility method” (FVM). This process, which was originally designed for fast breeder reactor fuel reprocessing, was studied mainly in the United States, Russia, France, Czech Republic, Belgium and Japan. The fluorination process was realized either in a fluidized-bed reactor (US, France, Belgium and Japan) or in a flame fluorinator (former Soviet Union and Czechoslovakia) [8,10–16]. Nowadays, this process is under further development only in the Czech Republic (in the frame of MSTR fuel cycle development) and in Japan in colabration with Russia (in the frame of fast reactor fuel cycle development) [17,18]. All fluoride volatilization studies have confirmed high efficiency of uranium recovery and promising possibilities of plutonium recovery. However, the efficiencies of individual minor actinides recovery have not yet been fully verified. 3. FLUORIDE VOLATILITY METHOD FVM is the most important fluoride technology applicable within MSTR fuel cycle. The separation process is based on the specific property of uranium, neptunium and plutonium that they form volatile hexafluorides, whereas most of fission products (lanthanides) and higher transplutonium elements, which are present in irradiated fuel, form non-volatile trifluorides. This property has led to the development of several technological processes based on fluoridation of irradiated fuel either by strong fluorinating agents like BrF3, BrF5, ClF3 or even by pure fluorine gas. Major activities were carried out in, Argonne, Oak Ridge and

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Brookhaven, United States, in Fontenay-aux-Roses, France, in Mol, Belgium, in Dimitrovgrad, former Soviet Union and in Rez, former Czechoslovakia. Original intentions of the development of this dry reprocessing method of spent fuel were motivated in the past by the assumed commercial utilization of fast breeder reactors. Their application in the power industry can be economically efficient only in the case of a closed fuel cycle. However, reprocessing of fast reactor spent fuel brings about a number of specific difficulties in comparison with the reprocessing of spent fuel from thermal reactors. They are caused, for example, by higher burn-up, shorter cooling time resulting in a higher amount of energy released by the fuel, higher concentration and amount of plutonium, different cladding material, presence of metallic sodium, different fission products composition, etc. Therefore, countries that were planning the introduction of fast reactors also attempted to develop suitable methods for reprocessing because the industrial hydrometallurgical PUREX process, employing organic extractants and solvents, was not suited for the fast reactor spent fuel. Hence, the most intensive effort in the development of the FVM was in 1960s and 1970s together with the development of fast breeder reactors. The possible application of the FVM as the principal “front-end” reprocessing technology within the MSTR fuel cycle was first considered in the second half of 1990s, when the technologies of “partitioning and transmutation” in connection to the advanced nuclear reactor systems came to be widely judged as a possible and realistic spent fuel management solution. In spite of this, the FVM did not reach industrial level, and to date, remains no more than in laboratory semi-pilot technological conditions. The partitioning technology based on the FVM dedicated to the front-end of MSTR fuel cycle could be very similar to the technology originally dedicated to the fast breeder reactor fuel reprocessing. The whole process consists of the following main operations: 1. Removal of the cladding material from spent fuel elements. 2. Transformation of the fuel into a powder form of a granulometric composition suitable for the fluorination reaction. 3. Fluorination of the fuel (the purpose of this operation is the separation of the uranium component from plutonium, minor actinide and most of fission products). 4. Purification of the products obtained. The first two steps represent preparatory stages for the FVM itself and they can be realized separately from the FVM. Suitable technology for the removal of the cladding material is melting in high-temperature furnace. The cladding material of oxide fuel is either stainless steel or zircalloy and both can be fully removed. Transformation of the fuel pellets into a powder is possible either by mechanical grinding or by partial oxidation of UO2 into U3O8. This chemical

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process is called voloxidation. Either flowing air or oxygen at 575–650°C is used as an oxidizing agent. The original voloxidation process was developed in ORNL in the US and then further developed by several countries [19]. However, the preparation of uniform powder with the required granulometry for subsequent fluorination by voloxidation technique seems to be quite difficult [16]. Early fluorination techniques of the FVM were fluidized-bed processes, but now the flame fluorination of powderized fuel is considered as the most promising unit operation for future industrial applications. This method of fluorination in the frame of the FVM was first used in 1980s in former Soviet Union and in Czechoslovakia [20]. Further description of main unit operations of the FVM is based on technologies realized in RIAR Dimitrovgrad and NRI Rez, where the FVM is now under further technological development. The scheme of a FVM technological line called FREGAT-2 is shown in Fig. 3. This semi-pilot technology was realized in hot cell of RIAR based on the international cooperation between former USSR and Czechoslovakia [16,21]. Czechoslovak companies made majority of equipment and apparatuses. Spent fuel reprocessing in the FREGAT-2 line emerged from the following principles: ●

Powdered ceramic (oxide) spent fuel is dosed in nitrogen atmosphere by a worm doser into the preheated reactor R11 – fluorinator and at the same time pure fluorine gas is also fed. A reaction takes place in the flame, in course of

Fig. 3. Scheme of FVM technological line FREGAT-2 operated in RIAR Dimitrovgrad in 1986–1988. R11, flame fluorinator; K12,13,14,41, condensers; R31, thermal decomposition reactor; K43, tube radiator (cooler); C15,32,46,51,61,62,64, sorption columns; C44, rectification column; H19, catcher tank.

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which the uranium oxides react to form gaseous UF6, most of the plutonium forms gaseous PuF6 and a great majority of fission products and trivalent minor actinides form solid fluorides that drop to the bottom of the fluorinator; only a small part of the fission products (Nb, Ru, Tc, Mo, I, Sb) reacts with fluorine under gaseous fluorides formation and will accompany uranium and plutonium along with rare gases. Three condensers for collecting volatile fluorides are situated next to the fluorinator. The first condenser, K12, is intended for uptaking of niobium and ruthenium fluorides; its working temperature is about 20°C. Uranium and plutonium fluorides condense together in condensers K13 and K14 working at 60 and 80°C, respectively. It was assumed that a complete fluorination of Pu would not take place. Therefore, secondary fluorination of fuel in the reactor R11 is introduced by circulating fluorine gas in the bottom part of the apparatus by means of a circulating pump. Non-volatile fluorides, constituted mainly of lanthanides and trivalent actinides (americium, curium), accumulate in the bottom part of fluorinator R11 on alumina bed during the fluorination reaction, and after termination of this unit operation are poured out from the reactor. Once the K13 and K14 condensers are filled with uranium and plutonium hexafluorides, they are heated and the evaporated fluorides are introduced into the reactor R31, where thermal decomposition of PuF6 to solid PuF4 takes place, whereas volatile UF6 passes through the reactor. In this way, uranium is separated from plutonium. UF6 is then purified by column rectification to remove mainly molybdenum, iodine and technetium fluorides. Sorption columns C15, C32 and C46, filled by NaF pellets and placed in line always next to condensers, play a role of safety equipments in which volatile fluorides could be trapped.

Flame fluorination reaction of a spent oxide fuel is a basic unit operation of the whole process. The reaction between the fuel powder and pure fluorine gas is spontaneous and highly exothermic. The usual temperature of ignition is above 250°C. Subsequently, the temperature in the flame can reach nearly 1700°C and therefore the walls of the reactor chamber and the whole reactor body have to be immediately cooled intensively. Principal fluorination reaction of main or significant fuel components are the following: uranium UO2 (s)  3F2 (g) → UF6 (g)  O2 (g),

Δr H 0298.15   1062.4 kJ/mol U

U3O8 (s)  9F2 (g) → 3UF6 (g)  4O2 (g),

Δr H 0298.15   955.8 kJ/mol U

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plutonium PuO2 (s)  2F2 (g) → PuF4 (s)  O2 (g),

Δr H 0298.15   722.4 kJ/mol Pu

PuO2 (s)  3F2 (g) → PuF6 (g)  O2 (g),

Δr H 0298.15   693.1 kJ/mol Pu

PuF4 (s)  F2 (g) x PuF6 (g),

Kp  [PuF6]/[F2]

lanthanides 2Ln2O3 (s)  6F2 (g) → 4LnF3 (s)  3O2 (g); minor actinides neptunium: NpO2 (s)  3F2 (g) → NpF6 (g)  O2 (g),

Δr H 0298.15   907.9 kJ/mol Np

NpO2 (s)  2F2 (g) → NpF4 (s)  O2 (g),

Δr H 0298.15   844.7 kJ/mol Np

NpF4 (s)  F2 (g) x NpF6 (g),

Kp  [NpF6]/[F2]

americium and curium: 2Am2O3 (s)  6F2 (g) → 4AmF3 (s)  3O2 (g), 2Cm2O3 (s)  6F2 (g) → 4CmF3 (s)  3O2 (g), The behavior of neptunium during flame fluorination varies that between uranium and plutonium. However, the thermal stability of neptunium hexafluoride is substantially higher than of plutonium hexafluoride. Other chemical properties of neptunium that cause both uranium and plutonium streams in the whole process could be contaminated by neptunium. Satisfactory solution of neptunium separation was not yet been found; however, the current effort in the further development of the FVM offers a chance to solve this problem. Uranium, plutonium and neptunium hexafluorides do not form a liquid phase at atmospheric pressure; their sublimation points are close: 56.5°C for UF6, 55.2°C for NpF6 and 62.2°C for PuF6. A certain possibility to separate neptunium from uranium and plutonium within the FVM technology is by sorption–desorption methods on sodium and magnesium fluorides [10]. Sodium fluoride is commonly used for the decontamination of UF6. Uranium, neptunium and plutonium hexafluorides are completely sorbed on NaF at 100°C. While uranium and neptunium

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hexafluorides can be completely desorbed at 400°C by passing fluorine gas through the bed, desorption of plutonium hexafluorides is impossible and forms complex PuF4.3NaF, which is thermally stable in fluorine gas flow. Partial separation of neptunium from uranium is possible only via irreversible sorption of NpF6 on MgF2 at 100°C. UF6 is not sorbed on MgF2, but sorption of NpF6 on this sorbent proceeds from 60 to 70%. Final purification of uranium hexafluorides from MoF6, TcF6, IF5 and SbF5, which tend to accompany UF6 through the system, could be done by rectification process. Distillation of UF6 is usually done in the temperature range from 75 to 90°C at pressure of about 2 atm in order to have uranium hexafluorides in liquid form. Suitable structural materials for FVM equipment are pure nickel and nickel alloys. The rates of corrosion of nickel by fluorine gas, anhydrous HF and volatile fluorides are acceptable up to 600–650°C [9]. Although pure nickel exhibits a very good corrosion resistance, owing to difficulties by welding of pure nickel material, the use of high nickel content alloys could be often more appropriate for the manufacture of several FVM apparatuses than the use of pure nickel [22]. The FREGAT-2 technological line, which was intended to verify FVM by reprocessing of spent fuel from Russian experimental fast reactor BOR-60, was so far the largest realized FVM technology (Fig. 4). The short-term capacity of flame fluorinator reached nearly 3 kg of spent fuel per hour. However, the fluorination had to be interrupted whenever condensers were fully filled by volatile fluorides. Unfortunately, the experimental program at the FREGAT-2 line was

Fig. 4. UF6–NpF6–PuF6 condensers of the FVM technological line in a laboratory of NRI Rez.

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Table 1 Separation efficiencies of selected spent fuel component using FVM Elements

Separation efficiency (%)

U

95–99.5

Pu

⬃ 98–99.5

Np

⬃ 60–70

Nb,Ru

⬃ 95–99

Am,Cm

Individually inseparable (in non-volatile fluoride stream)

FP forming solid fluorides

Individually inseparable (in non-volatile fluoride stream)

untimely closed in 1989 due to political and economical reasons and was not fully completed. In spite of this, the results permit to evaluate FVM as a suitable and promising technology for MSTR fuel cycle. Anticipated separation efficiencies of selected spent fuel components by using the FVM within the frame of MSTR are given in Table 1. 4. FUSED-SALT VOLATILIZATION PROCESS Fluoride volatilization of fused fluoride salt was used for uranium extraction from the fuel salt of MSRE [3]. This technology was realized at Oak Ridge National Laboratory in United States in 1968 during MSRE program. The goal of the extraction of uranium was the change of original 235U for 233U. Carrier molten fluoride salt, in which uranium was dissolved, was LiF–BeF2 eutectics. Uranium was recovered from the carrier salt as uranium hexafluoride by direct fluorination of the salt by fluorine gas. Special equipment for continual fluorination of the salt was designed and realized [23]. The crucial part of the equipment was the continuous fluorination reactor with frozen wall for corrosion protection. The main structural material of the fluorinator was pure nickel (Fig. 5). The carrier molten salt containing dissolved UF4 flowed into the top of the fluorination reactor and was contacted by a countercurrent stream of fluorine gas, which stripped out the uranium in the form of volatile UF6 according to the following reaction: UF4  F2 → UF6

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Fig. 5. Fused-salt fluorination reactor.

The working temperatures were in the range 500–550°C; the coolant used in the reactor walls was NaK. The off-gas containing uranium hexafluoride, fluorine gas and volatile fluorides of some fission products (mainly CrF4 and CrF5 here) passed through sorption column with NaF pellets heated to 400°C for removing chromium fluorides. UF6 was then trapped at the same sorbent at 100°C. Similar technology was under development in ORNL during ensuing MSBR program in the beginning of 1970s, which aimed mainly to design an on-line reprocessing scheme for MSRs working under 232Th–233U fuel cycle [24].

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The technology of uranium removal by fused salt volatilization is to date a widely discussed method and is considered as one of the possible techniques within the on-line fuel reprocessing of future MSRs with Th fuel. Here, the removal of uranium from the fuel salt of MSRs is a part of the protactinium removal process. Protactinium (Pa-233) forms in the reactor core from thorium (Th-232) by nuclear reactions. The half-life of radioactive decay of 232Pa is 27 days and its only daughter product is 233U. However, the arising protactinium, which is also a mighty neutron poison, has to be immediately removed from the fuel salt circulating in the reactor to pre-empt undesirable nuclear reaction. Suitable technology for protactinium removal from the fluoride molten salt carrier is “molten-salt/liquid metal reductive extraction process” occasionally called also as a “metal transfer process”. Because the fuel salt of MSR contains several percent of uranium in the form of uranium tetrafluoride (UF4), it is proved that a primary uranium removal by fluoride volatilization of fused salt can save the amount of reducing agent 7Li used for subsequent protactinium isolation. The simplified scheme of protactinium isolation process according to the ORNL proposal [3] is shown in Fig. 6. 5. PYROMETALLURGICAL SEPARATION PROCESSES OF MSR FUEL SALT CLEANUP The remaining two pyrometallurgical technologies can be considered to be suitable for on-line reprocessing of MSR spent fuel and cannot be called as typical fluoride technologies, in spite of the use of fluoride molten-salt medium. The first technology is molten-salt/liquid metal reductive extraction; the other covers electrochemical separation processes. A brief description of both technologies is done particularly for a better understanding of the complete MSR fuel on-line reprocessing.

Fig. 6. Simplified scheme of protactinium isolation from MSR fuel.

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During the MSRE and MSBR programs, which were in use in ORNL in 1960s and 1970s, the main process proposed for fluoride fuel salt cleanup was a metal transfer process, which is rather known as Molten-salt/liquid metal reductive extraction. The main purpose of this technology, within the MSR on-line reprocessing, is to remove fission products composed mainly of rare earth elements (lanthanides) from the circulating fuel salt [3,24]. Extraction by the liquid bismuth was proposed due to its suitable properties. Bismuth has a low melting point (271°C), a negligible vapor pressure in the temperature of interest (500–700°C) and good solubilities of lithium, thorium, protactinium, uranium and lanthanides. Bismuth is also essentially immiscible with molten halides. Reductive extraction between metal in molten salt and liquid metal phases can be expressed by the following general reaction: MXn  nLi (Bi) y M (Bi)  nLiX in which a metal halide MXn in the salt reacts with lithium from the bismuth phase to produce M in the bismuth phase and the respective lithium halide in the salt phase. There are three possible ways to influence the selectivity of the method: ●

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Technique utilizing dependency of distribution coefficients on used carrier molten salt and on concentration of reducing agent added in the form of alloy with molten metal (Bi). Technique utilizing selective reducing agents (this technique can be hardly used in MSR fuel reprocessing due to strict requirement on the chemical composition of carrier molten salt). Non-selective technique using direct addition of high concentration of reducing agent into the molten salt.

The distribution data obtained for elements included in MSR spent fuel between bismuth and carrier fuel salt (LiF–BeF2) showed that a multistage extraction process for stepwise separation of lanthanides from thorium is impossible, because the separation factors are close to unity. Therefore, another separation system between molten LiCl and liquid bismuth was chosen for separation of individual elements, but the common non-selective extraction of all metals dissolved in carrier fluoride salt remained to be as the first stage [3,24]. The present knowledge and a significant progress in the development of electrochemical separation methods of actinides and lanthanides also from molten fluoride salt media permits considerations realize that electrochemical processes should assert themselves within the spent fuel on-line reprocessing of future MSRs. The development of suitable technologies is in significant progress in several research laboratories in Europe, Japan, Korea, United States and

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Russia, and first flowsheets of MSR spent fuel reprocessing were brought out. A conceptual flowsheet of MSR on-line spent fuel reprocessing, associating the molten-salt/liquid metal reductive extraction technology and the electroseparation process in molten fluoride salt medium was proposed by research institution in the Czech Republic in the frame of national MSR technology development [25]. The flowsheet is shown in Fig. 7. The concept is based on primary nonselective molten-salt/liquid metal reductive extraction and subsequent selective electrochemical separation method. Li and molten Bi or Cd are planned to be used as reduction and extraction agents, respectively. The reason for such a combination of salt/metal extraction and electroseparation resides is, in fact, that the lanthanides should be removed prior to the actinides, which should remain in the main fuel stream and go back to the MSR. According to the experimentally measured thermodynamic properties of selected lanthanides and actinides, fluorides of the lanthanides are more stable and cannot be reduced neither electrochemically nor chemically prior to actinides without simultaneous reduction of the latter. The proposed flowsheet is based on the removal of all elements from the carrier salt in a form of metallic mixture, from which it would be possible to selectively remove only lanthanides by the electrochemical anodic dissolution method. The fluoride molten salt with satisfactory thermochemical stability has to be applied; the carrier melt composed, e.g. of LiF–CaF2 eutectic mixture has been proposed, because of the thermochemical stability of BeF2, else which is of the component of carrier fuel salt, is unsatisfactory for this purpose [26,27].

Fig. 7. Conceptual flowsheet of MSR on-line reprocessing technology based mainly on moltensalt/liquid metal reductive extraction and electroseparation processes. TRU, transuranium elements; FP, fission products (lanthanides).

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Anodic dissolution method is based on selective electrochemical oxidation of separated elements in reduced form. The differences in their deposition redox potentials provide the selectivity, while with the application of more positive potential, the less stable compounds can be formed and respective metal can be dissolved. It was experimentally proved that fluorides of the lanthanides are more stable than fluorides of the actinides, thus they can exist at more negative potentials. If the potential increases stepwise, the lanthanides are dissolved first at more negative potentials and the actinides remain in the metallic form. The instrumentation of the method is provided by two or three electrodes immersed in the carrier fluoride melt and connected to tuneable stabilized voltage source or potentiostat. Used electrodes are working liquid electrode with controlled potential, counter (auxiliary) electrode providing the charge transfer and possibly reference electrode in electrochemically more complicated systems. The separated metallic mixture can be directly connected as working electrode or inserted into the special conductive electrode basket. After the process starts, the dissolved elements are electrotransported to the counterelectrode and reduced there back to the metallic form [6]. The electrochemical method dealing with cathodic deposition can be used for the removal of dissolved ions, which are less stable than used carrier melt. The process consists of gradually decreasing the potentials of the working electrode and reducing the ions present; the lower the potential is, the more stable compounds are reduced. The instrumentation is similar to the above-mentioned case, however both the working and counterelectrodes could be solid. The proposed technology of the on-line reprocessing of MSR fuel should be further replenished, according to the ORNL proposal, by a helium bubbling system for removal of neutron poisoning gases, xenon and krypton, formed by the nuclear reaction and by the uranium isolation from decaying protactinium. 6. CONCLUSIONS Fluoride technologies applicable within the MSRs fuel cycle mentioned above have a significant chance for the future industrial implementation, because they offer some advantages, which cannot be achieved by hydrometallurgical separation technologies industrially used for reprocessing of spent fuel from currently operated nuclear reactors. The most notable advantages are the compactness of the methods, high radiation resistance of used inorganic chemicals, elimination of any neutron moderators, which cause the criticality problems. Unfortunately, the technologies have a pyrochemical or pyrometallurgical character, which, in combination with utilization of pure fluorine gas, anhydrous hydrogen fluoride or fused fluoride salt, bring about extreme requirements for structural materials. Another complication, typical for handling with the spent nuclear fuel, is caused by high radioactivity of processed material. Therefore the technologies

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have to be placed in hot cells or special shielded boxes and operated by remote control. These crew safety requirements naturally also make research and development of fluoride technologies in nuclear field difficult. However, the fluoride technologies are considered to be among the most promising technologies applicable within the future fuel cycles of advanced nuclear reactor types. REFERENCES [1]

[2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21] [22] [23] [24] [25] [26] [27]

A Technology Roadmap for Generation IV Nuclear Energy Systems, US DOE Nuclear Energy Research Advisory Committee and the Generation IV International Forum, December 2002, GIF-002-00, http://gif.inel.gov. H.G. MacPherson, Nucl. Sci. Eng., 90 (1985) 374. M.W. Rosenthal, P.N. Haubenreich, H.E. McCoy, and L.E. McNeese, At. Energy Rev., 9 (1971) 601. P.N. Haubenreich and J.R. Engel, Nucl. Appl. Technol., 8 (1970) 118. M.W. Rosenthal, P.R. Kasten, and R.B. Briggs, Nucl. Appl. Technol., 8 (1970) 107. J. Uhlir, P. Soucek, G. Modolo, E. Walle, and R. Nannicini, EC/EURATOM Report of the FP5 Project MOST, FIKW-CT-2001–00096, 08/2003 MOST-D7. J. Uhlir, V. Priman, and J. Vanicek, Proc. GLOBAL 2001 – Int. Conf. Back-End Fuel Cycle, CEA, Paris, France, 2001. M. Benedict, T.H. Pigford, and H.W. Levi, Nuclear Chemical Engineering, 2nd edn., McGraw-Hill, New York, 1981, pp. 310–311, Chap. 6. A.A. Jonke, At. Energy Rev., 3 (1965) 3. J.J. Schmets, At. Energy Rev., 8 (1970) 3. M. Bourgeois and B. Cochet-Muchy, Bulletin d’Informations Sci. et Technoques, 161 (1971) 41. M. Bourgeois, Bulletin d’Informations Sci. et Technoques, 188 (1974) 7. E. Yagi, S. Saito, and M. Horiuchi, JAERI Report No. M-6487, 1976, JAERI, Japan. E. Yagi, and M. Maeda et al., JAERI Report No. M-6488, 1976, JAERI, Japan. M.A. Demjanowich et al., RIAR Report No. 50 (565), 1982, RIAR Dimitrovgrad, Russia. P. Novy et al., NRI Report No. 9062Ch, 1989, NRI Rez, Czech Republic. M. Marecek, P. Novy, and J. Uhlir, Proc. GLOBAL 2001 – Int. Conf. Back-End Fuel Cycle, CEA, Paris, France, 2001. O. Amano et al., Proc. GLOBAL 2003, ANS, New-Orleans, Louisiana, USA, 2003. M. Benedict, T.H. Pigford, and H.W. Levi, Nuclear Chemical Engineering, 2nd edn., McGraw-Hill, New York, 1981, p. 476, Chap. 10. P. Novy, I. Peka, and I. Chochlovsky, NRI Report No. 9081Ch, T, 1989, NRI Rez, Czech Republic. J. Uhlir, Proc. 5th OECD/NEA IEM Actinide Fission Product Partitioning Transmutation Mol., Belgium, 1998. Z. Novak, NRI Report No. 5599M, 1980, NRI Rez, Czech Republic. ORNL Report No. 3936, 1966, Oak Ridge, Tennessee, USA. ORNL Report No. 5018, 1974, Oak Ridge, Tennessee, USA. J. Uhlir, V. Priman, and Z. Frejtich, Proc. GLOBAL 2003, ANS, New-Orleans, Louisiana, USA, 2003. H. Boussier et al., Proc. GLOBAL 2003, ANS, New-Orleans, Louisiana, USA, 2003. P. Soucek, F. Lisy, and R. Zvejskova, Proc. GLOBAL 2003, ANS, New-Orleans, Louisiana, USA, 2003.

567

Subject Index ab initio molecular dynamics, 4 absorption, 160 absorption ability, 169 absorption capacity, 162 acetonitrile, 227 Aciplex®, 476 actinides, 564 activated carbon, 359 activation, 20 additives, 313, 411 adhesion, 68 adsorption/dissolution mechanism, 185 affinity, 327 AFM, 16 Ag cathode, 434 Al-doped, 113 alkaline fuel cells, 494, 498, 500 alkylcarbonates, 137 alloy-based anode, 104 alloying/dealloying reactions, 107 aluminate polymers, 345 americium, curium, 557 amorphous Li2O, 106 amorphous phase, 324 analytical purposes, 182 anhydrous hydrogen fluoride (aHF), 351 anion receptors, 336 anion size, 240 anion transport, 240 anion-exchange membrane applications, 494 anode overvoltage, 10 apparent activation energy of ion transport, 238 apparent lithium salt diffusion coefficient, 241 application, 454 aprotic, 307 aprotic organic solvents, 253 Ar mixed O2, 422 argonCHF3, 545 argon-ion sputtering, 158 aromatic cations, 350

aromatic fluorinated macromolecules, 477 Arrhenius curves, 238, 245 Arrhenius equation, 150 Arrhenius’ law, 356 assignment of IR absorption peaks, 36 automobile market, 458 automotive, 454 average discharge potential, 128 Ba2 site, 431 basicity, 200, 213 batteries, 386 battery applications, 397 battery performances, 328 battery safety, 363 B-doped, 112 benzene, 477 BET method, 70 BET surface areas, 45, 51 bis[(perfluoroalkyl) sulfonyl]imide, 225 bis[(perfluoroalkylene)sulfonylimide] acid, 229 bis[(trifluoromethyl)sulfonyl]imide anion, 226 bismuth, 563 blends, 314 bond, 370 borate polymers, 345 boron-based anion receptors, 336 bubble evolution, 21 C1s and F1s XPS spectra, 91 C2F, 33 C5H5NHPF6, 143 capacitance, 360 capacities, 82, 386 capillary forces, 23 carbon, 369 carbon anodes, 3, 138 carbon beads, 86 carbon coating, 43 carbon electrode, 204 carbon nanotubes (CNTs), 89, 412

568

Subject Index

carbonized paper, 71 carbonized paper/TiN/pyrocarbon, 73 carboxylated perfluorovinyl ether, 475 carboxylic, 473 carrier mobility, 519 cassiterite twins, 518, 529 catalysts, 391 catalyzed graphitization, 74 cathode, 386 cathode surface, 276 cathode-active materials, 125 cathodic decomposition, 525 cationic transference number, 241 cavity microelectrode, 527 C–C, 381 CdF2, 545 CdO, 538 CdS, 536 CdS:F, 547 CdSO4, 538 CdTe. Solar, 536 cell voltage, 434 ceramic fuel cells, 419 C–F, 375 C–F bonding nature, 91 CF2, 375 CF3, 390 CF3 groups, 199, 202 CFx, 12 CFx, 385 C–H bonding, 415 13 C–NMR, 393 chair-type, 385 charge – discharge curves, 71, 73, 75 charge – discharge experiments, 91 charge/discharge curves, 47 charge/discharge cycles, 132 charge/discharge efficiency, 130 charge/discharge properties, 126 charged Li0.5CoO2, 276 charge–discharge coulombic efficiencies, 302 charge–discharge measurement, 408 chemical change, 475 chemical diffusion coefficients, 42

chemical modifications, 473, 487, 500 chemical stability, 443 chemical structure, 446 CHF2COOLi, 273 CHF3, 540, 543 chloroaluminate salt, 352 CHP system, 458 chronoamperometry, 204, 215 classical molecular dynamics, 4 ClF3, 128, 129 close-spaced sublimation, 536 (co)polymerisation, 478 CO tolerance, 458 co-deposition of low crystalline carbon, 76 coin cell, 281 cointercalated HF, 38, 42 co-intercalation, 411 coke, 384 combined heat and power, 454 compatibility, 314 complexes F(HF)n, 6 components, 458 conductimetric titration, 142 conduction in fluoride-based conductors, 427 conduction mechanism, 327 conduction network, 62 conductivity, 145, 224, 326, 427, 469, 486, 487, 491, 492, 493, 494, 498, 535 conductivity and viscosity, 353 conductivity values, 487 configurational entropy, 239 contact angle, 15, 25 contact ion pairs, 153 cooperative rearrangement, 238 copolymerisation of fluoromonomers, 490 copolymerisations, 473, 476 corrosion of aluminum, 204, 215, 258 cost, 455 couloumbic efficiency, 82, 213 coupling step, 230 covalent, 377

Subject Index

covalent C–F bonds, 36 covalent nature, 409 covalent-type graphite intercalation compounds, 398 cross-linker, 234 cross-linking agents, 498 crossover, 470 crystalline complexes, 224 crystalline-phase, 324 crystallinity, 312 crystallinity of the carbon, 404 crystallite sizes La and Lc, 401 CSS, 546 cube root law, 148 current collectors, 62 current density, 495 curvature radius, 23 CVD, 61, 516, 529 cycleability, 86, 131 cyclic voltammetry, (CV) 53, 91, 203, 204, 213, 295 cyclic voltammograms, 129, 413 cycling behavior, 282 cycling efficiencies, 290 cycling efficiency, 214 d.o.g., 499 dangling, 382 D-band, 18, 45 decomposition, 276 defects, 383, 403, 427 degenerated tin oxide films, 520 degradation mechanism, 453 degradations, 325, 443 degree of crystallinity, 234 degree of grafting, 492 dehydration, 452 dendrite, 302 density, 288 deposition rate, 514 depression, 324 develop in situ, i.e. in laboratories, a good expertise in polymer science and electrochemistry, 330 dianions, 247 dielectric constant, 289

569

diffusion, 461, 543 diffusion coefficients, 163, 166, 322, 328, 366 dilithium salts, 236, 244 dimeric salt, 232 direct bandgap, 522 direct copolymerisation, 472 direct fluorination, 285 direct methanol fuel cell (DMFC), 474, 487 direct radical (co)polymerisation, 477, 500 direct radical terpolymerisation, 484 discharge behavior, 38 discharge capacities, 48, 128, 387, 40 discharge potentials 40 discharge ratio, 100 discharge reaction, 100 discharged product, 100 discharging performance, 98 dispersions, 447 dissociation, 134, 370 dissociation of lithium salts, 254 DMC, 296 DMF, 233 DMFC, 460 DMFC technology, 460 doped SnO2, 107 doping, 513 double layer 409 Dow®, 476 Drude oscillator model, 523 DSC, 338 DSC thermograms, 234 durability, 453 dye-sensitized nanocrystalline solar cells, 527 earth fluorides, 434 EC/DMC, 280 EC-DEC, 272 edge surface, 56 effect of temperature, 449 efficiency, 537, 547 elaboration process, 311 electric double-layer capacitor, 359

570

Subject Index

electric resistance, 491 electric vehicles (EV), 267 electrical efficiency, 442 electrical power, 461 electrical resistance, 13 electrically conductive pigments, 529 electrochemical, 386 electrochemical devices, 349, 439 electrochemical generation of chlorine, 526 electrochemical impedance spectrometry, 177 electrochemical Li insertion, 91 electrochemical lithium storage, 90 electrochemical properties, 90, 195, 480, 486, 487, 491, 494, 499 electrochemical sensor, 526 electrochemical stabilities, 203 electrochemical stability, 309 electrochemical wastewater treatment, 526 electrochemical window, 153, 357, 360 electroconductive substrates, 63 electrode, 369 electrode reactions, 461 electrolysis, 25 electrolyte, 174, 306, 386 electrolyte purity, 191 electrolytes, 195, 211, 213, 215, 253 electromotive force, 432 electron depleted regions, 520 electron probe, 523 electron spin resonance or paramagnetic resonance, 421 electronic insulation, 444 electronic insulator, 453 electron-withdrawing groups, 345 electro-osmotic, 461 electroosmotic drag, 452 electrospinning, 316 electrostatic spray pyrolysis, 517 elevated temperatures, 361 EMF, 422 EMIm cations, 353 EMIm(CF3SO2)2N, 366

EMIm(HF)2.3F, 366 end-closed SWNTs, 94 energy, 369 energy density, 457 energy of activation, 146 environmentally safe, 2, 454 EPR, 375 equilibrium glass transition temperature, 246 equivalent weight, 446 ESR, 425 ETFE-g-poly(M), 495 ETFE-g-PSSA, 498 ethyl difluoroacetate, 268 ethylene carbonate (EC), 290 exchange current, 526 exfoliation, 401, 411 external aid-process plasticizer, 313 F/C ratio, 94 F2, 128, 129, 385 F ion, 421, 425 FCVI, 62 F-doped nanocrystalline SnO2, 118 F-doped SnO2, 115 F-doping, 120 19 F, 176 19 F-NMR, 229, 374 Fe, 74 FEP-g-PSSA, 498 fermi energy, 118 fill factor, 537 fillers, 313 first coulombic efficiencies, 55, 56 first principle, 133 fission materials, 549 Flemion®, 445, 476 fluoride and hydrofluoride-based electrolytes, 429 fluoride technologies, 566 fluoride volatility method, 554 fluorides, 370, 423, 552 fluorinated, 390 fluorinated carboxylic acid esters, 268 fluorinated electrolyte salt, 258 fluorinated MWNTs, 90

Subject Index

fluorinated polymers, 470, 472 fluorinated single-wall carbon nanotubes (fluorotubes), 96 fluorination, 79, 89, 370, 555 fluorination of single-walled carbon nanotubes, 93 fluorination treatment, 127 fluorine – graphite intercalation compound, 31 fluorine, 1, 195, 247, 369, 541, 544 fluorine atom, 288, 293 fluorine contents, 78 fluorine doping, 104 fluorine gas, 26, 134 fluorine–carbon nanocomposite surface, 130 fluorine-for-hydrogen substitution, 212 fluoroacids, 352 fluoroanions, 169 fluoroborates, 153 fluorocarbon, 378 fluorocomplex anion, 353 fluorographite, 377 fluorohydrogenate salts, 351 fluorophosphates, 153 fluoropolymers, 169 fluorotin alkoxides, 115 foil-type current collector, 65 fossil fuel, 454 “front-end” reprocessing technology, 555 FT–IR, 272, 375 FTO, 514, 529 FTO powders, 527 fuel cell, 360, 419, 429, 469, 476, 499 fuel cell membranes, 476, 495, 499, 500 fuel cell operation, 441 fuel cell performance, 456 fuel cell vehicle, 458 fuel crossover, 453 fullerenes, 371 galvanostatic discharge – charge curves, 92 Galvanostatic discharge, 388 gas bubble, 24 gas permeation, 453

571

G-band, 18, 45 GC–MS analysis, 296 gel phase, 327 gelled polymer electrolyte, 309 Generation IV International Forum, 551 geometric surface area, 65 Germanium, 540 g-factor, 383 GIC, 14, 375 glass transition temperature, 234 glassy state, 246 GO, 399 grafting from, 490 grain boundaries, 519, 545 graphene, 393 graphite, 105, 154 graphite electrode, 159 graphite fluoride (CF)n, 90 graphite fluorides, 31, 38, 373 graphite intercalation compounds, 397 graphite oxide, 15, 397 graphite-based anodes, 83 graphitized, 371 GSM discharges, 330 hard carbon, 407 hard-carbon powder, 84 heat management, 456 heat-treated petroleum cokes, 53 Henri Moissan, 3 hexafluorophosphate (LiPF6), 173 hexafluoropropene, 312 HF, 4, 178, 372, 515 HFP, 474 high capacity, 181 high lithium transference, 226 high oxidation state transition metal complex fluoride, 33 high temperatures, 175 high-energy radiation, 491 higher discharge potentials, 405 highly fluorinated graphite, 33 highly polar electrolytes, 308 HiPco – SWNTs, 95 HOMO energies, 299 HOPG, 12

572

humidification, 361, 456 hybrid, 385 hybrid conductors, 435 hybrid electric vehicle, 458 hybrid proton, 428 hybridization, 379 hydration, 448 hydration level, 451 hydride ion (H), 430 hydride ion vacancy, 431 hydride ions, 422 hydrodynamic radius, 145 hydrofluoride solid, 430 hydrogen, 414, 421, 422 hydrogen adsorption, 188 hydrogen electrode positive, 432 hydrogen storage, 412, 455, 460 hydrolysis, 127, 134, 175, 209 hydrolysis of anions, 262 hydrolytic and thermal stability, 195 hydrolytic instability, 205 hydroperoxides, 490 hydrothermal conditions, 107 Hyflon®, 476 hyperfine, 384 ICVI, 62 IEC, 494, 495, 500 IEM, 495 IF6, 375 IF7, 376 IFO, 547 imide compounds, 256 impedance diagrams, 21 impedance measurements, 42 In2O3, 538, 541 indium oxide (IGO), 540 inductive loop, 188 inflammable, 363 influence of cointercalated HF, 40 inner-surface fluorination, 91 Inter Penetrated Network, 314 interaction parameters, 319, 320 intercalation, 303, 384 intercalation of lithium ion, 409 interfacial potential barrier, 525

Subject Index

interfacial space charge, 525 intermediate discharge product, 42 iodine, 377 ion – polymer interactions, 167 ion cluster, 448 ion interactions, 166 ion mobility, 209 ion pairs, 149 ion-conducting, 441 ion-conductive polymer electrolytes, 335 ion-exchange capacity, 446 ion-exchange membrane, 493 ion-exchange membranes for electrolysis, 476 ionic, 386 ionic conductivity, 174, 235, 244, 254, 336, 338, 345, 442, 499 ionic conductivity of aprotic solvent, 255 ionic exchange capacity, 487 ionized impurities, 521 ionomer membranes, 486 ion-pairing, 200 irreversible reactions, 86 ITO, 540 K2MnF6, 33 K2NiF6, 33 K2PdF6, 33 KAgF4, 33 Karl – Fisher, 176 KF–2HF, 4, 19, 26 kinetic absorption, 166 kinetic aspects, 322 kinetic parameters, 163 kinetics study, 177 lactone, 137 lanthanides, 554, 565 lattice constant, 94 lattice energy, 224 lattice parameters, 35, 75 layers, 383 lengths, 372 Li cells, 494 Li metal, 270 Li(C5H5N)PF6, 144 Li(CH3CN)4PF6, 134, 142

Subject Index

Li(Py)PF6, 144 Li/ fluorotube cell, 98 Li/carbon monofluoride (CFx) cell, 96 Li/LiCoO2 coin cells, 302 Li/SnO2 cells, 117 Li2CoPO4F, 127 Li2SO4, 420 Li4Ti5O12, 364 LiAsF6, 140 LiBF4, 140 LiBOB, 264 LiCF3SO3, 256 LiCoO2 electrode, 204 LiF, 180, 391 LiF–BeF2, 563 LiF–CaH2(–Al2O3), 431 lifetime, 453 LiF–MgF2 system, 429, 430, 435 Li-ion batteries, 119 Li-ion cells, 267 LiMxOyFz, 127 LiPF6, 125, 140, 191, 196, 202, 203, 205 liquid anhydrous hydrogen fluoride, 134 liquid electrolytes, 139, 307 liquid methanol, 461 Li–Sn alloys, 105 LiTFSI, 140 LiTFSI, 226, 246 lithiated carbon anode, 281 lithiated cobalt oxide, 138 lithiated graphite anode, 176 lithium, 137 lithium batteries, 103, 191, 253, 285, 306 lithium battery, 363, 386 lithium electrode cycling efficiencies, 289 lithium hexafluorophosphate (LiPF6), 133 lithium hexafluorophosphate, 140 lithium hydroxide, 142 lithium oxyfluorophosphate, 134 lithium-imide salts, 270 lithium-ion batteries, 173, 195, 335 lithium-ion battery, 125 lithium-ion rechargeable batteries, 491 lithium-ion secondary batteries, 61

573

lithium-ion transference numbers, 336, 338, 340, 345 lithium-ion-conductive polymer electrolytes, 335 lithium-polymer, 305 longer run time, 463 Lorentzian fits, 116 low or zero emissions, 440 LUMO energies, 300 macroinitiators, 487 macroporous PVdF, 309, 315 macroscaled, 62 macroscopic or microscopic degradation, 323 magnetron sputtering, 515 MCMB, 281 Me4Sn, 516 mean molar activity coefficients, 243 mechanical coupling, 238 mechanical properties, 314 MEK, 316 melting point, 324 melting point and glass transition temperature, 353 membrane electrode assembly (MEA), 441 membrane-electrode-assembly, 470 membranes, 474, 491, 493, 494, 495 membranes for chlorine-alkali electrolysis, 475 membranes for fuel cells, 494, 501 membranes in chlor-alkaly, 476 mercury porosimetry, 318 metal (Ni) foam, 65 metal transfer process, 562 metallic oxides, 535 metathesis, 351 methanol crossover, 487 methanol permeabilities, 476 methanol permeability, 452 methanol transport, 461 methyl difluoroacetate, 267 microcrystalline tin dioxide, 113 microstructure, 317, 445 mixed solvent electrolyte, 277

574

Subject Index

mobility, 200 modern electronic technology, 103 molar conductivity, 147, 150, 196 molecular orbital calculation, 133 molten salt breeder reactor, 550 Molten-salt reactors, 549 molten-salt/liquid metal reductive extraction, 563 morphology, 448 Mössbauer spectroscopy, 116 MSRE program, 560 multi-wall carbon nanotubes (MWNTs), 89 NaF, 426 NaF–CaF2, 425 Nafion®, 360, 439, 445, 474, 476, 486, 487, 492, 493, 494, 495 Nafion® membranes, 476 nanocrystalline materials, 93, 120 nanometer-sized FTO, 518 nanoparticle, 515 nanoscaled structure, 68 nanothickness carbon, 129 nascent fluorine, 22 natural graphite, 281 N-butyl-N-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide, 359 neptunium, 558 neutralization, 352 new cathode-active materials, 127 NF3, 128, 129 NF–MF2, 423 NH4F, 515 Ni, 74 NMP, 316 NMR, 176, 328 nonaromatic cations, 350 non-blocking electrodes, 242 non-ideal behavior, 243 non-solvent, 315 non-volatile fluorides, 557 non-volatile trifluoride, 554 n-type semiconducting electrodes, 119 nuclear energy, 1 nuclear power, 551

Nyquist plot, 187 O2, 422 O2 defect, 428 Oak Ridge National Laboratory, 550 OCV, 390, 434 OCV value, 91 oligomeric anions, 226, 247 oligomeric chains, 247 open-circuit voltages, 362 open-end SWNTs, 94 opening of edge plane, 57 orientation of carbon layers, 403 overpotential, 405 overswelling, 319 oxidation, 393 oxidative destruction, 526 oxygen, 422 oxygen defects, 127 oxygen ions, 436 oxygen vacancies, 541 31 P- and 19F-NMR, 135 partial fluorination, 285 passivation, 215 passivation film, 155, 175, 289 PC, 281 PCVI, 62 PEG, 226 PEG cross-linking, 243 PEMFC stacks, 455 PEO, 224, 226 perfluorination, 196 perfluoroalkoxyalkyl vinyl ethers, 476 perfluorocyclobutane (PFCB), 480 perfluorovinyl ethers, 475 performance criteria, 195, 215 permability to methanol, 469 peroxides, 490, 492 PF5, 269 PF6, 210 PFA (perfluoroalkoxy), 286 PFA-g-PSSA, 499 PFA-g-PSSA membrane, 499 PFCB, 482 PFSA, 444 phase inversion, 315

Subject Index

phase transitions, 235 phase-inversion mechanism, 316 phase-inversion processes, 315 phosphonic, 473 phosphonic acid, 475, 476, 480 phosphonic acid TFS, 479 photocurrent response, 527 photovoltaic modules, 535 physical properties, 447 planar defects, 518 plasma fluorination, 45, 51 plasma frequency, 523 plasma-assisted CVD, 517 plasma-fluorinated graphite, 49 plasticizers, 241 plutonium hexafluoride, 558 PMMA, 314 POF2OH, 178 POF3, 178 polarization behaviors, 361 polarization curves, 407 poly(carbon monofluoride), 397 poly(dicarbon monofluoride), 397 polyanionic lithium salts, 236, 244, 248 polyanionic series, 231, 240 polydispersity index, 232 polyethers, 223 polyfluorides, 5 polyfluorination, 196 polymer electrolyte fuel cells, 475 polymer electrolytes, 306, 335, 338, 491 polymer matrix plasticization, 248 polymer-gel electrolytes, 139 polyolefins, 308 polyparaphenylenes, 487 polyvinylidene fluoride (PVdF), 306 pore diameter, 318 pore wetting, 308 pores, 163 porosity, 167, 317, 318, 327 porous separator, 308 portable, 460 portable devices, 462 post-treatment, 374 potentiostatic charge, 132

575

powdered ceramic, 556 power density, 407, 495 power output, 455 power sources, 267 precursors 514 preferred orientation, 518 primary lithium battery, 31 primary lithium cells, 90 primary lithium cells with fluorinated single-wall carbon nanotube, 93 probe, 191 production of solar cells, 542 propylene carbonate (PC), 83 propylene carbonate solution, 410 protactinium, 562 proton carriers, 429 proton conductivity, 448 proton reduction, 183 proton transfer, 450 proton transport mechanism, 450 proton-exchange membrane fuel cell (PEMFC), 469, 470, 478, 480, 482, 487 proton-exchange membrane fuel cells, 439 proton-exchange membranes for fuel cell applications, 472, 493 protonic conduction, 473 protons, 436 pseudo-lattice model, 148 Pt anode, 434 PTFE-g-PSSA, 495 PTFE-g-PSSA copolymers, 495 puckered sp3 graphene layers, 38 pulsed-gradient spin-echo (PGSE) NMR, 355 PV cells, 528 PV module production, 536 PVdF, 330 PVDF-g-PSSA, 498 PVDF-g-PSSA membranes, 492 PVDF–HFP copolymer, 159 pyridinium hexafluorophosphate, 142 pyrocarbon, 68, 74, 79 pyrocarbon coating, 83

576

Subject Index

pyrometallurgical separation, 562 R value, 79 Raman shifts, 45, 52 Raman spectra, 74 Raman spectroscopy, 152, 17 rate of absorption, 162 RBM (radial breathing mode), 95 RDF analysis, 9 rechargeable batteries, 305 rechargeable lithium batteries, 223 reciprocal of the half-width of (002) line, 401 reduction, 388 reductive decomposition, 181 reformer, 455 relaxation time, 523 renewable energy sources, 454 reproducibility, 236 residual carbons, 398 residual currents, 215 resistive heaters, 529 resistivities, 495 resistivity, 513 resonance, 225 resonant nuclear reaction, 523 reversibility, 391 reversible capacity, 86, 93 reversible modifications, 324 RF sputtering, 539 Rietveld analysis, 127 room-temperature ionic liquids, 349 room-temperature molten salts, 349 Rutherford backscattering, 523 safety, 267 safety aspect, 308 salt, 321 saturation, 160 Sb2Te3, 537, 538 Sb-doped SnO2, 112 scanning electron microscopy, 204 scattering mechanism, 520 secondary (rechargeable) lithium batteries, 32 secondary ion mass spectroscopy, 523 segmental motion, 241

SEI, 175 SEI formation, 53 SEI layer, 154, 158 SEM images, 84 semiconductor, 513 semi-crystalline, 310 semi-ionic, 372 semi-ionic C–F bond, 33, 35 several types of crystalline phases, 310 SF6, 517 shut-down effect, 308, 325 SiC, 63 Si-doped SnO2, 111 silica, 160 silica fillers, 322 simulation, 7 single-ion-conducting polymer electrolytes, 343 single-wall carbon nanotubes (SWNTs), 89 six F ions, 426 skin on the surface, 317 slow scan voltammetry, 177 Sn clusters, 107 Sn1xMoxO2 mixed oxides, 109 SnBuCl3, 515 SnCl4, 515 SnO, 543 SnO2, 543 SnO2 doped with Mo, 107 soda-lime glass, 538 SOFC, 419 soft anion, 256 soft carbon, 71, 407 soft chemistry route, 517 solar cells, 365, 542, 545 sol–gel, 517 solid carbon–fluorine, 10 solid electrolyte interface (SEI), 262, 289, 410 solid electrolyte interphase (SEI), 270 solid polymer electrolytes, 223 solvent, 315 solvent impurities, 215 solvent separated pairs, 153

Subject Index

solvent/polymer interactions, 309 sp2, 379 sp3, 377 specific and volumetric energies, 329 specific conductivities, 289 specific resistance, 495 spin, 384 spray pyrolysis, 514, 529 stability in oxidation, 312 stability of the electrolyte salt, 261 stage, 1 35 stage transitions, 183 stationary, 454, 458 statistical or random copolymers, 311 STEM image, 99 step-growth polymerization, 229 steric hindrance, 166 STM, 11 structural modifications, 323 structural parameters of the host graphite, 398 structure, 319 subband-gap states, 120 sulfonimides, 225 sulfonylimide, 228 sulphonated perfluorovinyl ethers, 475 sulphonated polystyrenes, 470 sulphonated PVDF-g-PS, 491 sulphonic, 473 sulphonic acid, 492 sulphonic acid–TFS, 480 sulphonyl fluoride perfluorovinyl ether, 473 superacid lithium salt, 307 superweak anions, 195 surface analysis, 259 surface area, 52, 70 surface composition, 51 surface fluorination, 129, 32, 43 surface fluorine concentrations, 44 surface modification, 32 surface oxidation, 43 surface tension, 23 surface treatment, 411 surface-fluorinated carbon materials, 43

577

surface-fluorinated natural graphite, 44 surface-fluorinated petroleum coke, 51 surface-fluorinated samples, 82 swelling, 167, 499 swelling behaviour, 319 swelling selectivity, 320 swelling tests, 492 synthesis, 445 synthetic methods, 350 target price, 457 TCO, 547 TCO thin films, 537 TEM and SEM images, 96 temperature ceramic fuel cells, 420 template-synthesized MWNTs, 91 terpolymerised, 475 tetrafluoroethane-β-sultone, 474 TFS, 477, 478 TFVOB, 477, 478, 480 TGA, 231 TG-DSC, 268 232 Th–233U cycle, 551 theoretical capacity of graphite, 49 thermal cyclopolymerisation, 480, 486 thermal decomposition, 176, 196, 401 thermal degradation, 310 thermal engine, 469 thermal stabilities, 215 thermal stability, 169, 174, 487 thermocyclodimerisation, 484 thermodynamical, 319 thermomechanical properties, 310 thermostability, 477 thermostable polymers, 472 ThO2, 552 thorium tetrafluorides, 549, 552 TiC, 63 TiN, 68 tin dioxide, 513 tin oxide, 104 tortuosity, 327 Tosflex®, 476 transference number, 328 transparency, 535 transparent, 541

578

Subject Index

transport, 445 transported charge, 423 triflic acid, 228 trifluoromethane, 540 trifluorovinyl ethers functionalised, 482 triple ion, 210 triple-ion formation, 196, 210 tris(pentafluorophenyl)borane (TPFB), 336 tunnelling, 525 twins, 518 two-stage sorption, 322 U3O8, 555 UF6, 1 uninterrupted power supply, 463 unstable vs. lithium, 312 uranium hexafluorides, 559 VdF polymers, 330 viscosity, 145, 196, 200, 202, 209, 210, 288 Vogel – Tamman – Fulcher (VTF), 180 volatile UF6, 554 voltage, 388 voltammetry, 391 voltammogram, 393 VTF, 246 VTF equation, 151 Walden plot, 355 Walden’s rule, 355 Warburg behavior, 188 water, 456 water balance, 451 water diffusion coefficient, 451 water scaling, 529 water uptake, 495, 499 water-contaminated electrolytes, 184 weakly coordinating, 209, 210 weakly ion pairing, 209, 210 wettability, 14

window layer, 544 wood/TiN/pyrocarbon, 69 XPS, 272 XPS analysis, 158 XPS spectra, 129 X-ray diffraction, 399, 413, 421 X-ray photoelectron spectroscopy, 538 X-ray structure, 199, 202, 210 XRD, 16, 79, 380 ZnO, 543 ZnO:Al (AZO), 543 ZnO:F film, 544 [(α,β,β-trifluorovinyl)oxy] benzene (TFVOB), 480 α form, 310 α phase, 319 α,β,β-trifluorostyrene (TFS), 477, 478, 479 α,β,β-trifluorovinyl benzyl ethers, 486 α-F–γ-BL, 286 β-F–γ-BL 286 γ-BL, 286 γ-butyrolactone, 145 γ-F–γ-BL, 286 γ-valerolactone, 159 1-ethyl-3-methylimidazolium chloride, 351 1M LiPF6/EC-DMC, 281 2.61V, 364 4-[(trifluorovinyl)oxy]bromobenzene, 482 4-[(α,β,β-trifluorovinyl)oxy] bromobenzene, 485 4C batteries, 330 4-fluoro-ethylene carbonate (FEC), 290 4-TFVOB phosphonic acid, 482 4-TFVOB sulphonic acid (TFVOBSA), 486