Evaluation of electrolytes for redox flow battery applications

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1 Electrochimica Acta 52 (2007) Evaluation of electrolytes for redox flow battery applications M.H. Chakrabarti a, R.A.W. Dryfe b, E.P.L. Roberts a, a School of Chemical Engineering and Analytical Science, The University of Manchester, P.O. Box 88, Manchester M60 1QD, UK b School of Chemistry, The University of Manchester, P.O. Box 88, Manchester M60 1QD, UK Received 17 May 2006; received in revised form 1 August 2006; accepted 17 August 2006 Available online 26 September 2006 Abstract A number of redox systems have been investigated in this work with the aim of identifying electrolytes suitable for testing redox flow battery cell designs. The criteria for the selection of suitable systems were fast electrochemical kinetics and minimal cross-contamination of active electrolytes. Possible electrolyte systems were initially selected based on cyclic voltammetry data. Selected systems were then compared by charge/discharge experiments using a simple H-type cell. The all-vanadium electrolyte system has been developed as a commercial system and was used as the starting point in this study. The performance of the all-vanadium system was significantly better than an all-chromium system which has recently been reported. Some metal organic and organic redox systems have been reported as possible systems for redox flow batteries, with cyclic voltammetry data suggesting that they could offer near reversible kinetics. However, Ru(acac) 3 in acetonitrile could only be charged efficiently to 9.5% of theoretical charge, after which irreversible side reactions occurred and [Fe(bpy) 3 ](ClO 4 ) 2 in acetonitrile was found to exhibit poor charge/discharge performance Elsevier Ltd. All rights reserved. Keywords: Redox flow battery; Vanadium; Chromium; Ru(acac) 3 ; [Fe(bpy) 3 ](ClO 4 ) 2 1. Introduction Redox flow batteries are electrochemical energy storage devices that utilise the oxidation and reduction of two soluble redox couples for charging and discharging. They differ from conventional batteries in that the energy-bearing chemicals are not stored within at the electrode surface, but in separate liquid reservoirs and pumped to the power converting device for either charging or discharging [1,2]. Due to the use of two soluble redox couples, solid-state reactions with their accompanying morphological changes at the electrodes are absent [3]. Thus, there are no fundamental cycle life limitations associated with these processes such as shedding or shape changes, which usually occur in conventional storage batteries. Despite these advantages, the redox flow battery has not been widely exploited to date. One disadvantage of the systems developed to date is the use of two separate redox species in the halfcells, leading to the potential for cross-contamination of active Corresponding author. Tel.: ; fax: address: edward.roberts@manchester.ac.uk (E.P.L. Roberts). electrolytes by transport through the membrane. For example, there has been little recent interest in the development of the iron/chromium redox flow cell due to this problem [4]. To redress this issue, an all-chromium redox electrolyte was investigated at the University of Manchester and the charge/discharge characteristics of a laboratory scale battery were reported [5,6]. Prior to this, other workers have performed extensive investigations on the all-vanadium redox system [3,7 10] and patented the technology [11]. In addition, an all-neptunium system has been evaluated [12], although the hazards of working with radioactive electrolytes are likely to limit the practical application of this system. Several prototype vanadium systems have been investigated successfully [13 16] and some systems are well on their way to commercial success [17]. Despite such achievements, batteries employing aqueous electrolytes have a low energy content. The energy output from the battery is proportional to the potential window of operation available from the background electrolyte. The operating potential window of aqueous electrolytes is limited due to water electrolysis [5]. Organic electrolytes, which offer a wider potential window, have been investigated in this study. In addition, species have been selected which minimize the effect of elec /$ see front matter 2006 Elsevier Ltd. All rights reserved. doi: /j.electacta

2 2190 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) trolyte cross-contamination. One approach is to use a single system which offers three oxidation states, so that the discharged species is the same on each side of the cell. Such a system would have the advantage that any cross-contamination would only lead to some self-discharge, and little or no cell balancing or electrolyte processing would be required. An approach whereby the same cation is used but with different ligands on each side of the cell has been suggested [18], but this has not been considered in this study. Electrolyte systems can be selected on the basis of the following properties, which are generally desirable for redox flow batteries [18,19]: fast kinetics at the electrode electrolyte interface; a relatively large open circuit potential; reasonable cost; high solubility in the process electrolyte. In this study, the following series of redox systems in acetonitrile electrolyte were selected which apparently offered fast electrode kinetics (based on literature data, e.g. [19] and [20]) and the potential to operate with a single electrolyte using a species with three oxidation states. (i) Ruthenium organic complexes A number of ruthenium organic complexes which can be both oxidized and reduced electrochemically have been reported in the literature, and some of these have been suggested as suitable candidates for a redox flow battery [21]. Tris(2,2 -bipyridine) ruthenium (II) tetrafluoroborate [Ru(bpy) 3 (BF 4 ) 2 ] has exhibited fast kinetics [19]. In addition, this system offers the possibility of cell voltages of up to 2.6 V, much higher than is possible in aqueous battery systems [19]. Ruthenium acetylacetonate [Ru(acac) 3 ] has also been reported as offering fast oxidation and reduction kinetics [22] and a possible cell voltage of around 1.75 V. (ii) Tris(2,2 -bipyridine) iron(ii) perchlorate This species can be oxidized and reduced [19] and offers a possible cell potential of 2.4 V. This compound is available commercially and is significantly cheaper than the ruthenium complexes. (iii) Rubrene Rubrene, a neutral organic species, can be oxidized and reduced electrochemically [23]. The redox potentials of these reactions offer a possible cell potential of around 2.3 V. Again this compound is available commercially, although it is significantly more expensive than the other redox species. These systems are compared to the all vanadium redox flow battery system, which has previously been investigated in detail ([3,7 11]) and has been commercialized in recent years [17]. In this system, vanadium in four different oxidation states is used: V(II)/V(III) at the negative electrode and V(IV)/V(V) at the positive electrode. Each system was first tested by cyclic voltammetry in order to evaluate the electrode kinetics. While cyclic voltammetry can give an indication of the reversibility of redox couples, further experiments are needed to demonstrate that selected systems can be used for energy storage. For example, a redox couple may be reversible, but the charged species may be unstable over long timescales, which would not be detected by cyclic voltammetry. In this study systems which were found to exhibit fast kinetics were tested for their charge/discharge performance in a simple H-type cell. These experiments aimed to determine whether the selected systems could be used for energy storage and to provide a preliminary indication of the relative performance of each system. 2. Experimental 2.1. Electrolytes Vanadium electrolytes were prepared from vanadium (IV) sulphate (>99.99% purity, Aldrich), with the V(II)/V(III) couple generated by electro-reduction. Sulphuric acid was used as the background electrolyte. Reagent grade tris(2,2 -bipyridine) ruthenium (II) chloride is available from Aldrich. Since oxidation of the chloride salt was known to be irreversible [24], the tetrafluouroborate salt [Ru(bpy) 3 (BF 4 ) 2 ] was prepared by addition of NaBF 4 in acetonitrile and precipitation of NaCl. Ruthenium acetylacetonate [Ru(acac) 3, 97% purity, Aldrich], tris(2,2 -bipyridine) iron(ii) perchlorate (reagent grade, GFS) and rubrene (reagent grade, Aldrich) were used for the preparation of the respective electrolytes (Caution: perchlorate salts are potentially explosive and should be handled with appropriate care). Tetraethyl-ammonium tetrafluoroborate and tetraethyl-ammonium perchlorate were used as the background electrolyte. To remove dissolved oxygen, electrolytes were sparged for at least 10 min with oxygen-free dry argon (aqueous electrolytes) or nitrogen (organic electrolytes). Water was removed from the organic electrolytes using zeolite 4A (Merck) to a moisture level of below wt% Cyclic voltammetry A graphite rod (Goodfellow) of surface area 0.06 cm 2 was used as the working electrode for cyclic voltammogram experiments with the vanadium battery electrolytes. A glassy-carbon electrode (I.J. Cambria Scientific) of surface area 0.07 cm 2 was used for cyclic voltammetry in organic media. The electrode was polished with alumina washed with de-ionised water and acetone following the procedure described in literature [21]. The reference electrode used in aqueous solutions was the saturated calomel electrode along with a salt bridge. Organic media required the use of a silver wire quasi-reference electrode (AgQRE). A platinum counter electrode was used in each case. Cyclic voltammetry was conducted using a standard threeelectrode cell, with a Autolab/PGSTAT30 potentiostat for potential control. All solutions were de-aerated prior to experiments. The solution headspace was purged with inert gas for the duration of experiments.

3 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) Fig. 1. Schematic diagram of a glass cell apparatus for small-scale charge/ discharge tests of redox couples Charge/discharge experiments A schematic diagram of the H-type test cell is shown in Fig. 1. The constant current (galvanostatic) charge/discharge characteristics of the redox couples were used to evaluate their performance in a prototype redox flow battery. Each electrolyte compartment contained 40 ml of electrolyte, except for the initial charging of the all-vanadium system, where twice as much electrolyte (80 ml) is required in the anodic compartment. After the first charging of the all-vanadium half of the electrolyte in the anodic compartment was removed [3]. Graphite felt electrodes (Sigratherm GFA 10) were employed for charge/discharge experiments in the H-type glass cell. The graphite felt electrode had dimensions of 30 mm 15 mm 10 mm. Graphite rods were used as current collectors. An Ultrex TM (Membranes-International Ltd.) anion exchange membrane was used for vanadium charge/discharge tests and a Neosepta AHA membrane (Eurodia Industrie SA) was used for organic charge/discharge experiments in all cases. Membranes were pre-conditioned by exposing them to the required test solution for at least 6 h prior to experiments. The circular area of the membrane exposed to the electrolyte in the cell had a diameter of 27.5 mm. The charge/discharge experiments were carried out under constant current conditions. The current was selected on the basis of preliminary experiments. For the vanadium electrolytes, which had a much higher conductivity than the organic electrolytes, a high charging current of 100 ma was used. Although this led to a high cell voltage during charging ( 4 V) with the likelihood of side reactions, the aim was to attempt to charge the cell close to its maximum capacity. For the organic electrolytes, much lower charging currents were used in order to minimize ohmic losses and to evaluate whether efficient charging could be achieved. Discharge currents were selected to ensure that significant cell voltages were obtained. For the vanadium system, a discharge current of 2 ma was found to be suitable while with the lower conductivity organic electrolytes, it was necessary to use lower discharge currents (0.5 ma or lower). The applied current was controlled using a galvanostat and the total cell potential and the potential of each electrode were monitored [relative to saturated calomel electrode (SCE) or silver quasi reference electrode (AgQRE)] throughout each experiment. Since a silver quasi reference electrode was used in the organic electrolytes the absolute value of the electrode potentials is not meaningful, and the measurements can only be used to monitor significant changes in the potential of each electrode. In aqueous experiments mass transport was provided by means of sparging the solution with argon gas, while for organic solutions a magnetic stirrer was used. 3. Results and discussion 3.1. Cyclic voltammetry The results of cyclic voltammetry experiments for the vanadium system suggested that the kinetics of the V(II)/V(III) couple were relatively fast, while the V(V)/V(IV) redox couple was found to be irreversible, consistent with results reported in literature [25 27]. All of the organic electrolyte systems studies demonstrated reasonably fast (in most cases reversible) kinetics for both oxidation and reduction reactions. The results suggested that all four systems could be oxidized or reduced, confirming their suitability for a redox flow battery with the same species occurring in the discharged state. Table 1 compares the species against each other based on the possible open circuit potential of a battery based on the system, their solubility, electrochemical reaction kinetics, and cost. The information on the kinetics was obtained from cyclic voltammetry experiments and from literature data. The Ru(acac) 3 and Fe(bpy) 3 (ClO 4 ) 2 systems were selected for further study on the basis of their superior solubility and fast kinetics. The relatively low cost of the Fe(bpy) 3 (ClO 4 ) 2 makes this system particularly attractive. Cyclic voltammograms for the Ru(acac) 3 and Fe(bpy) 3 (ClO 4 ) 2 systems are shown in Figs. 2 and 3, illustrating the combination of multiple redox couples with fast kinetics. The cyclic voltammogram for Fe(bpy) 3 (ClO 4 ) 2 indicates that the species can be reduced at least twice, consistent with previous studies [19]. The reduction reactions are presumed to be single electron reductions of Fe(bpy) 3 2+ to Fe(bpy) 3 + and Fe(bpy) 3. Table 1 The characteristics of the redox species studied in acetonitrile Chemical Expected open circuit potential (V) Solubility in solvent Reaction kinetics [19,22 24] Approximately cost per mmol Ru(acac) High Reversible 12 [i] Ru(bpy) 3 (BF 4 ) Poor Quasi-reversible 21 [i] [Fe(bpy) 3 ](ClO 4 ) Moderate Reversible 2 [ii] Rubrene 2.33 Poor Reversible 37 [i] Costs were obtained from Sigma Aldrich [i] and GFS Chemicals [ii].

4 2192 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) Fig. 2. Cyclic voltammograms recorded at 0.1 V s 1 at a GC electrode in: (a) 2 mm Ru(acac) 3 and 0.05 M TEABF 4 in acetonitrile and (b) 0.05 M TEABF 4 in acetonitrile Charge/discharge of the vanadium redox system During the first charging of the VOSO 4 electrolyte, the V(IV) species must be reduced to V(II) at the cathode and oxidized to V(V) at the anode. Consequently, as same electrolyte concentration was used in each compartment, twice as much electrolyte (80 ml) was used in the anodic compartment during the first charging [9]. For subsequent cycles, equal volumes of electrolyte were used in each compartment [V(III)/V(II) and V(IV)/V(V)]. The charge/discharge reactions for the second and subsequent charge/discharge cycles of the VOSO 4 electrolyte are shown below: Positive half-cell [V(IV)/V(V)]: VO 2+ + H 2 O VO H + + e Negative half-cell [V(III)/V(II)]: V 3+ + e V 2+ Fig. 4. Charge/discharge potential time profile of 0.1 M VOSO 4 solution in 2 M H 2 SO 4 using graphite felt electrodes and Ultrex TM anion exchange membrane. Constant charging current of 100 ma for 300 min followed by constant discharge at 2 ma constant current. Electrode potentials were measured relative to a SCE. The charge/discharge of 0.1 M VOSO 4 was performed using an Ultrex TM AEM. The voltage profile during the charge and discharge (second cycle) is shown in Fig. 4. The cell was charged at 100 ma for 300 min, and was discharged at 2 ma until the cell voltage dropped to zero. A high charging current was used in order to attempt to fully charge the cell, with around 2.3 times the theoretical charge passed. This high charging current also led to a relatively high voltage during charging ( 4 V). The open circuit voltage after charging was high at 1.61 V, and the cell voltage remained above 1 V during most of the discharge process. The overall efficiency was found to be 5.6% (18.4% charge efficiency and 31% voltage efficiency) and Wh of energy was recovered from the 80 ml of charged electrolyte. The low overall efficiency obtained is a consequence of the high charging current used. The charge recovered is around 86% of the theoretical capacity, so that significant side reactions must have occurred during charging. Furthermore the average voltage during discharge is around 1.2 V, >80% of the theoretical potential which can be achieved with the vanadium system. In spite of the poor overall efficiencies obtained, the discharge results indicate that the vanadium system can achieve high efficiencies, as expected [28]. Furthermore, the results compare favourably to the allchromium system studied by Bae et al. [6] in a similar cell. The cell potential was lower during charging in this study, but only 14% of this potential was due to ohmic drop (estimated based on the electrolyte conductivity and cell geometry), compared to the 50% reported for the all-chromium system [6]. The lower ohmic loss is associated with the use of the high conductivity 2MH 2 SO 4 supporting electrolyte Charge/discharge of Ru(acac) 3 and [Fe(bpy) 3 ](ClO 4 ) 2 systems Fig. 3. Cyclic voltammograms recorded at 0.1 V s 1 at a GC electrode in: (a) 2 mm Ru(acac) 3 and 0.05 M TEABF 4 in acetonitrile and (b) 0.05 M TEABF 4 in acetonitrile. Charge/discharge of electrolytes consisting of ruthenium acetylacetonate [Ru(acac) 3 ] and tetraethylammonium tetrafluoroborate (TEABF 4 ) were carried out in the H-type glass cell using graphite felt electrodes and a Neosepta anion exchange membrane. The reactions occurring at the electrodes are shown below:

5 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) Fig. 5. Potential time profile during charging of 0.1 M Ru(acac) 3 with 1 M TEABF 4 in acetonitrile using 1 ma constant current to 8.3% SOC in a stirred H-type glass cell with graphite felt electrodes and Neosepta anion exchange membrane. Electrode potentials were measured relative to a AgQRE. Fig. 7. Potential time profile during recharging of 0.1 M Ru(acac) 3 and 1 M TEABF 4 in acetonitrile at a constant current of 1 ma to 12% SOC in a stirred H-type glass cell with graphite felt electrodes and Neosepta anion exchange membrane. Electrode potentials were measured relative to an AgQRE. Positive electrode: [Ru(acac) 3 ] [Ru(acac) 3 ] + + e Negative electrode: [Ru(acac) 3 ] + e [Ru(acac) 3 ] Fig. 5 shows the charging profile to 8.3% state of charge (SOC). A rapid rise in potential occurred beyond 7.6% SOC possibly due to a side reaction at the positive electrode or an increasing concentration overpotential. Note that the electrode potentials shown in Fig. 5 (and the subsequent charge/discharge data in Figs. 6 10) were measured relative to a silver quasi reference electrode and hence the absolute values of the potentials are not meaningful. However, the electrode potential data show that the rapid rise in the cell potential was associated with an increase in potential at the positive electrode. Fig. 6 shows the discharge of the charged Ru(acac) 3 at a constant current of 0.5 ma. The variations in the potential were due to the addition of solvent during the long discharging process (to make-up for solvent evaporation). With the low currents used, a relatively high energy Fig. 8. Potential time profile during discharge of the recharged 0.1 M Ru(acac) 3 and 1 M TEABF 4 in acetonitrile at 0.5 ma in a stirred H-type glass cell with graphite felt electrodes and Neosepta anion exchange membrane. Electrode potentials were measured relative to an AgQRE. efficiency was obtained, as shown in Table 2. The open circuit potential was significantly lower than obtained in the vanadium system, although the state of charge was much lower in this case. During discharge the cell potential was relatively low, at around 0.7 V. Fig. 6. Potential time profile during discharge of the charged 0.1 M Ru(acac) 3 with 1 M TEABF 4 in acetonitrile in a stirred H-type glass cell with graphite felt electrodes and Neosepta anion exchange membrane at a constant current of 0.5 ma. Electrode potentials were measured relative to an AgQRE. Fig. 9. Potential time profile during charging of 0.05 M [Fe(bpy) 3 ](ClO 4 ) 2 and 0.5 M TEAP in acetonitrile at 0.5 ma to 3% SOC in a stirred H-type glass cell with graphite felt electrodes and Neosepta anion exchange membrane. Electrode potentials were measured relative to an AgQRE.

6 2194 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) Table 2 A comparison of the key results obtained from the charge/discharge experiments conducted on Ru(acac) 3 and [Fe(bpy) 3 ](ClO 4 ) 2 in a stirred H-type glass cell Electrolyte OCP after charge (V) Energy efficiency (%) Cell energy output (mwh) 0.1 M Ru(acac) 3 in 1 M TEABF M Ru(acac) 3 in 1 M TEABF 4 (re-charge) M [Fe(bpy) 3 ](ClO 4 ) 2 in 0.5 M TEAP The electrolytes were re-charged at 1 ma constant current to a SOC of 12%, and the observed potentials are shown in Fig. 7, with the corresponding discharge at 0.5 ma shown in Fig. 8. Once again, a rapid increase in both cell and positive electrode potential was observed during charging, but this increase was delayed by around 200 min compared to the first charging. Although the reasons for this delay are unknown, it may have been due to an increase in solution temperature (resulting in a decrease in the activation overpotential) or an increase in the concentration of electroactive species due to solvent evaporation (decreased activation and concentration overpotentials). The delay could also have been caused by an increase in the rate of stirring, enhancing mass transport and thereby reducing the concentration overpotential. After charging to this higher state of charge, the open circuit potential rose to 1.3 V, however the cell voltage during discharge was again low. The reactions that are expected to occur at the positive and negative half-cells during charging and discharging of the [Fe(bpy) 3 ](ClO 4 ) 2 system are: Positive electrode: Fe(bpy) 3 2+ Fe(bpy) e Negative electrode: Fe(bpy) e Fe(bpy) 3 + The charge/discharge profiles for this system in the H-type cell are shown in Figs. 9 and 10. A lower charging current of 0.5 ma was used, considering the lower concentration of the active species (0.05 M due to solubility limitations). As it was difficult to determine whether the cell potential was due to a reversible electrochemical reaction or a side reaction, charging was only carried out to 3% SOC. After charging was completed, the electrolytes were discharged at a constant current of 0.1 ma (Fig. 10). It was necessary to use this low discharge current since it was found that with higher currents the cell voltage fell rapidly to zero. It should be noted that the lower state of charge used will lead to a lower concentration of the active species during discharge, so the performance would be expected to be poorer. The electrolytes in each half-cell were topped up periodically during the experiments. The energy efficiency obtained was very low when compared to the performance obtained from the Ru(acac) 3 compound under similar conditions (see Table 2). In addition, the cell potential obtained was less stable, falling below 0.5 V after less than 200 min and subsequently falling to zero. The results suggest that the charged species were unstable in the electrolyte or that the current efficiencies were very low. Although relatively low efficiencies were observed with all systems, this is largely associated with the H-cell design. As has been reported for the vanadium system [7 11], much higher efficiencies can be achieved with a practical flow cell design. Quantitative comparison of the three systems is difficult since differences in each system necessitated the use of a range of operating conditions. However, the results suggest that the Ru(acac) 3 system is superior to the [Fe(bpy) 3 ](ClO 4 ) 2 system. Higher efficiencies were obtained with this system and around times more energy was recovered during discharge. Although the [Fe(bpy) 3 ](ClO 4 ) 2 offers the possibility of higher open circuit potentials and lower cost, low energy efficiencies were observed, probably due to low current efficiencies. Because of the potential advantages of the [Fe(bpy) 3 ](ClO 4 ) 2, it is recommended that the causes and possible remedies for the low current efficiencies be investigated. In addition, further studies with a flow cell with improved transport conditions and lower ohmic losses should be carried out to determine the viability of both systems. 4. Conclusions Fig. 10. Potential time profile during discharge of the charged 0.05 M [Fe(bpy)3](ClO4)2 and 0.5 M TEAP in acetonitrile solution at 0.1 ma in a stirred H-type glass cell with graphite felt electrodes and Neosepta anion exchange membrane. Electrode potentials were measured relative to an AgQRE. Metal organic species may offer high efficiency, high cell potential systems for redox flow battery applications. Results in a simple H-type cell indicate that high efficiencies can be achieved with a ruthenium acetylacetonate system, which has high solubility and stability in an acetonitrile electrolyte. Since the two redox couples revert to the same species on discharge, cross-over will not reduce the cycle life time and complex electrolyte reprocessing is not required. Charging the system generates species

7 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) with opposite charge, so some loss of efficiency will occur by transport through either cation or anion exchange membrane materials. However, a redox flow battery utilizing a low cost microporous membrane can be envisaged. Further evaluation in a flow cell is recommended. Acknowledgements Funding for this research was provided by the Engineering and Physical Sciences Research Council (EPSRC). The authors would also like to give special thanks to Dr. N. Stevens and Dr. C.H. Bae for their valuable input in this work. References [1] I.M. Ritchie, O.T. Siira, Eightth Biennial Congress of International Solar Energy, 1983, p [2] M. Bartolozzi, J. Power Sources 27 (1989) 219. [3] M. Skyllas-Kazacos, F. Grossmith, J. Electrochem. Soc. 134 (1987) [4] A. Paulenova, S.E. Creager, J.D. Navratil, Y. Wei, J. Power Sources 109 (2002) 431. [5] C.H. Bae, Ph.D. Thesis, University of Manchester Institute of Science and Technology, UK, [6] C.H. Bae, E.P.L. Roberts, R.A.W. Dryfe, Electrochim. Acta 48 (2002) 279. [7] E. Sum, M. Rychcik, M. Skyllas-Kazacos, J. Power Sources 16 (1985) 85. [8] E. Sum, M. Skyllas-Kazacos, J. Power Sources 15 (1985) 179. [9] M. Skyllas-Kazacos, M. Rychcik, R.G. Robins, A.G. Fane, M.A. Green, J. Electrochem. Soc. 133 (1986) [10] M. Skyllas-Kazacos, D. Kasherman, D.R. Hong, M. Kazacos, J. Power Sources 35 (1991) 399. [11] M. Skyllas-Kazacos, M. Rychick, R. Robins, US Patent 4,786,567, US, [12] T. Yamayura, N. Watanabe, Y. Shiokawa, J. Alloys Compd. 408 (2006) [13] D.A.J. Rand, R. Woods, R.M. Dell, Batteries for Electric Vehicles, John Wiley & Sons Inc., NY, [14] I. Tsuda, K. Nozaki, K. Sakuta, K. Kurokawa, Sol. Energy Mater. Sol. Cells 47 (1997) 101. [15] A. Shibata, K. Sato, Power Eng. J. 13 (1999) 130. [16] C. Fabjan, J. Garche, B. Harrer, L. Jorissen, C. Kolbeck, F. Philippi, G. Tomazic, F. Wagner, Electrochim. Acta 47 (2001) 825. [17] VRB Power Systems Inc., [18] Y.D. Chen, K.S.V. Santhanam, A.J. Bard, J. Electrochem. Soc. 128 (1981) [19] M. Morita, Y. Tanaka, K. Tanaka, Y. Matsuda, T. Matsumura-Inoue, Bull. Chem. Soc. Jpn. 61 (1988) [20] A. Pighin, B.E. Conway, J. Electrochem. Soc. 122 (1975) 619. [21] Y. Matsuda, K. Tanaka, M. Okada, Y. Takasu, M. Morita, T. Matsumura- Inoue, J. Appl. Electrochem. 18 (1988) 909. [22] A. Endo, Y. Hoshino, K. Hirakata, Y. Takeuchi, K. Shimizu, Y. Furushima, H. Ikeuchi, G.P. Sato, Bull. Chem. Soc. Jpn. 62 (1989) 709. [23] K. Itoh, K. Honda, M. Sukigara, J. Electroanal. Chem. 110 (1980) 277. [24] N. Tokel, A.J. Bard, J. Am. Chem. Soc. 94 (1972) [25] H. Kaneko, K. Nozaki, Y. Wada, T. Aoki, A. Negishi, M. Kamimoto, Electrochim. Acta 36 (1991) [26] S. Zhong, M. Skyllas-Kazacos, J. Power Sources 39 (1992) 1. [27] M. Gattrell, J. Park, B. MacDougal, J. Apte, S. McCarthy, C.W. Wu, J. Electrochem. Soc. 151 (2004) A123. [28] M. Rychcik, M. Skyllas-Kazacos, J. Power Sources 19 (1987) 45.