Electrochimica Acta 52 (2007) 2189–2195

Evaluation of for 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 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](ClO4)2 in acetonitrile was found to exhibit poor charge/discharge performance. © 2006 Elsevier Ltd. All rights reserved.

Keywords: Redox flow battery; Vanadium; Chromium; Ru(acac)3; [Fe(bpy)3](ClO4)2

1. Introduction electrolytes by transport through the membrane. For example, there has been little recent interest in the development of the Redox flow batteries are electrochemical energy storage iron/chromium redox flow cell due to this problem [4]. Toredress devices that utilise the oxidation and reduction of two soluble this issue, an all-chromium redox electrolyte was investigated at redox couples for charging and discharging. They differ from the University of Manchester and the charge/discharge charac- conventional batteries in that the energy-bearing chemicals are teristics of a laboratory scale battery were reported [5,6]. Prior not stored within at the surface, but in separate liquid to this, other workers have performed extensive investigations reservoirs and pumped to the power converting device for either on the all-vanadium redox system [3,7–10] and patented the charging or discharging [1,2]. Due to the use of two soluble redox technology [11]. In addition, an all-neptunium system has been couples, solid-state reactions with their accompanying morpho- evaluated [12], although the hazards of working with radioac- logical changes at the are absent [3]. Thus, there are tive electrolytes are likely to limit the practical application of this no fundamental cycle life limitations associated with these pro- system. Several prototype vanadium systems have been investi- cesses such as shedding or shape changes, which usually occur gated successfully [13–16] and some systems are well on their in conventional storage batteries. way to commercial success [17]. Despite these advantages, the redox flow battery has not been Despite such achievements, batteries employing aqueous widely exploited to date. One disadvantage of the systems devel- electrolytes have a low energy content. The energy output from oped to date is the use of two separate redox species in the half- the battery is proportional to the potential window of opera- cells, leading to the potential for cross-contamination of active tion available from the background electrolyte. The operating potential window of aqueous electrolytes is limited due to water [5]. Organic electrolytes, which offer a wider poten- ∗ Corresponding author. Tel.: +44 161 306 8849; fax: +44 161 306 4399. tial window, have been investigated in this study. In addition, E-mail address: [email protected] (E.P.L. Roberts). species have been selected which minimize the effect of elec-

0013-4686/$ – see front matter © 2006 Elsevier Ltd. All rights reserved. doi:10.1016/j.electacta.2006.08.052 2190 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) 2189–2195 trolyte cross-contamination. One approach is to use a single give an indication of the reversibility of redox couples, further system which offers three oxidation states, so that the discharged experiments are needed to demonstrate that selected systems can species is the same on each side of the cell. Such a system would be used for energy storage. For example, a redox couple may be have the advantage that any cross-contamination would only lead reversible, but the charged species may be unstable over long to some self-discharge, and little or no ‘cell balancing’ or elec- timescales, which would not be detected by cyclic voltammetry. trolyte processing would be required. An approach whereby the In this study systems which were found to exhibit fast kinetics same cation is used but with different ligands on each side of the were tested for their charge/discharge performance in a simple cell has been suggested [18], but this has not been considered in H-type cell. These experiments aimed to determine whether the this study. selected systems could be used for energy storage and to pro- Electrolyte systems can be selected on the basis of the fol- vide a preliminary indication of the relative performance of each lowing properties, which are generally desirable for redox flow system. batteries [18,19]: 2. Experimental • fast kinetics at the electrode–electrolyte interface; • a relatively large open circuit potential; 2.1. Electrolytes • reasonable cost; • high solubility in the process electrolyte. Vanadium electrolytes were prepared from vanadium (IV) sulphate (>99.99% purity, Aldrich), with the V(II)/V(III) couple In this study, the following series of redox systems in ace- generated by electro-reduction. Sulphuric acid was used as the tonitrile electrolyte were selected which apparently offered fast background electrolyte. electrode kinetics (based on literature data, e.g. [19] and [20]) Reagent grade tris(2,2-bipyridine) ruthenium (II) chloride and the potential to operate with a single electrolyte using a is available from Aldrich. Since oxidation of the chloride salt species with three oxidation states. was known to be irreversible [24], the tetrafluouroborate salt [Ru(bpy)3(BF4)2] was prepared by addition of NaBF4 in ace- (i) Ruthenium organic complexes tonitrile and precipitation of NaCl. Ruthenium acetylacetonate  A number of ruthenium organic complexes which can [Ru(acac)3, 97% purity, Aldrich], tris(2,2 -bipyridine) iron(II) be both oxidized and reduced electrochemically have been perchlorate (reagent grade, GFS) and rubrene (reagent grade, reported in the literature, and some of these have been Aldrich) were used for the preparation of the respective elec- suggested as suitable candidates for a redox flow battery trolytes (Caution: perchlorate salts are potentially explosive and [21]. Tris(2,2-bipyridine) ruthenium (II) tetrafluoroborate should be handled with appropriate care). Tetraethyl-ammonium [Ru(bpy)3(BF4)2] has exhibited fast kinetics [19]. In addi- tetrafluoroborate and tetraethyl-ammonium perchlorate were tion, this system offers the possibility of cell voltages of up used as the background electrolyte. to 2.6 V, much higher than is possible in aqueous battery To remove dissolved oxygen, electrolytes were sparged for at systems [19]. Ruthenium acetylacetonate [Ru(acac)3] has least 10 min with oxygen-free dry argon (aqueous electrolytes) also been reported as offering fast oxidation and reduction or nitrogen (organic electrolytes). Water was removed from the kinetics [22] and a possible cell voltage of around 1.75 V. organic electrolytes using zeolite 4A (Merck) to a moisture level (ii) Tris(2,2-bipyridine) iron(II) perchlorate of below 0.005 wt%. This species can be oxidized and reduced [19] and offers a possible cell potential of 2.4 V. This compound is avail- 2.2. Cyclic voltammetry able commercially and is significantly cheaper than the ruthenium complexes. A graphite rod (Goodfellow) of surface area 0.06 cm2 was (iii) Rubrene used as the for cyclic voltammogram exper- Rubrene, a neutral organic species, can be oxidized and iments with the vanadium battery electrolytes. A glassy-carbon reduced electrochemically [23]. The redox potentials of electrode (I.J. Cambria Scientific) of surface area 0.07 cm2 was these reactions offer a possible cell potential of around used for cyclic voltammetry in organic media. The electrode 2.3 V. Again this compound is available commercially, was polished with alumina washed with de-ionised water and although it is significantly more expensive than the other acetone following the procedure described in literature [21]. redox species. The reference electrode used in aqueous solutions was the saturated calomel electrode along with a salt bridge. Organic These systems are compared to the all vanadium redox flow media required the use of a silver wire quasi-reference elec- battery system, which has previously been investigated in detail trode (AgQRE). A platinum counter electrode was used in each ([3,7–11]) and has been commercialized in recent years [17]. case. In this system, vanadium in four different oxidation states is Cyclic voltammetry was conducted using a standard three- used: V(II)/V(III) at the negative electrode and V(IV)/V(V) at electrode cell, with a Autolab/PGSTAT30potentiostat for poten- the positive electrode. tial control. All solutions were de-aerated prior to experiments. Each system was first tested by cyclic voltammetry in order The solution headspace was purged with inert gas for the dura- to evaluate the electrode kinetics. While cyclic voltammetry can tion of experiments. M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) 2189–2195 2191

nificant 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 satu- rated calomel electrode (SCE) or silver quasi reference electrode (AgQRE)] throughout each experiment. Since a silver quasi ref- erence electrode was used in the organic electrolytes the absolute Fig. 1. Schematic diagram of a glass cell apparatus for small-scale charge/ value of the electrode potentials is not meaningful, and the mea- discharge tests of redox couples. surements can only be used to monitor significant changes in the potential of each electrode. In aqueous experiments mass 2.3. Charge/discharge experiments transport was provided by means of sparging the solution with argon gas, while for organic solutions a magnetic stirrer was A schematic diagram of the H-type test cell is shown in used. Fig. 1. The constant current (galvanostatic) charge/discharge characteristics of the redox couples were used to evaluate their 3. Results and discussion performance in a prototype redox flow battery. Each electrolyte compartment contained 40 ml of electrolyte, 3.1. Cyclic voltammetry except for the initial charging of the all-vanadium system, where twice as much electrolyte (80 ml) is required in the anodic com- The results of cyclic voltammetry experiments for the vana- partment. After the first charging of the all-vanadium half of the dium system suggested that the kinetics of the V(II)/V(III) electrolyte in the anodic compartment was removed [3]. couple were relatively fast, while the V(V)/V(IV) redox cou- Graphite felt electrodes (Sigratherm® GFA 10) were ple was found to be irreversible, consistent with results reported employed for charge/discharge experiments in the H-type in literature [25–27]. glass cell. The graphite felt electrode had dimensions of All of the organic electrolyte systems studies demonstrated 30 mm × 15 mm × 10 mm. Graphite rods were used as current reasonably fast (in most cases reversible) kinetics for both oxi- collectors. An UltrexTM (Membranes-International Ltd.) anion dation and reduction reactions. The results suggested that all exchange membrane was used for vanadium charge/discharge four systems could be oxidized or reduced, confirming their tests and a Neosepta® AHA membrane (Eurodia Industrie SA) suitability for a redox flow battery with the same species occur- was used for organic charge/discharge experiments in all cases. ring in the discharged state. Table 1 compares the species Membranes were pre-conditioned by exposing them to the against each other based on the possible open circuit poten- required test solution for at least 6 h prior to experiments. The tial of a battery based on the system, their solubility, electro- circular area of the membrane exposed to the electrolyte in the chemical reaction kinetics, and cost. The information on the cell had a diameter of 27.5 mm. kinetics was obtained from cyclic voltammetry experiments The charge/discharge experiments were carried out under and from literature data. The Ru(acac)3 and Fe(bpy)3(ClO4)2 constant current conditions. The current was selected on the systems were selected for further study on the basis of their basis of preliminary experiments. For the vanadium electrolytes, superior solubility and fast kinetics. The relatively low cost of which had a much higher conductivity than the organic elec- the Fe(bpy)3(ClO4)2 makes this system particularly attractive. trolytes, a high charging current of 100 mA was used. Although Cyclic voltammograms for the Ru(acac)3 and Fe(bpy)3(ClO4)2 this led to a high cell voltage during charging (∼4 V) with the systems are shown in Figs. 2 and 3, illustrating the combination likelihood of side reactions, the aim was to attempt to charge the of multiple redox couples with fast kinetics. The cyclic voltam- cell close to its maximum capacity. For the organic electrolytes, mogram for Fe(bpy)3(ClO4)2 indicates that the species can be much lower charging currents were used in order to minimize reduced at least twice, consistent with previous studies [19]. The ohmic losses and to evaluate whether efficient charging could be reduction reactions are presumed to be single electron reductions 2+ + achieved. Discharge currents were selected to ensure that sig- of Fe(bpy)3 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)3 1.77 High Reversible £12 [i] Ru(bpy)3(BF4)2 2.62 Poor Quasi-reversible £21 [i] [Fe(bpy)3](ClO4)2 2.41 Moderate Reversible £2 [ii] Rubrene 2.33 Poor Reversible £37 [i]

Costs were obtained from Sigma–Aldrich [i] and GFS Chemicals [ii]. 2192 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) 2189–2195

Fig. 4. Charge/discharge potential–time profile of 0.1 M VOSO4 solution in 2 M −1 TM Fig. 2. Cyclic voltammograms recorded at 0.1 V s at a GC electrode in: (a) H2SO4 using graphite felt electrodes and Ultrex anion exchange membrane. 2 mM Ru(acac)3 and 0.05 M TEABF4 in acetonitrile and (b) 0.05 M TEABF4 Constant charging current of 100 mA for 300 min followed by constant discharge in acetonitrile. at 2 mA constant current. Electrode potentials were measured relative to a SCE.

The charge/discharge of 0.1 M VOSO was performed using 3.2. Charge/discharge of the vanadium redox system 4 an UltrexTM AEM. The voltage profile during the charge and dis- charge (second cycle) is shown in Fig. 4. The cell was charged at During the first charging of the VOSO electrolyte, the V(IV) 4 100 mA for 300 min, and was discharged at 2 mA until the cell species must be reduced to V(II) at the and oxidized to voltage dropped to zero. A high charging current was used in V(V) at the . Consequently, as same electrolyte concentra- order to attempt to fully charge the cell, with around 2.3 times tion was used in each compartment, twice as much electrolyte the theoretical charge passed. This high charging current also led (80 ml) was used in the anodic compartment during the first to a relatively high voltage during charging (∼4 V). The open charging [9]. For subsequent cycles, equal volumes of electrolyte circuit voltage after charging was high at 1.61 V, and the cell were used in each compartment [V(III)/V(II) and V(IV)/V(V)]. voltage remained above 1 V during most of the discharge pro- The charge/discharge reactions for the second and subsequent cess. The overall efficiency was found to be 5.6% (18.4% charge charge/discharge cycles of the VOSO electrolyte are shown 4 efficiency and 31% voltage efficiency) and 0.092 Wh of energy below: was recovered from the 80 ml of charged electrolyte. The low overall efficiency obtained is a consequence of the high charg- Positive half-cell [V(IV)/V(V)]: ing current used. The charge recovered is around 86% of the theoretical capacity, so that significant side reactions must have 2+ + + − VO + H2O  VO2 + 2H + e occurred during charging. Furthermore the average voltage dur- ing discharge is around 1.2 V, >80% of the theoretical potential which can be achieved with the vanadium system. In spite of Negative half-cell [V(III)/V(II)]: the poor overall efficiencies obtained, the discharge results indi- cate that the vanadium system can achieve high efficiencies, as 3+ + −  2+ V e V expected [28]. Furthermore, the results compare favourably to the all- chromium 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 2MH2SO4 supporting electrolyte.

3.3. Charge/discharge of Ru(acac)3 and [Fe(bpy)3](ClO4)2 systems

Charge/discharge of electrolytes consisting of ruthenium acetylacetonate [Ru(acac)3] and tetraethylammonium tetraflu- oroborate (TEABF4) were carried out in the H-type glass cell ® Fig. 3. Cyclic voltammograms recorded at 0.1 V s−1 at a GC electrode in: (a) using graphite felt electrodes and a Neosepta anion exchange 2 mM Ru(acac)3 and 0.05 M TEABF4 in acetonitrile and (b) 0.05 M TEABF4 membrane. The reactions occurring at the electrodes are shown in acetonitrile. below: M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) 2189–2195 2193

Fig. 5. Potential–time profile during charging of 0.1 M Ru(acac)3 with 1 M Fig. 7. Potential–time profile during recharging of 0.1 M Ru(acac)3 and 1 M TEABF4 in acetonitrile using 1 mA constant current to 8.3% SOC in a stirred TEABF4 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 H-type glass cell with graphite felt electrodes and Neosepta® anion exchange membrane. Electrode potentials were measured relative to a AgQRE. 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 refer- Fig. 8. Potential–time profile during discharge of the recharged 0.1 M Ru(acac)3 and 1 M TEABF4 in acetonitrile at 0.5 mA in a stirred H-type glass cell with ence electrode and hence the absolute values of the potentials are graphite felt electrodes and Neosepta® anion exchange membrane. Electrode not meaningful. However, the electrode potential data show that potentials were measured relative to an AgQRE. the rapid rise in the cell potential was associated with an increase in potential at the positive electrode. Fig. 6 shows the discharge efficiency was obtained, as shown in Table 2. The open circuit of the charged Ru(acac)3 at a constant current of 0.5 mA. The potential was significantly lower than obtained in the vanadium variations in the potential were due to the addition of solvent system, although the state of charge was much lower in this case. during the long discharging process (to make-up for solvent During discharge the cell potential was relatively low, at around evaporation). With the low currents used, a relatively high energy 0.7 V.

Fig. 6. Potential–time profile during discharge of the charged 0.1 M Ru(acac)3 Fig. 9. Potential–time profile during charging of 0.05 M [Fe(bpy)3](ClO4)2 and with 1 M TEABF4 in acetonitrile in a stirred H-type glass cell with graphite felt 0.5 M TEAP in acetonitrile at 0.5 mA to 3% SOC in a stirred H-type glass electrodes and Neosepta® anion exchange membrane at a constant current of cell with graphite felt electrodes and Neosepta® anion exchange membrane. 0.5 mA. Electrode potentials were measured relative to an AgQRE. Electrode potentials were measured relative to an AgQRE. 2194 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) 2189–2195

Table 2 A comparison of the key results obtained from the charge/discharge experiments conducted on Ru(acac)3 and [Fe(bpy)3](ClO4)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 TEABF4 0.85 74 11.0 0.1 M Ru(acac)3 in 1 M TEABF4 (re-charge) 1.30 57 12.0 0.05 M [Fe(bpy)3](ClO4)2 in 0.5 M TEAP 1.50 6 0.6

The electrolytes were re-charged at 1 mA constant current 0.5 mA was used, considering the lower concentration of the to a SOC of 12%, and the observed potentials are shown in active species (0.05 M due to solubility limitations). As it was Fig. 7, with the corresponding discharge at 0.5 mA shown in difficult to determine whether the cell potential was due to a Fig. 8. Once again, a rapid increase in both cell and positive elec- reversible electrochemical reaction or a side reaction, charging trode potential was observed during charging, but this increase was only carried out to 3% SOC. was delayed by around 200 min compared to the first charg- After charging was completed, the electrolytes were dis- ing. Although the reasons for this delay are unknown, it may charged at a constant current of 0.1 mA (Fig. 10). It was neces- have been due to an increase in solution temperature (resulting sary to use this low discharge current since it was found that in a decrease in the activation overpotential) or an increase in with higher currents the cell voltage fell rapidly to zero. It the concentration of electroactive species due to solvent evap- should be noted that the lower state of charge used will lead oration (decreased activation and concentration overpotentials). to a lower concentration of the active species during discharge, The delay could also have been caused by an increase in the rate so the performance would be expected to be poorer. The elec- of stirring, enhancing mass transport and thereby reducing the trolytes in each half-cell were topped up periodically during the concentration overpotential. After charging to this higher state experiments. The energy efficiency obtained was very low when of charge, the open circuit potential rose to 1.3 V, however the compared to the performance obtained from the Ru(acac)3 com- cell voltage during discharge was again low. pound under similar conditions (see Table 2). In addition, the cell The reactions that are expected to occur at the positive potential obtained was less stable, falling below 0.5 V after less and negative half-cells during charging and discharging of the than 200 min and subsequently falling to zero. The results sug- [Fe(bpy)3](ClO4)2 system are: gest that the charged species were unstable in the electrolyte or that the current efficiencies were very low. Positive electrode: Although relatively low efficiencies were observed with all systems, this is largely associated with the H-cell design. As 2+ 3+ − Fe(bpy)3  Fe(bpy)3 + e has been reported for the vanadium system [7–11], much higher efficiencies can be achieved with a practical flow cell design. Negative electrode: Quantitative comparison of the three systems is difficult since differences in each system necessitated the use of a range of oper- 2+ − + ating conditions. However, the results suggest that the Ru(acac)3 Fe(bpy)3 + e  Fe(bpy)3 system is superior to the [Fe(bpy)3](ClO4)2 system. Higher efficiencies were obtained with this system and around 20–30 times more energy was recovered during discharge. Although the The charge/discharge profiles for this system in the H-type [Fe(bpy) ](ClO ) offers the possibility of higher open circuit cell are shown in Figs. 9 and 10. A lower charging current of 3 4 2 potentials and lower cost, low energy efficiencies were observed, probably due to low current efficiencies. Because of the poten- tial advantages of the [Fe(bpy)3](ClO4)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

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 solu- bility and stability in an acetonitrile electrolyte. Since the two 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 redox couples revert to the same species on discharge, cross-over H-type glass cell with graphite felt electrodes and Neosepta® anion exchange will not reduce the cycle life time and complex electrolyte repro- membrane. Electrode potentials were measured relative to an AgQRE. cessing is not required. Charging the system generates species M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) 2189–2195 2195 with opposite charge, so some loss of efficiency will occur by [10] M. Skyllas-Kazacos, D. Kasherman, D.R. Hong, M. Kazacos, J. Power transport through either cation or anion exchange membrane Sources 35 (1991) 399. materials. However, a redox flow battery utilizing a low cost [11] M. Skyllas-Kazacos, M. Rychick, R. Robins, US Patent 4,786,567, US, 1988. microporous membrane can be envisaged. Further evaluation in [12] T. Yamayura, N. Watanabe, Y. Shiokawa, J. Alloys Compd. 408 (2006) a flow cell is recommended. 1260. [13] D.A.J. Rand, R. Woods, R.M. Dell, Batteries for Electric Vehicles, John Acknowledgements Wiley & Sons Inc., NY, 1994. [14] I. Tsuda, K. Nozaki, K. Sakuta, K. Kurokawa, Sol. Energy Mater. Sol. Cells 47 (1997) 101. Funding for this research was provided by the Engineering [15] A. Shibata, K. Sato, Power Eng. J. 13 (1999) 130. and Physical Sciences Research Council (EPSRC). The authors [16] C. Fabjan, J. Garche, B. Harrer, L. Jorissen, C. Kolbeck, F. Philippi, G. would also like to give special thanks to Dr. N. Stevens and Dr. Tomazic, F. Wagner, Electrochim. Acta 47 (2001) 825. C.H. Bae for their valuable input in this work. [17] VRB Power Systems Inc., http://www.vrbpower.com/, 2006. [18] Y.D. Chen, K.S.V. Santhanam, A.J. Bard, J. Electrochem. Soc. 128 (1981) 1460. References [19] M. Morita, Y. Tanaka, K. Tanaka, Y. Matsuda, T. Matsumura-Inoue, Bull. Chem. Soc. Jpn. 61 (1988) 2711. [1] I.M. Ritchie, O.T. Siira, Eightth Biennial Congress of International Solar [20] A. Pighin, B.E. Conway, J. Electrochem. Soc. 122 (1975) 619. Energy, 1983, p. 1732. [21] Y. Matsuda, K. Tanaka, M. Okada, Y. Takasu, M. Morita, T. Matsumura- [2] M. Bartolozzi, J. Power Sources 27 (1989) 219. Inoue, J. Appl. Electrochem. 18 (1988) 909. [3] M. Skyllas-Kazacos, F. Grossmith, J. Electrochem. Soc. 134 (1987) 2950. [22] A. Endo, Y. Hoshino, K. Hirakata, Y. Takeuchi, K. Shimizu, Y. Furushima, [4] A. Paulenova, S.E. Creager, J.D. Navratil, Y. Wei, J. Power Sources 109 H. Ikeuchi, G.P. Sato, Bull. Chem. Soc. Jpn. 62 (1989) 709. (2002) 431. [23] K. Itoh, K. Honda, M. Sukigara, J. Electroanal. Chem. 110 (1980) 277. [5] C.H. Bae, Ph.D. Thesis, University of Manchester Institute of Science and [24] N. Tokel, A.J. Bard, J. Am. Chem. Soc. 94 (1972) 2862. Technology, UK, 2001. [25] H. Kaneko, K. Nozaki, Y. Wada, T. Aoki, A. Negishi, M. Kamimoto, Elec- [6] C.H. Bae, E.P.L. Roberts, R.A.W. Dryfe, Electrochim. Acta 48 (2002) 279. trochim. Acta 36 (1991) 1191. [7] E. Sum, M. Rychcik, M. Skyllas-Kazacos, J. Power Sources 16 (1985) 85. [26] S. Zhong, M. Skyllas-Kazacos, J. Power Sources 39 (1992) 1. [8] E. Sum, M. Skyllas-Kazacos, J. Power Sources 15 (1985) 179. [27] M. Gattrell, J. Park, B. MacDougal, J. Apte, S. McCarthy, C.W. Wu, J. [9] M. Skyllas-Kazacos, M. Rychcik, R.G. Robins, A.G. Fane, M.A. Green, J. Electrochem. Soc. 151 (2004) A123. Electrochem. Soc. 133 (1986) 1057. [28] M. Rychcik, M. Skyllas-Kazacos, J. Power Sources 19 (1987) 45.