A study of the Fe(III)/Fe(II)–triethanolamine complex redox
couple for redox flow battery application
Y.H. Wen
a,b,∗, H.M. Zhang
a, P. Qian
a, H.T. Zhou
a, P. Zhao
a, B.L. Yi
a, Y.S. Yang
baFull Cell R&D Center, Dalian Institute of Chemical Physics, Chinese Academy of Sciences, Dalian, Liaoning 116023, China bResearch Institute of Chemical Defense, Beijing 100083, China
Received 14 July 2005; received in revised form 2 October 2005; accepted 27 October 2005 Available online 1 December 2005
Abstract
The electrochemical behavior of the Fe(III)/Fe(II)–triethanolamine(TEA) complex redox couple in alkaline medium and influence of the con-centration of TEA were investigated. A change of the concon-centration of TEA mainly produces the following two results. (1) With an increase of the concentration of TEA, the solubility of the Fe(III)–TEA can be increased to 0.6 M, and the solubility of the Fe(II)–TEA is up to 0.4 M. (2) In high concentration of TEA with the ratio of TEA to NaOH ranging from 1 to 6, side reaction peaks on the cathodic main reaction of the Fe(III)–TEA complex at low scan rate can be minimized. The electrode process of Fe(III)–TEA/Fe(II)–TEA is electrochemically reversible with higher reaction rate constant than the uncomplexed species. Constant current charge–discharge shows that applying anodic active materials of relatively high con-centrations facilitates the improvement of cell performance. The open-circuit voltage of the Fe–TEA/Br2cell with the Fe(III)–TEA of 0.4 M, after
full charging, is nearly 2.0 V and is about 32% higher than that of the all-vanadium batteries, together with the energy efficiency of approximately 70%. The preliminary exploration shows that the Fe(III)–TEA/Fe(II)–TEA couple is electrochemically promising as negative redox couple for redox flow battery (RFB) application.
© 2005 Elsevier Ltd. All rights reserved.
Keywords: Redox flow battery; Anolyte; Iron–triethanolamine complex; Bromine
1. Introduction
With increasing depletion of traditional fossil fuels and decreasing air quality, the development and utilization of renew-able energy sources are of great importance. For a stochastic supply of renewable energy such as solar or wind, reliable and inexpensive electricity storage is of crucial importance[1–3]. Of all the new energy storage technologies currently under devel-opment around the world, the redox flow battery (RFB) appears to offer great promise as a low cost, high-efficiency large-scale energy storage system[4,5].
Redox flow batteries are stationary storage batteries that operate by continually pumping two electrolytes past a pair of high-surface-area electrodes that are separated by an ioni-cally conductive spacer. Energy is stored and harvested via the oxidation/reduction reactions of redox-active solutes in the two electrolytes. It is desirable that the redox flow battery has
effi-∗Corresponding author.
E-mail address:wen yuehua@126.com (Y.H. Wen).
cient charge and discharge performance, a large storage capacity of electricity, and a high open circuit voltage. Therefore, the active materials are required to have two kinds of reversible or rapid redox couples, a large solubility in the electrolyte, a large potential difference between two couples. In addition, it is preferable that the electrode potentials of the two couples exist within the electrochemical potential window of the solvents[6]. Since the redox flow cell concept was first proposed, a number of redox flow batteries have been fabricated and developed. Up to now, more and more attentions have been paid to all-vanadium RFB[7–10].
The specific energy of the cell is mainly determined by redox potential differences and concentration of species in the elec-trolyte. For optimization of the all-vanadium RFB constitution of the electrolyte, its concentration influence on the redox species properties has to be considered[11–13]. But, the open-circuit voltage for each single cell after full charging is about 1.5 V, which is relatively low. With an aim of increasing the cell voltage further, the Ce(IV)/Ce(III) couple, which has a standard potential of 1.74 V, has been investigated in order to assess its suitability as positive active species in the catholyte of a RFB[4,10,11,15]. 0013-4686/$ – see front matter © 2005 Elsevier Ltd. All rights reserved.
J.G. Ibanez et al.[16]examined the electrochemistry of transi-tion metal complexes including iron–TEA for electrochemical applications by cyclic voltammetry. It was shown that the anodic peak potential of the Fe(III)–TEA was as low as−1.05 V (ver-sus SCE) on a mercury drop or Pt electrode, but the solubility of Fe(III)–TEA is limited to 0.15 M. In 1991, Bechtold et al.[17] investigated the iron–TEA complex as a mediator for the reduc-tion of the organic compounds due to its low redox potential. In addition, alkali stable ironII/III-complexes have been studied for the indirect cathodic reduction of vat dyes in different dyeing processes[14,18].
The present experimental work was performed with the aim of evaluating the Fe(III)/Fe(II)–triethanolamine(TEA) complex redox couple in alkaline electrolyte for use in a RFB. The Fe(III)/Fe(II)–TEA complex redox couple is attractive for RFB technology because of its very negative redox potential, which should result in a battery with a high cell voltage and thus a great energy storage capacity. For example, the vanadium RFB exhibits a cell voltage of approximately 1.3 V, whereas a cell based on the Fe(III)/Fe(II)–TEA complex and Br−/Br2 redox couple is predicted to have a cell potential of about 1.9 V, which is significantly larger.
2. Experimental 2.1. Reagents
Fe2 (SO4)3·7H2O and FeSO4,triethanolamine(TEA) obtained from Kermel chemicals of Tianjing, Bodi chemicals of Tianjing and Bazhoushi chemicals of Tianjing. Sodium hydroxide from Xinxi chemicals of Shenyang, Sodium chloride from Tianhe chemicals of Tainjing, were all analytical grade. All solutions were prepared with distilled water.
2.2. Preparation of redox couple solution
The iron Fe(III)/Fe(II) to ligand (TEA) ratio was from 1/1 to 1/10; the concentration of the complex was 50–600 mM, the supporting electrolyte was 0.4 mol/L aqueous sodium chlo-ride; 1–3 M NaOH was used to adjust the pH of the solution to about 13. The determination of the solubility requires achiev-ing the equilibrium between solution and solid state at a given temperature. Instead of several hours, the achievement of true equilibrium takes many days of shaking at constant tempera-ture. The experiments indicated that the equilibration time for the Fe–TEA system is about 7 days.
2.3. Apparatus and procedure
The pH of the solution was measured with a calibrated pH meter (Shanghai instruments PHS-25, China). The cyclic voltammogram was measured by the CHI660 electrochemical station (CH Corporation, USA). All solutions were purged with prepurified nitrogen prior to experimentation. The curves of cur-rent versus potential were recorded in a three-compartment cell, with a graphite rod (area, 0.103 cm2) as an inert working elec-trode, the auxiliary electrode was a large area graphite sheet
electrode. All potentials were expressed relative to an aqueous saturated calomel electrode (SCE), which was connected with the electrochemical cell through a salt bridge full of saturated potassium chloride solution.
Cell charge–discharge tests were performed on small test cells which employed graphite felt electrodes contacted against two graphite sheets that acted as current-collectors. The elec-trode and membrane areas were 5 and 6 cm2, respectively. The two half-cell electrolytes were separated by a sheet of Nafion 117 cation-exchange membrane in the Na+form to prevent the bulk mixing of the two solutions as they were pumped through the cell. Two Xishan pumps (China) were used to pump each half-cell electrolyte through the corresponding half-cell cav-ity where the charge–discharge reactions occurred. The anolyte was 50 ml of 2 M NaBr, while the catholyte was 50 ml of 0.2–0.5 M Fe(III)–TEA + 3 M NaOH in 0.4 M NaCl medium. Prior to charging the cell, the negative electrolyte was purged with nitrogen gas at 0.2 l min−1for 30 min and then sealed in the tank for electrolyte storage. The galvanostatic charge and dis-charge of the test cell were carried out using a CT2001A (Land, instruments) multi-channel generator. The applied current for the charge and discharge was 100 mA. While the equipment was running, the total cell potentials were monitored using a com-puter through the A/D boards. After full charging, open-circuit voltage of the battery was measured.
3. Results and discussion 3.1. Solubilities
Literature data[12]suggest that the solubility of Fe(III)–TEA complex is limited to 0.15 M when the metal to ligand ratio was 2, which is too low for use in the RFB. It is believed that the sol-ubility of the iron–TEA complex is expected to depend strongly on the concentrations of TEA and NaOH. To predict saturation and precipitation of iron species in alkaline electrolyte, a sys-tematic study of solubility of both ferric and ferrous complexes with TEA in sodium hydroxide was undertaken to determine the optimal conditions of TEA and NaOH concentration where the iron species are the most stable.
The saturation concentrations of ferric and ferrous complexes with TEA were determined at room temperature at different concentrations of TEA which ranged from 0.1 to 2 M, adding dif-ferent NaOH concentrations to adjust the PH of solution. Results are listed in Table 1. It can be seen that the solubility of the Fe(III)–TEA complex increases continuously with the TEA con-centration rising, as well as the concon-centration of NaOH added. But, at high TEA concentrations beyond 1.4 M, a large amount of NaOH precipitates from the solution if the Fe(III)–TEA complex concentration is higher than 0.5 M, leading to the unavailability of the Fe(III)–TEA complex in alkaline electrolyte. Thus, the maximum concentration of Fe(III)–TEA complex that is about 0.6 M when the TEA concentration is 1.4 M, adding about 3 M NaOH to adjust the pH. Also, the solubility of the Fe(II)–TEA complex can be up to 0.4 M with the TEA concentration increas-ing.
Table 1
Concentrations of iron in saturated solutions of Fe(III)–TEA and Fe(II)–TEA complexes at various concentrations of TEA and NaOH
TEA (mol/L) NaOH (mol/L) Fe(II)–TEA (mol/L) NaOH (mol/L) Fe(III)–TEA (mol/L) 0.1 – <0.1 1 0.1 0.3 2 0.1 2 0.2 0.6 3 0.15 3 0.35 0.8 3 0.20 3 0.5 1.0 3 0.25 3 0.55 1.4 3 0.3 3 0.6 2.0 3 0.4 3 0.5
3.2. Voltammetric behavior of Fe(III)/Fe(II)–TEA in alkaline medium
Fig. 1 shows the cyclic voltammograms of 0.05 M Fe(III)–TEA complex formed in TEA ligand solution with vari-ous concentrations containing 1 M NaOH, 0.4 M NaCl at a scan rate of 10 mV/s. FromFig. 1, it was found that the influence of TEA ligand concentration mainly concerns two aspects: (1) when the concentration of TEA is as low as 0.1 M, a small shoul-der can be observed on the cathodic peak, which diminishes with cycle number and TEA concentration increasing; (2) in low con-centration of TEA, the cathodic peak of this complex is distorted because of the side reaction which seems to be the hydrogen evolution. But, with cycling number rising, it decreases sharply. With the increase of TEA concentration, this side reaction peak on the cathodic process is gradually eliminated. Thus, this two irreversible redox process may be explained by the existence of complexes with different stabilities in the solution containing
Fig. 2. Cyclic voltametric curves for 0.05 M Fe(III)–TEA complex formed in TEA ligand solution containing 1 M NaOH, 0.4 M NaCl at a scan rate of 100 mV/s. Concentrations of the TEA ligand: (1) 0.1 M; (2) 0.15 M; (3) 0.27 M; (4) 0.40 M.
low concentration of TEA. Thus, stable Fe(III)–TEA complexes can be obtained in high concentration of TEA, namely, the molar ratio of Fe(III) ion to TEA is up to 8/1. In addition, the pre-peak and side reaction pre-peak in the cathodic main pre-peak disappear almost completely at higher scan rates such as 100 mV/s from the reduction of Fe(III)–TEA complex due to the slow reaction rate of the side reactions as shown inFig. 2.
Fig. 3shows cyclic voltammograms of 0.05 M Fe(III)–TEA complex formed in 0.4 M TEA solution at various concentrations of NaOH at a scan rate of 10 mV/s. It can be seen that in too low
Fig. 1. Cyclic voltametric curves for 0.05 M Fe(III)–TEA complex formed in TEA ligand solution containing 1 M NaOH, 0.4 M NaCl at a scan rate of 10 mV/s. Concentrations of the ligand TEA: (1) 0.1 M; (2) 0.15 M; (3) 0.27 M; (4) 0.40 M.
Fig. 3. Cyclic voltammograms for 0.05 M Fe(III)–TEA complex formed in 0.4 M TEA solution at a graphite electrode and at a scan rate of 10 mV/s, adding different concentrations of NaOH. Concentration of NaOH and the PH of the solution: (1) 0.5 M, 13.85; (2) 1 M, 13.96; (3) 1.5 M, 13.97; (4) 2.5 M, 13.78.
or too high concentration of NaOH, the irreversible side reaction peak on the cathodic process appears, especially when adding very low concentration of NaOH to adjust the pH. It is possibly due to the formation of complexes with different stabilities in the solution. Whatever the concentration of NaOH is, the pH of the solutions is between 13 and 14, indicating that the ligand – TEA acts as a PH buffer. Too low concentration of NaOH, namely the ratio of NaOH to TEA is below 1, cannot be enough for the stable Fe(III)–TEA complex to be formed, whereas too high concentration of NaOH, namely the ratio of NaOH to TEA is above 6, results in a slightly decrease in the PH of the solution, making the side reaction of hydrogen evolution tend to take place at lower potential. So, the adding amount of NaOH should be controlled by the ratio of NaOH to TEA, ranging from 1 to 6. 3.3. Kinetics of Fe(III)–TEA/Fe(II)–TEA electrochemical processes on graphite carbon electrodes
The electrochemical kinetics of the Fe(III)–TEA/Fe(II)–TEA redox couple at Pt or mercury drop electrodes have been studied by different authors[14]at room temperature.
Carbon electrodes are of technological importance and are widely employed in RFB technology; therefore, their use is the subject of our interest. In this work, the cyclic voltammetry with a graphite electrode is performed at room temperature with solu-tions of 0.05 M Fe(III)–TEA complex containing 1 M NaOH, 0.4 M NaCl at a range of potential scan rates () as shown in Fig. 4.
Fig. 4shows the typical characteristics of a reversible one-electron process, exhibiting anodic and cathodic peak potential
separation,Ep60 mV (at 25◦C) which almost does not change with potential scan rates. It indicates that the Fe(III)–TEA com-plex redox reaction is very fast. The dependence of peak currents for the oxidation and reduction processes of the Fe(III)–TEA complex on the square root of the scan rate is shown in Fig. 5. For both the reduction of Fe(III)–TEA to Fe(II)–TEA atE=−1.063 V versus SCE and for the oxidation process at E=−1.003 V versus SCE, a good linear response is obtained, indicative of relatively fast electrode kinetics and a diffusion con-trolled reaction. Also, the equal anodic and cathodic peak cur-rents can be observed, namely, theip,a/ip,cis 1, whereip,aandip,c
Fig. 4. Cyclic voltammograms recorded at a graphite electrode for 50 mM Fe(III)–TEA complex at different potential scan rates. Scan rates: (1) 10; (2) 50; (3) 100; (4) 200; (5) 300; (6) 400 mV/s.
Fig. 5. Dependence of peak currents for the oxidation and reduction processes of the Fe(III)–TEA complex on the square root of the scan rate.
Table 2
Comparison of kinetic constants for Fe(III)–TEA/Fe(II)–TEA couples with the uncomplexed species Ligands Medium E0 (V vs. SCE) K0 (cm/s×103) D (cm2/s×106) Aquoa 0.5 M H 2SO4 0.45 1.6 2.5 TEA 1 M NaOH −1.0 63 1.63 aReference[16].
are the anodic and cathodic peak current, respectively. It is fur-ther demonstrated that the electrode process of the Fe(III)–TEA complex is reversible.
For a reversible system, the heterogeneous electron transfer rate constants,k0, can be obtained from the variation ofEpwith by using methods developed by Nicholson[19,20], according to which it is possible to determine the redox kinetic potential values (), where
Ψ = k0
(πaD0)1/2 (1)
where a =nF/RT and D0 is the diffusion coefficient of the Fe(III) or Fe(II) ion in the medium employed. The diffusion coefficients of Fe(III) and Fe(II) are assumed to be the same and be determined from the equation of the peak currentsipaoripc.
Table 2 lists the diffusion coefficient of the uncomplexed Fe(III)/Fe(II) species [16]for comparison and the complexed
Fe(III)–TEA species. With theseD0and-values, thek0-value is obtained from the Eq. (1) and also listed in Table 2. It is shown that the coordinated Fe(III)–TEA species have smaller D0-values than the uncomplexed aquo-species. But, thek0-value of the complexed Fe(III)–TEA is almost one order of magnitude larger than that of the aquo-species, suggesting the RFB employ-ing the Fe(III)/Fe(II)–TEA as negative active species will have high voltage efficiency.
3.4. Charge–discharge performance of Fe–TEA/Br2redox
flow battery
Performance of a RFB employing the Br2/Br− couple as catholyte active species and the Fe(III)/Fe(II)–TEA complexes with different concentrations as anolyte ones was evaluated with constant current charge–discharge tests and open-circuit voltage measurements, respectively. The Fe(III)–TEA complexes with concentrations ranging from 0.2 to 0.5 M in 0.4 M NaCl and 3 M NaOH was employed to determine charge discharge perfor-mance of the Fe–TEA/Br2redox flow battery.
Fig. 6 presents charge-discharge cycling curves for the Fe–TEA/Br2RFB employing Fe(III)/Fe(II)–TEA with different concentrations as negative electrolyte and 2 M NaBr as positive electrolyte at a current density of 20 mA cm−2. FromFig. 6, it can be seen that the charge–discharge voltage of the bat-tery is very high, facilitating the redox flow batbat-tery to obtain large energy density. When the Fe(III)–TEA complex with low concentration (≤0.3 M) is employed as negative active species, the cell resistance of the battery is relatively large and the capacity is reduced quickly with cycling number increasing. With an increase in concentration of the Fe(III)–TEA complex, charge–discharge performance of the battery is improved with higher discharge capacity and stability. When the Fe(III)–TEA complex concentration is up to 0.4 M, the discharge capacity is almost constant after cycling for 5 times, indicating prefer-able charge–discharge performance. When the Fe(III)–TEA complex concentration increases to 0.5 M, charge–discharge performance of the battery starts to decay. It indicates that the optimal concentration of the Fe(III)–TEA complex is 0.4 M.
To investigate the reason why performance of the battery with low concentration of the Fe(III)–TEA complex is poor, the conductivity of the electrolytes of various concentrations is measured and listed inTable 3. It is found that the conductiv-ity of the electrolyte is reduced with the active ion and ligand concentration increasing due to an increase in the solution vis-Table 3
Comparison for the theoretical and actual capacity of the Fe–TEA/Br2test cellsa
Fe(III)–TEA (mol/L) Theoretical capacity (Qt) (mAh) Charge capacity (Qc) (mAh) Discharge capacity (Qd) (mAh) Loss ratio of capacity (r)b(%) Conductivity (ms/cm) 0.2 268 250 143 46.6 84.4 0.3 402 346 184 54.2 78.8 0.4 536 466 346 35.4 74.3 0.5 670 575 387 42.2 68.9
aThe charge/discharge capacity is the value of the first cycle. b r= (Q
Fig. 6. Charge–discharge cycling curves for the Fe–TEA/Br2 redox flow battery employing Fe(III)/Fe(II)–TEA with different concentrations as negative redox
couple. Concentration of Fe(III)–TEA: (1) 0.2 M; (2) 0.3 M; (3) 0.4 M; (4) 0.5 M. cosity. This indicates that the ohmic resistance of the cell would not be attributed to the poor cell performance for the battery with low concentration of the Fe(III)–TEA complex. Rather, it may be mainly due to concentration polarization that would be highest at the lower active material concentrations. In addi-tion, on one hand, the pH of the catholyte falls as the system is periodically cycled resulting in H+ ions diffusing and being transported electrically into the anolyte with a lowing of the PH of the anolyte leading to a decrease in the solubility of the Fe(III)–TEA complex. on the other hand, active species may transfer across the membrane. The cross-contamination would lead to a considerable decrease in the discharging voltage and capacity with cycling number. In the case of the concentrated Fe(III)–TEA complex, the influence of the acid and the diffu-sion of active species are relatively minimized possibly because of the high solution viscosity leading to an increase in ionic transferring resistance across the membrane. However, when the Fe(III)–TEA complex concentration increases to 0.5 M, the limited solubility of the Fe(II)–TEA complex leads to a decrease in charge–discharge performance of the battery.
Data obtained from the charge–discharge test of the cells are summarized inTables 3 and 4, where various cell efficiencies are the average values over 5 charge/discharge cycles. It is observed that compared with the theoretical capacity, the loss ratio of capacity is the lowest for the cell with the 0.4 M Fe-TEA com-plex indicative of optimal cell performance. Further,Table 4 presents that open circuit voltages of the RFB with different concentrations of the Fe(III)–TEA as negative active species are above 1.94 V after full charging. Though the conductivity of the electrolytes with low concentration active species is relatively high, the voltage and current efficiencies of the battery with the Fe(III)–TEA complex of low concentrations(≤0.3 M) are lower than that of the battery with the Fe(III)–TEA complex of high concentrations due to concentration polarization and relatively serious cross-mixing of active materials. For the battery with the Fe(III)–TEA complex of 0.4 M, the highest voltage efficiency of 84%, energy efficiency of approximately 70% together with an open-circuit voltage of almost 2.0 V can be obtained. This is attributed to the fast electrode reaction kinetics and very low electrode potential of the Fe(III)/Fe(II)–TEA complex.
Table 4
Open circuit voltage and efficiencies of the Fe–TEA/Br2test cells
Fe (III)–TEA (mol/L) Open circuit voltage (V) Current efficiencya(%) Voltage efficiencya(%) Energy efficiencya(%)
0.2 1.94 60.7 79.9 48.5
0.3 1.96 51.7 77.3 40.0
0.4 1.98 82.4 83.9 69.1
0.5 1.95 72.7 82.9 60.2
4. Conclusion
In this study, the voltammetric behaviors of the Fe(III)– TEA/Fe(II)–TEA as negative redox couples in the RFB, elec-trode kinetics and electrolytic solubility were investigated. The following conclusions can be drawn.
The electrode process of Fe(III)–TEA/Fe(II)–TEA is elec-trochemically reversible with higher electron transfer rate con-stant than the uncomplexed Fe(III)/Fe(II) couple. Relatively, high TEA concentration is favorable electrochemically for the Fe(III)–TEA/Fe(II)–TEA redox couple with no side reaction peaks on the cathodic main reaction at low scan rate. Apply-ing the ligand TEA with high concentration, the solubility of the Fe(III)–TEA complex as high as 0.6 M can be obtained. And, the solubility of the Fe(II)–TEA complex is also up to 0.4 M.
Results obtained from the charge–discharge performance of test cells employing Br2/Br− as catholyte active species and the Fe(III)/Fe(II)–TEA as anolyte ones, demonstrate that applying relatively high concentration active materials as anolyte ones facilitates the improvement of cell performance with the range of the solubility. Due to low concentration polarization and the cross-contamination inhibiting to some extent, the battery with the Fe(III)–TEA complex of 0.4 M could deliver a energy efficiency of approximately 70% at a current density of 20 mA cm−2, together with an open circuit cell voltage of approximately 2.0 V. Summarizing the preliminary experimental data, the Fe(III)–TEA/Fe(II)–TEA redox couple is an attractive system for use in a RFB. It has a formal potential of approximately −1.0 V, making it suitable for use as a negative redox couple with the positive Br2/Br−system which has a formal potential of approximately 1.0 V. However, further experiments are needed to increase the long-term stability and solubility of the electrolyte.
Acknowledgements
The authors are grateful to the scientific research and innova-tion fund of the knowledge innovainnova-tion program from the Chinese academy of science (No. K2002D3) for the financial support. References
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