evaluation of electrolytes for redox flow battery applications 2007

7/25/2019 Evaluation of Electrolytes for Redox Flow Battery Applications 2007

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Electrochimica Acta 52 (2007) 21892195

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, UKb 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 celldesigns. 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 simpleH-type cell. Theall-vanadiumelectrolyte systemhas been developed as a commercial systemand was used asthe 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 metalorganic 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)2in 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

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

chargingor discharging [1,2]. Duetotheuseoftwosolubleredox

couples, solid-state reactions with their accompanying morpho-

logical changes at the electrodes are absent [3].Thus, there are

no fundamental cycle life limitations associated with these pro-

cesses 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 devel-

oped to date is the use of two separate redox species in the half-

cells, leading to the potential for cross-contamination of active

Corresponding author. Tel.: +44 161 306 8849; fax: +44 161 306 4399.

E-mail address: [email protected](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 dueto this problem [4]. To redress

this issue, an all-chromium redox electrolyte was investigated at

the University of Manchester and the charge/discharge charac-

teristics of a laboratory scale battery were reported[5,6].Prior

to this, other workers have performed extensive investigations

on the all-vanadium redox system [3,710] and patented the

technology[11].In addition, an all-neptunium system has been

evaluated[12],although the hazards of working with radioac-

tive electrolytesare likelyto limit thepractical application of this

system. Several prototype vanadium systems have been investi-

gated successfully[1316]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 opera-

tion 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 poten-

tial window, have been investigated in this study. In addition,

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
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2190 M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) 21892195

trolyte cross-contamination. One approach is to use a single

system which offers three oxidation states, so that thedischarged

species is the same on each side of the cell. Such a system would

have the advantage thatany cross-contaminationwould onlylead

to some self-discharge, and little or no cell balancing or elec-

trolyte 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 fol-

lowing properties, which are generally desirable for redox flow

batteries[18,19]:

fast kinetics at the electrodeelectrolyte 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 ace-

tonitrile electrolyte were selected which apparently offered fastelectrode 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(BF4)2] has exhibited fast kinetics[19].In addi-

tion, this system offers the possibility of cell voltages of upto 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 avail-

able 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 around2.3 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

([3,711]) 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 areneeded 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 pro-

vide 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(BF4)2] was prepared by addition of NaBF4 in ace-

tonitrile 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 elec-

trolytes (Caution: perchlorate salts are potentially explosive and

should be handled withappropriatecare). 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 0.005 wt%.

2.2. Cyclic voltammetry

A graphite rod (Goodfellow) of surface area 0.06 cm2 was

used as the working electrode for cyclic voltammogram exper-

iments with the vanadium battery electrolytes. A glassy-carbon

electrode (I.J. Cambria Scientific) of surface area 0.07 cm2 was

used for cyclic voltammetry in organic media. The electrodewas 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 elec-

trode (AgQRE). A platinum counter electrode was used in each

case.

Cyclic voltammetry was conducted using a standard three-

electrode cell, with a Autolab/PGSTAT30 potentiostat for poten-

tial control. All solutions were de-aerated prior to experiments.

The solution headspace was purged with inert gas for the dura-

tion of experiments.


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M.H. Chakrabarti et al. / Electrochimica Acta 52 (2007) 21892195 2191

Fig. 1. Schematic diagram of a glass cell apparatus for small-scale charge/

discharge tests of redox couples.

2.3. 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, wheretwice as much electrolyte (80 ml) is required in the anodic com-

partment. 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

30mm 15mm 10 mm. Graphite rods were used as current

collectors. An UltrexTM (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 therequired 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 elec-

trolytes, 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 beachieved. Discharge currents were selected to ensure that sig-

nificant cell voltages were obtained. For the vanadium system, a

dischargecurrent of 2 mA wasfoundto 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 andthe total cell potentialand

the potential of each electrode were monitored [relative to satu-

rated calomel electrode(SCE) or silverquasi reference electrode

(AgQRE)] throughout each experiment. Since a silver quasi ref-

erenceelectrode was used in theorganic electrolytesthe absolute

value of the electrode potentials is not meaningful, and the mea-

surements 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 vana-

dium system suggested that the kinetics of the V(II)/V(III)

couple were relatively fast, while the V(V)/V(IV) redox cou-

ple was found to be irreversible, consistent with results reported

in literature[2527].

All of the organic electrolyte systems studies demonstrated

reasonably fast (in most cases reversible) kinetics for both oxi-

dation 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 occur-

ring in the discharged state. Table 1 compares the species

against each other based on the possible open circuit poten-tial of a battery based on the system, their solubility, electro-

chemical 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(ClO4)2systems were selected for further study on the basis of their

superior solubility and fast kinetics. The relatively low cost of

the Fe(bpy)3(ClO4)2 makes this system particularly attractive.

Cyclic voltammograms for the Ru(acac)3and Fe(bpy)3(ClO4)2systems are shown inFigs. 2 and 3, illustrating the combination

of multiple redox couples with fast kinetics. The cyclic voltam-

mogram for Fe(bpy)3(ClO4)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 reductionsof 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[19,2224] 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 SigmaAldrich [i] and GFS Chemicals [ii].


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Fig. 2. Cyclic voltammograms recorded at 0.1 V s1 at a GC electrode in: (a)

2 mM Ru(acac)3 and 0.05M TEABF4 in acetonitrile and (b) 0.05M TEABF4in acetonitrile.

3.2. Charge/discharge of the vanadium redox system

During the first charging of the VOSO4electrolyte, 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 concentra-

tion 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 VOSO4 electrolyte are shown

below:

Positive half-cell [V(IV)/V(V)]:

VO2++H2O VO2++2H++ e

Negative half-cell [V(III)/V(II)]:

V3++ e V2+

Fig. 3. Cyclic voltammograms recorded at 0.1 V s1 at a GC electrode in: (a)

2 mM Ru(acac)3 and 0.05M TEABF4 in acetonitrile and (b) 0.05M TEABF4

in acetonitrile.

Fig.4. Charge/dischargepotentialtime profile of 0.1 M VOSO4solutionin 2 M

H2SO4 using graphite felt electrodes and UltrexTM anion exchange membrane.

Constantcharging current of 100mA for 300min followedby constantdischarge

at 2 mA constant current. Electrode potentials were measured relative to a SCE.

The charge/discharge of 0.1 M VOSO4was performed using

an UltrexTM AEM. Thevoltageprofile duringthe charge anddis-

charge (second cycle) is shown in Fig. 4. The cell was charged at100 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 chargepassed. 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 pro-

cess. The overall efficiency was found to be 5.6% (18.4% charge

efficiency and 31% voltage efficiency) and 0.092 Wh of energy

was recovered from the 80 ml of charged electrolyte. The low

overall efficiency obtained is a consequence of the high charg-

ing 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 dur-

ing 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 indi-

cate that the vanadium system can achieve high efficiencies, as

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 lowerohmic loss is associated with the use of the high conductivity

2 M H2SO4supporting electrolyte.

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

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

using graphite felt electrodes and a Neosepta anion exchange

membrane. The reactions occurring at the electrodes are shown

below:


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Fig. 5. Potentialtime profile during charging of 0.1 M Ru(acac)3 with 1M

TEABF4 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.

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 inFig. 5(and the subsequent charge/discharge

data in Figs. 610) were measured relative to a silver quasi refer-

ence electrode and hence theabsolute values of the potentials arenot 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. 6shows 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 lowcurrents used, a relatively high energy

Fig. 6. Potentialtime profile during discharge of the charged 0.1M Ru(acac)3with 1M TEABF4in 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. 7. Potentialtime profile during recharging of 0.1 M Ru(acac)3 and 1 M

TEABF4 in acetonitrile at a constant current of 1 mA to 12% SOC in a stirred


membrane. Electrode potentials were measured relative to an AgQRE.

Fig.8. Potentialtime profile during dischargeof the recharged0.1 M Ru(acac)3and 1 M TEABF4 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 inTable 2.The open circuit

potential was significantly lower than obtained in the vanadium

system, although the state of chargewas much lower in this case.

During discharge the cell potential was relatively low, at around

0.7 V.

Fig. 9. Potentialtime profile during charging of 0.05M [Fe(bpy)3](ClO4)2and

0.5M 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.


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Table 2

A comparison of the key results obtained from the charge/discharge experiments conducted on Ru(acac)3 and [Fe(bpy)3](ClO4)2in a stirred H-type glass cell

Electrolyte OCP after charge (V) Energy efficiency (%) Cell energy output (mWh)

0.1 M Ru(acac)3in 1 M TEABF4 0.85 74 11.0

0.1 M Ru(acac)3in 1 M TEABF4(re-charge) 1.30 57 12.0

0.05M [Fe(bpy)3](ClO4)2in 0.5 M TEAP 1.50 6 0.6

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 elec-

trode potential was observed during charging, but this increase

was delayed by around 200 min compared to the first charg-

ing. 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 evap-

oration (decreased activation and concentration overpotentials).

The delay could also have been caused by an increase in the rateof 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](ClO4)2system are:

Positive electrode:

Fe(bpy)32+

Fe(bpy)33+

+ e

Negative electrode:

Fe(bpy)32+

+ e Fe(bpy)3+

The charge/discharge profiles for this system in the H-type

cell are shown inFigs. 9 and 10.A lower charging current of

Fig. 10. Potentialtime profile during discharge of the charged 0.05 M

[Fe(bpy)3](ClO4)2and 0.5 M TEAPin acetonitrile solutionat 0.1 mA in a stirred


membrane. Electrode potentials were measured relative to an AgQRE.

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 dis-

charged at a constant current of 0.1 mA (Fig. 10).It was neces-

sary 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 elec-trolytes 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)3com-

poundundersimilar 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 sug-

gest 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[711],much higher

efficiencies can be achieved with a practical flow cell design.

Quantitative comparison of the three systems is difficult sincedifferences in each systemnecessitated theuse of a range of oper-

ating conditions. However, the results suggest that the Ru(acac)3system is superior to the [Fe(bpy)3](ClO4)2 system. Higher

efficiencies were obtained with this system and around 2030

timesmore energy was recovered duringdischarge. Although the

[Fe(bpy)3](ClO4)2 offers the possibility of higher open circuit

potentials and lower cost, lowenergy 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 becarried out to determine the viability of both systems.

4. Conclusions

Metalorganic species may offer high efficiency, high cell

potentialsystems forredoxflow battery applications. Results in a

simple H-typecell indicate that high efficiencies canbe achieved

with a ruthenium acetylacetonate system, which has high solu-

bility and stability in an acetonitrile electrolyte. Since the two

redox couples revert to thesame species on discharge, cross-over

will not reduce the cycle lifetime and complex electrolyte repro-

cessing is not required. Charging the system generates species


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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.

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