An inexpensive aqueous flow battery for large scale electrical energy storage based on water-soluble organic

Journal of The Electrochemical Society, 161 (9) A1371-A1380 (2014) A1371
An Inexpensive Aqueous Flow Battery for Large-Scale Electrical
Energy Storage Based on Water-Soluble Organic Redox Couples
Bo Yang,∗
Lena Hoober-Burkhardt,∗
Fang Wang,a
G. K. Surya Prakash,∗∗
and S. R. Narayanan∗∗∗,z
Loker Hydrocarbon Research Institute, Department of Chemistry, University of Southern California, Los Angeles,
California 90089, USA
We introduce a novel Organic Redox Flow Battery (ORBAT), for meeting the demanding requirements of cost, eco-friendliness,
and durability for large-scale energy storage. ORBAT employs two different water-soluble organic redox couples on the positive
and negative side of a flow battery. Redox couples such as quinones are particularly attractive for this application. No precious
metal catalyst is needed because of the fast proton-coupled electron transfer processes. Furthermore, in acid media, the quinones
exhibit good chemical stability. These properties render quinone-based redox couples very attractive for high-efficiency metal-free
rechargeable batteries. We demonstrate the rechargeability of ORBAT with anthraquinone-2-sulfonic acid or anthraquinone-2,6-
disulfonic acid on the negative side, and 1,2-dihydrobenzoquinone- 3,5-disulfonic acid on the positive side. The ORBAT cell uses a
membrane-electrode assembly configuration similar to that used in polymer electrolyte fuel cells. Such a battery can be charged and
discharged multiple times at high faradaic efficiency without any noticeable degradation of performance. We show that solubility and
mass transport properties of the reactants and products are paramount to achieving high current densities and high efficiency. The
ORBAT configuration presents a unique opportunity for developing an inexpensive and sustainable metal-free rechargeable battery
for large-scale electrical energy storage.
© The Author(s) 2014. Published by ECS. This is an open access article distributed under the terms of the Creative Commons
Attribution 4.0 License (CC BY, http://creativecommons.org/licenses/by/4.0/), which permits unrestricted reuse of the work in any
medium, provided the original work is properly cited. [DOI: 10.1149/2.1001409jes] All rights reserved.
Manuscript submitted March 18, 2014; revised manuscript received June 6, 2014. Published June 19, 2014. This was Paper 468
presented at the San Francisco, California, Meeting of the Society, October 27–November 1, 2013.
The integration of electrical energy generated from solar and wind
power into the grid is faced with the challenge of intermittent electric-
ity output from these sources. This challenge can be effectively met
by storing the electricity during times of excess production and re-
leasing the electrical energy to the grid during times of peak demand.
Rechargeable batteries are very attractive for energy storage because
of their high energy efficiency and scalability.1–3
Since grid-scale elec-
trical energy storage requires hundreds of gigawatt-hours to be stored,4
the batteries for this application must be inexpensive, robust, safe and
sustainable. None of today’s mature battery technologies meet all of
these requirements. The vanadium redox flow battery is one such bat-
tery technology that has reached an advanced level of development for
grid-scale applications.5
However, the limited resources of vanadium,
the high expense associated with the cell materials, and the toxicity
hazard of using large quantities of soluble vanadium, have been the
major challenges to the widespread adoption of the vanadium redox
flow battery.2,6,7
Aiming to overcome these disadvantages, we have
demonstrated for the first time an aqueous redox flow battery that
uses water-soluble organic redox couples at both electrodes that are
metal-free. Such a battery has the potential to meet the demanding
cost, durability, eco-friendliness, and sustainability requirements for
grid-scale electrical energy storage. We have termed this battery an
Organic Redox Flow Battery (ORBAT).
In ORBAT, two different aqueous solutions of water-soluble or-
ganic redox substances such as quinones are circulated past elec-
trodes. The positive and negative electrodes are separated by a
proton-conducting polymer electrolyte membrane (Figure 1). In the
schematic, the positive electrode uses a solution of 1,2-benzoquinone-
3,5-disulfonic acid (BQDS) and the negative electrode uses a solution
of anthraquinone-2-sulfonic acid (AQS). By choosing the appropriate
organic redox couples for the positive and negative electrodes, we
project that a cell voltage as high as 1.0 V can be achieved.8,9
The
quinones have a charge capacity in the range of 200–490 Ah/kg, and
cost about $5–10/kg or $10–20/kWh, leaving ample scope for achiev-
ing the US Department of Energy’s target of $100/kWh for the entire
∗
Electrochemical Society Student Member.
∗∗
Electrochemical Society Active Member.
∗∗∗
Electrochemical Society Fellow.
a
Present address: Department of Chemistry, Massachusetts Institute of Technology,
77 Massachusetts Avenue, Cambridge, Massachusetts 02139, USA.
z
E-mail: sri.narayan@usc.edu
battery system. ORBAT is unique in that it does not use any heavy
metals such as vanadium, chromium or zinc, and also avoids volatile
organic solvents such as those used in lithium batteries. Additionally,
the redox reactions will not require any expensive precious metal cata-
lysts or metallic electrodes. With the potential of being simultaneously
inexpensive and environmentally-friendly such a battery presents an
attractive solution for grid-scale energy storage.
While oxidation-reduction properties of organic compounds such
as quinones have been known to electrochemists for many years,
hardly a handful of these material pairs have been exploited for en-
ergy storage, especially as rechargeable batteries. Furthermore, to
our knowledge none of these compounds have been studied for a
large-scale energy storage system. In 2009, Xu and Wen studied a
battery that uses 1,2-benzoquinone disulfonic acid as a positive elec-
trode solution and combined it with a conventional lead sulfate neg-
ative electrode.10
We made a public announcement of the idea of
a battery with a dual-aqueous feed of water soluble organic redox
couples as part of the ARPA-E project press release in March 2013
and presented technical results on the operation of such a battery
at the Electrochemical Society Meeting in October 2013.11,12
More
recently, in January 2014, Aziz et al. reported the use of aqueous so-
lutions of anthraquinone-2,7-disulfonic acid at the negative electrode
in a redox flow battery with a bromine flow electrode on the posi-
tive side and highlighted the possibility of avoiding heavy metals.13
These quinone/bromine flow cells perform at high current densities of
1 A/cm2
demonstrating the possibility of using the organic redox cou-
Figure 1. Schematic of aqueous organic redox flow battery (ORBAT).
) unless CC License in place (see abstract).ecsdl.org/site/terms_useaddress. Redistribution subject to ECS terms of use (see198.27.89.56Downloaded on 2014-10-20 to IP

A1372 Journal of The Electrochemical Society, 161 (9) A1371-A1380 (2014)
ples for achieving high power densities. However, the bromine-based
positive electrode presents some challenges with respect to handling
and use of bromine, and also bromine crossover from the positive
to the negative electrode. The use of a water-soluble organic redox
couple at the positive electrode that we have studied here has the clear
prospect of overcoming the foregoing problems. Non-aqueous solu-
tions of organic redox substances have been considered by Rasmussen
and Brushett.14,15
In this publication, we focus on understanding the
performance characteristics of a dual-aqueous feed organic flow bat-
tery and discuss the materials, challenges, and directions for further
research in this area.
Factors Governing the Operation of ORBAT
Charge-transfer processes.— The principal basis of an efficient
rechargeable redox flow battery is the rapid kinetics of charge-
transfer at the positive and negative electrodes. Many organic redox
couples, especially from the quinone family, undergo rapid proton-
coupled electron transfer without the need for dissociating high en-
ergy bonds. Consequently, these redox couples have relatively high
rate constants for the charge-transfer process. In general, we expect
molecules with conjugated carbon-carbon bonds and keto- and enol-
groups that allow for the delocalization and rearrangement of the
pi-electrons to undergo these redox transformations with extraordi-
nary facility.16
Typical organic redox couples are 1,2-benzoquinone-
3,5-disulfonic acid (BQDS) and anthraquinone-2-sulfonic acid (AQS)
(Eqs. 1 and 2).
HO3S SO3H
OH
OH
+ 2e-
+ 2H+
O
O
HO3S SO3H
E0
= + 0. 85 V
[1]
O
O
SO3H
OH
OH
SO3H
+ 2e-
+ 2H+
E0
= + 0.09 V
[2]
Such quinone-based redox couples have rate constants that are 2–
3 orders of magnitude higher than that of the vanadium ions in the
commercial vanadium redox system.17,18
Overpotential losses from
charge-transfer are expected to be low with these organic redox cou-
ples. As dissociation and rearrangement of C-C and C-H bonds do not
occur in these electrochemical reactions, high-surface area conductive
metal-free electrode surfaces, such as those based on carbon black, are
sufficient to support the charge-transfer process. No precious metal
electro-catalyst is required. Selecting the appropriate compounds with
fast rate constants (on the order of 10−3
to 10−4
cm s−1
) is necessary
when considering the technical viability of ORBAT.
Cell voltage.— The standard reduction potential of the redox cou-
ple is characteristic of the molecule and its specific substituent groups.
Since the standard reduction potential is also related to the electronic
energy of the molecular orbitals, the voltage of a redox flow cell is de-
termined by the difference in energy of the highest occupied molecular
orbital (HOMO) of redox couple used as the negative electrode mate-
rial and the lowest unoccupied molecular orbital (LUMO) of the redox
couple used as the positive electrode material. Previous experimen-
tal studies on quinones show that electron-withdrawing substituent
groups lower the energy levels of the HOMO and LUMO, while
electron-donating substituents raise the levels.19
Thus, we can use
substituent groups to selectively tune the standard reduction potential
of the quinone compounds to achieve the desired cell voltage. We may
use the Gibbs free-energy change for the reduction of the redox couple
with hydrogen, as estimated from quantum mechanical calculations,
to determine the standard reduction potential of the redox couples.
Being an aqueous battery, the voltage range for ORBAT is limited
by the oxygen evolution reaction at the positive electrode and the hy-
drogen evolution reaction at the negative electrode. Consequently, a
maximum cell voltage of 1.23 V is to be expected at room temperature.
However, by inhibiting the kinetics of the hydrogen evolution and oxy-
gen evolution reactions, we may achieve higher cell voltages. In this
respect, the non-aqueous systems have an advantage of being able to
provide a wider range of cell voltage. However, the inherent advantage
of lower cost and higher level of safety presented by aqueous systems
is particularly attractive for large-scale energy storage applications.
Mass transport processes.— To realize a low-cost battery, we must
be able to operate at high current densities without compromising volt-
age efficiency. With the rapid charge transfer processes in quinones,
the current density will be limited by mass transport of the reactants
and products. If a current density of 100 mA cm−2
is required at a
reactant concentration of 1 M, we may use the steady-state Nernst
diffusion layer model to estimate the mass transport coefficients to be
approximately 5 × 10−4
cm s−1
.20
Consequently, to maintain a cur-
rent density of 100 mA cm−2
with a redox couple that has a solubility
of 1 M and a diffusion coefficient of about 1 × 10−6
cm2
s−1
, the
solutions must be circulated past the electrodes so as to maintain a
thin diffusion layer of approximately 2 × 10−3
cm. Such a diffusion
layer thickness is in the practical range of values observed with rapid
circulation. Flow-through electrodes can lead to further reduction in
diffusion layer thickness.21,22
Thus, the diffusion coefficient and solu-
bility of the redox couple are principal properties for which the values
must be as high as possible for reaching the performance and cost
targets for large-scale energy storage applications.
In general, the un-substituted quinones exhibit limited solubility
in water. However, the solubility of the quinones can be increased
substantially by the incorporation of sulfonic acid and hydroxyl sub-
stituents. For example, benzoquinone has a solubility of 0.1 M, while
BQDS has a solubility of approximately 1.7 M at 25◦
C. Further
increases in solubility can be achieved by raising the temperature.
Therefore, achieving solubility values as high as 2 M is practical even
within the quinone family of compounds. Other organic redox couples
with high aqueous solubility include carboxylic acids and aromatic
heterocyclic compounds.8
Reactivity and long-term cycling.— Under acidic conditions, the
electro-reduction of quinones occurs often by a concerted proton trans-
fer and electron transfer process5
and no radical species are produced
as part of this redox process. Sometimes, the mechanisms could in-
volve sequential steps of protonation and electron transfer.23
Addi-
tionally, alkaline environments favor the formation of anion radicals
that tend to be reactive. Such formation of radicals in alkaline medium
is commonly encountered at about a pH of 9 and have lifetimes long
enough to be studied by electron spin resonance spectroscopy10,23–25
Another significant benefit of acidic over alkaline systems arises from
the absence of free cations besides the proton (or hydronium ion) in
the solutions. Thus, a proton exchange membrane electrolyte with
high ionic conductivity can be used. Inexpensive hydrocarbon mem-
branes such as polystyrenesulfonic acid, sulfonated polyetherether-
ketone and sulfonated polyethersulfone can also be used instead of
Nafion.26,27
Furthermore, the proton exchange membrane will inhibit
the transport of any anionic chemical species across the membrane
avoiding crossover of reactants from one side of the cell to another,
thereby avoiding self-discharge. Consequently, the long-cycle life re-
quirement of large-scale energy storage systems is more likely to be
realized with acidic systems.
For continuous operation at temperatures as high as 60◦
C the se-
lected organic compounds must have sufficient hydrolytic stability.
The quinones are generally quite stable in contact with oxygen. How-
ever, when the reduced form of the quinones (hydroquinones) in so-
lution come in contact with oxygen from air, the reduced form can be
re-oxidized to the quinone form. No permanent loss of material prop-

erties occurs as a result of such re-oxidation, and this process is just
like self-discharge. This self-discharge process can be prevented by
excluding access of the hydroquinones to oxygen. The quinones and
hydroquinones are generally stable at elevated temperatures of 60◦
C,
permitting faster diffusion to be achieved. In contrast, the vanadium
systems are very sensitive to such elevated temperature operation be-
cause of the numerous reactions between water and the vanadium oxo
ions. In some cases insoluble products such as vanadium pentoxide
are produced that negatively impact cycle life.5
Faradaic efficiency.— The potential for faradaic efficiency losses
are generally associated with the negative electrode in these types
of systems due to the proximity of the hydrogen evolution potential.
In the case of AQS, the standard reduction potential is about 100
millivolts positive to that of the normal hydrogen electrode (NHE).
Consequently, hydrogen evolution cannot occur readily during charge.
Additionally, the hydrogen evolution reaction is substantially inhib-
ited on carbon materials. This situation in ORBAT is to be contrasted
with the vanadium redox battery; the vanadium(III)/vanadium(II) cou-
ple operates at a potential approximately 0.350 volts negative to NHE
and hydrogen evolution occurs during normal charging of the battery.
Furthermore, the deposition of metallic impurities on the electrode fa-
cilitates hydrogen evolution and thus lowers the overall efficiency. By
avoiding the use of any soluble metal ions and only using carbon-based
electrodes ORBAT can achieve a faradaic efficiency close to 100%.
In the present study, we focus on understanding the performance
of ORBAT based on 1,2-benzoquinone-3,5-disulfonic acid (BQDS)
at the positive electrode with either anthraquinone-2-sulfonic acid
(AQS) or anthraquinone-2,6-disulfonic acid (AQDS) at the negative
electrode. To this end, we have measured the kinetic parameters for
charge-transfer, and diffusion coefficients for various redox couples in
acidic media. We have fabricated membrane-electrode assemblies and
measured the properties of flow cells over multiple cycles of charge
and discharge, and analysed these results in terms of the fundamental
properties discussed above.
Experimental
Electrochemical characterization.— Measurement of kinetic pa-
rameters and diffusion coefficients was conducted in a three-electrode
cell consisting of a rotating glassy carbon disk working electrode, a
platinum wire counter electrode, and a mercury/mercuric sulfate ref-
erence electrode (Eo
= +0.65 V). The quinones, in either the fully
reduced or fully oxidized form, were dissolved in 1 M sulfuric acid
to a concentration of 1 mM. The solutions were de-aerated and kept
under a blanket of argon gas throughout all the experiments. All mea-
surements were conducted in the potentiodynamic mode (Versastat
300 potentiostat) at a scan rate of 5 mV s−1
over a range of rotation
rates (500 rpm to 3000 rpm). Impedance measurements were also
made at each rotation rate. Cyclic voltammetry was conducted on a
static glassy carbon electrode at a scan rate of 50 mV s−1
.
Charge/discharge cycling of full ORBAT cells.— A flow cell was
constructed using fuel cell hardware that has graphite end plates (Elec-
trochem Inc.) and an electrode active area of 25 cm2
. The reactant was
circulated using peristaltic pumps (Masterflex) at a flow rate of approx-
imately 0.5–1.0 liter min−1
. Membrane electrode assemblies (MEA)
needed for the cell were fabricated in house using procedures similar
to those previously reported for direct methanol fuel cells.27
Specif-
ically, two sheets of carbon paper (Toray 030-non-teflonized) were
coated with an ink containing 0.1 g of Vulcan XC-72 carbon black
and 0.3 g of Nafion. The coated electrodes were hot pressed with a
Nafion 117 membrane to form a MEA. All full cell experiments were
carried out at 23◦
C. Two glass containers served as reservoirs for the
solutions of the redox couples. An argon flow was maintained above
these solutions to avoid reaction of the redox couples with oxygen.
The current-voltage characteristics of the cells were measured at vari-
ous states of charge. Charge/discharge studies were carried out under
constant current conditions.
Quantum mechanics-based calculations.— To determine E1/2 val-
ues, we used density functional theory to calculate the standard Gibbs
free energy change ( Go
) for the reduction of the oxidized form of
the redox couple. The calculations were performed at the B3LYP/
6-31+G(d,p) level of theory with thermal correction28
and implicit
consideration of water-solvation.29
The free energy correction for the
standard state of 1 atm in the gas phase and 1 M upon solvation was
applied, i.e. Gsolution = Ggas + 1.9 kcal · mol−1
at 298 K. Consid-
ering the lower pKa value of benzenesulfonic acid (pKa(sulfonic acid) =
–2.8),30
quinone sulfonic acid derivatives are expected to dissociate to
sulfonates in 1 M sulfonic acid aqueous solution. Go
was calculated
based on the reduction of quinone derivatives with H2. The stan-
dard electrode potential for the redox couple was deduced from Eo
=
– Go
/nF, where n is the number of protons involved in the reaction
and F is the Faraday constant.
Results and Discussion
Cyclic voltammetric measurements on AQS and AQDS show a sin-
gle step electrochemical reaction involving two electrons (Figures 2a
and 2b). Peak separations suggested AQDS was less kinetically re-
versible than AQS. The cyclic voltammograms for BQDS showed a
rapid oxidation step, but a slower reduction step (Figure 2c). The shape
of the reduction peak for BQDS suggested a possible slow chemical
step following electron transfer. Such a slow step is consistent with
the hydration process leading to the conversion of the hydroquinone to
1,4-benzoquinone-2-hydroxy-3,5-disulfonic acid, as reported by Xu
and Wen.10
The reversible potentials estimated from the anodic and
cathodic scans for the three compounds were in agreement with the
quantum mechanical calculations (Table I). The facile proton and elec-
tron transfer processes occurring on a glassy carbon electrode in the
absence of any catalyst confirmed an outer-sphere type of mechanism.
Linear sweep voltammetric measurements at a rotating disk elec-
trode at various rotation rates (Figure 3a–3c) showed that the limiting
current, Ilim, was found to depend linearly on the square root of the
rotation rate, ω, as per the Levich equation (Eq. 3).
Ilim = 0.62nF AD2/3
o ω1/2
ν−1/6
C∗
[3]
Where n is the number of electrons transferred, F, the Faraday
constant, A, electrode area, Do, the diffusion coefficient, ν, the kine-
matic viscosity of the solution and C*, the bulk concentration of the
reactants. For n = 2, an active electrode area of 0.1925 cm2
, and a
kinematic viscosity of the electrolyte of 0.01 cm2
s−1
, we were able
to evalulate the diffusion coefficient from the slope of the straight line
plots in Figure 3d.
To determine the kinetic parameters for the charge-transfer pro-
cess, namely the rate constant and the apparent transfer coeffcient, the
logarithm of the kinetic current (after correction for mass-transport
losses) was plotted against the observed overpotentials greater than
100 mV, where the Tafel equation is applicable (Eq. 4 and Figure 3e).
I
1 − I
Ilim
= Iex
CO
C∗
O
exp −
αnF(E − Erev)
RT
−
CR
C∗
R
exp
(1 − α)nF(E − Erev)
RT
[4]
Where I is the current, Ilim is the limiting current, Iex is the exchange
current density, CO and CR are the concentration of the oxidized and
reduced species at the surface of the electrode, CO* and CR* are
the bulk concentrations of the oxidized and reduced species, α is the
transfer coefficient, n is the number of electrons transferred, F is the
Faraday constant, E-Erev is the overpotential, R is the gas constant,
and T is the temperature.
The rate constant, ko, was obtained from the exchange current
density (Eq. 5).
ko = Iex /nF AC∗
[5]

Figure 2. Cyclic voltammograms at a scan rate of 5 mV s−1 on a glassy carbon electrode in 1 M sulfuric acid containing (a)1 mM anthraquinone-2-sulfonic acid,
(b)1 mM anthraquinone-2,6-disulfonic acid and (c)1 mM 1,2-benzoquinone-3,5-disulfonic acid.
Besides BQDS, AQS and AQDS, we have also measured the
current-overpotential curves for hydroquinone and hydroquinone sul-
fonic acid (see supplementary material). Solubility of anthraquinone
in 1 M sulfuric acid was too low to obtain any reliable data.
The half-wave potential values (Table I) are consistent with the
values reported in the literature for the various compounds tested.29,31
It is clear that the addition of aromatic rings has a marked effect
of lowering the standard reduction potential and half-wave potential.
The addition of sulfonic acid groups tends to increase the standard
reduction potential, which is consistent with the lowering of molecular
orbital energies by electro-withdrawing groups.
To understand the changes in the standard reduction potentials we
have used quantum mechanics to calculate the free-energy change in
the reaction of the oxidized form of the redox couple with hydrogen. If
Go
is the Gibbs free energy change under standard conditions, then
− Go
/nF is the standard electrode potential for the redox couple,
Table I. Standard reduction potentials for selected quinones.
Standard Reduction Potentials
Redox Couple Experimental (E1/2 values) vs. NHE Eo(formal) Theoretical
Hydroquinone 0.67 0.68 0.70
Hydroquinone sulfonic acid 0.82 0.70 0.77
1,2-benzoquinone -3,5-disulfonic acid 1.1 0.87 0.85
Anthraquinone Insoluble Insoluble 0.05
Anthraquinone-2-sulfonic acid 0.13 0.15 0.09
Anthraquinone-2,6-sulfonic acid 0.05 0.19 0.12

Figure 3. Linear sweep voltammetric data (scan rate of 5 mV s−1) at a glassy carbon rotating disk electrode for 1 mM concentration of BQDS (a), AQS (b) and
AQDS (c) at the rotation rates indicated. Electrode potentials are versus a mercury sulfate reference electrode (E0 = +0.65). (d) Levich plot of the square root of
rotation rate vs the limiting current for AQS(♦), BQDS ( ), and AQDS (o). (e) Mass transport-corrected current-voltage plot for BQDS, AQS, and AQDS.
where n is the number of protons involved in the reaction and the F
is the Faraday constant. The values of E1/2 from experiments follow
the trends predicted by the theoretical calculations (Table I). The
strong correlation between experimental and theoretical predictions
suggest that such free energy calculations can be used to predict the
trends in E1/2 values of the redox compounds prior to experimental
testing, potentially enabling the discovery of new redox couples by
this computational approach.
The values of diffusion coefﬁcients (Table II) are about an or-
der of magnitude smaller in aqueous solutions than in non-aqueous

Table II. Electrochemical properties of the redox couples determined from rotating disk electrode experiments. MSE refers to the mercury sulfate
reference electrode (Eo = +0.65 V).
Exchange Diffusion Transfer
E1/2 vs. MSE, Current Density, Coefficient, Coefficient, Constant,
Redox Couple (Volt) (A cm−2) (cm2 s−1) αn Solubility (cm s−1)
Hydroquinone 0.02 5.09E-5 5.03E-6 0.508 0.53 M 2.36E-3
Hydroquinone sulfonic acid 0.17 1.10E-5 4.28E-6 0.418 0.8 M 5.52E-4
1,2-benzoquinone -3,5-disulfonic acid 0.45 3.00E-6 3.80E-6 0.582 1 M 1.55E-4
Anthraquinone-2- sulfonic acid −0.52 1.96E-5 3.71E-6 0.677 0.2 M 2.25E-4
Anthraquinone 2,6-disulfonic acid −0.60 2.97E-6 3.40E-6 0.426 0.5 M 1.52E-4
solvents such as acetonitrile.32
In aqueous solutions, the observed ex-
tent of decrease in the values of diffusion coefficients with increase in
molecular mass is about 6 × 10−9
cm2
s−1
per unit of molecular mass.
This coefficient is an order of magnitude lower than that observed in
acetonitrile. Thus, besides the effect of molecular mass, the molecu-
lar diameters resulting from the solvation and the interaction of ionic
groups with water through hydrogen bonding have a significant effect
on the diffusion coefficient values in aqueous solutions.
Rate constants are within the range of values found widely in
the literature for quinones.33,34
As sulfonic acid groups are added to
the ring, the intra-molecular hydrogen bonding interactions in the
quinone molecules increase. This intra-molecular hydrogen bonding
plays a critical role in the rate limiting step of proton-coupled electron
transfer,35
due to the increased stability of the compound and increased
cleavage energy required for concerted proton and electron transfer.
This stability provides a competition between the resident hydro-
gen atom and the incoming proton for interaction with the carbonyl
oxygen. According to our calculations, hydroquinone sulfonic acid
preferentially adopts a conformation allowing the formation of intra-
molecular hydrogen bonding, which leads to a stabilization energy of
1.6 kcal mol−1
. Similarly, intra-molecular hydrogen bonding provides
extra stabilization of other hydroquinone sulfonic acid derivatives
(Eq. 6 and Eq. 7). Thus, the intra-molecular hydrogen bonding could
explain the lowering of the rate constants observed with the addition
of sulfonic acid groups.
[6]
[7]
The quinone-based redox systems have been extensively reported
in the literature36–38
and it is well known that these systems undergo
proton-coupled electron transfer. The rate constants for charge transfer
were generally quite high, at least an order of magnitude higher than
that observed for the vanadium redox couples.17
The value of the
transfer coefficients being close to 0.5 and the high values of rate
constants suggest an “outer-sphere” process.
While the rate constants for the various compounds were at least
one order of magnitude greater than that of vanadium system, the dif-
fusion coefficients were comparable to that of vanadium,18
making the
quinone redox couples very attractive from the standpoint of electrode
kinetics compared to the vanadium redox flow battery system.
The Nernst diffusion layer model allows us to estimate the limiting
current for the oxidation and reduction processes (Eq. 8).
Ilim = nFC∗ D
δ
[8]
Where Ilim is the limiting current density, n is the number of moles
of electrons transferred per mole of reactants, F is the Faraday con-
stant (96485 C mole−1
), C* is the concentration, DO is the diffusion
coefficient, and δ is the diffusion layer thickness.
For a diffusion layer thickness of 50 microns, a diffusion coeffi-
cient of 3.8 × 10−6
cm2
s−1
, and a bulk concentration of 0.2 M, we
predict from Eq. 8 a limiting current density at room temperature to
be approximately 30 mA cm−2
. Further increase in limiting current
density can be achieved by increasing the concentration of reactants,
reducing the diffusion layer thickness, and by increasing the diffu-
sion coefficient. Higher concentrations and diffusion coefficients are
achieved by raising the operating temperature while a lower diffu-
sion layer thickness can be achieved by increased convective mass
transport to the surface of the electrode.
We have operated flow cells with aqueous solutions of 0.2 M BQDS
at the positive electrode and 0.2 M AQS or 0.2M AQDS at the negative
electrode. In these cells the electrodes consisted of Toray paper coated
with high-surface area carbon black bonded to the Nafion membrane.
These cells did not show any noticeable change in capacity over at
least 12 cycles of repeated charge and discharge (Figure 4). This result
confirmed that the quinones in aqueous acid solution are chemically
stable to repeated cycling. The capacity realized at a current density of
10 mA cm−2
was over 90% of the capacity contained in the solutions.
The use of Toray paper electrodes on either side of the cell as current
collecting surfaces presented a barrier to convective transport, setting
the diffusion layer thickness to as high as 150 microns, reducing the
limiting current density and lowering the cell voltage.
Increasing the mass transport of reactants and products improved
the current density and cell voltage significantly. In one configuration
of the electrodes, the increase in mass transport was achieved by
punching nine equally-spaced holes 1 cm in diameter in the Toray
paper electrodes to allow the flow of redox active materials to shear
directly past the carbon black layer bonded to the membrane. The
increased current and voltage observed as a result of the change in the
access of the redox materials to the electrode (Figure 5) confirmed
that the kinetics of the electrode reactions are largely controlled by
the mass transport of the reactants and products.
The dependency of cell voltage on current density when measured
as a function of the state of charge of ORBAT confirmed that the mass
transport of reactants had a significant impact on the operating cell
voltage (Figure 6a). The power density of the cell decreased signifi-
cantly below 50% state-of-charge. The alternating current impedance
of the cells measured as a function of frequency also confirmed that
the mass transport limitations increased as the state-of-charge de-
creased. The slopes of the Nyquist plot at low frequencies (< 1 Hz)
were found to steadily increase with decreasing state-of-charge, while
the rest of the impedance spectrum remained almost unaffected
(Figure 6b), suggesting an increasing thickness of the diffusion lay-
ers with a decrease in the state-of-charge. Therefore, it is clear from

Figure 4. 12 charge and discharge curves with a 25 cm2
redox flow cell, 0.2 M BQDS and 0.2 M AQS, 1 M
sulfuric acid, charge and discharge at 200 mA, flow rate
of 0.5 liters min−1.
Figure 5. The effect of electrode structure on cell per-
formance: 9 equally spaced holes 1 cm in diameter were
punched in the Toray paper electrodes. 25 cm2 redox
flow cell, 0.2 M BQDS, 0.2 M AQS, 1 M sulfuric acid,
charge-discharge at 50 mA, flow rate 0.25 liter min−1
the results in Figure 6 that at all current densities, mass transport
limitations made a significant contribution to the overpotential losses.
In an effort to understand the results presented in Figure 6, we have
analysed the effect of state-of-charge on the current-voltage character-
istics using a simplified one-dimensional model (see Appendix). The
assumptions in this analysis are based on experimental findings from
RDE studies and flow cell studies that show that the charge-transfer
reactions are facile and that mass transport processes determine the
cell voltage during operation.
The analysis yields the following relationship between the ob-
served cell voltage and the discharge current as a function of state-of-
charge.
Vcell = E0
c −E0
a
RT
nF
ln
Q2
(1−Q)2
−2Id
RT
nF
1
nFmt AQCi −Id
+
1
nFmt A(1 − Q)Ci + Id
− Id Rohmcell [9]
Where Vcell is the cell voltage during discharge and Ecnot and Eanot
are the standard reduction potentials for the two redox couples used
at the cathode and anode, respectively. Id is the discharge current
and Q is the state-of-charge with values between 0 to 1. Cinitial is the
starting concentration of the reactants at 100% state-of-charge; Cinitial
is assumed in this analysis to be the same at both electrodes. A is the
area of the electrode, and mt is the mass transport coefficient defined
as the diffusion coefficient divided by the diffusion layer thickness.
R is the universal gas constant, F is the Faraday constant, T is the
temperature, and n is the number of electrons in the redox reaction.
Eq. 9 has been graphed (Figure 7) for various states-of-charge using
experimentally determined parameters for the BQDS and AQS system
(Table III). Comparison of Figure 7 with the experimental data in Fig-
ure 6 shows general agreement of the trend predicted by the analysis
with the observed experimental results; decrease in state-of-charge
resulted in a decrease of discharge current at any particular voltage,
leading to a significant reduction in discharge rate capability at low
states-of-charge. However, the experimental current-voltage curves
were nearly linear at the high states of charge and the experimental
values of cell voltage at low current values decreased substantially
with decreasing state-of-charge. These deviations from the analytical
expression suggest that there are additional resistance elements under
dynamic operating conditions that are not captured in the simplified
analysis. We list at least two other effects that can cause substantial
changes to the observed voltage:
1. Electro-osmotic drag of water molecules (estimated to be about 3
molecules per proton) occurs across the membrane during passage
of current. These water molecules either appear at or are removed
from the diffusion layer at each electrode causing changes to the
pH and concentration of reactants and products. These concen-
tration changes at the interface will contribute to a reduction in
cell voltage. For example, at the cathode during discharge, wa-
ter molecules could be added to the diffusion layer causing the
pH to increase and, consequently, the electrode potential to de-
crease. Correspondingly, water molecules will be removed from

Figure 6. 25 cm2 redox flow cell, 0.2 M BQDS, 0.2 M AQS, 1 M sulfuric
acid (a) Cell voltage-current density curves as a function of state-of-charge
(5% difference each run). (b) Impedance spectroscopy data from 10 kHz to
10 mHz on the cells at various states-of-charge.
Table III. Parameters used in the analysis of current-voltage
curves as a function of state-of-charge.
Parameter Value
Standard Reduction Potential of Cathode (Ecnot), V +0.45
Standard Reduction Potential of Anode (Eanot), V −0.52
Initial concentration of reactants (Cinitial) moles cm−3 2 E-4
Diffusion coefficient of cathode and anode reactants
and products (D) cm2 s−1
4E-6
Diffusion layer thickness (δ) cm 4E-3
Geometric Area of the Electrode (A), cm2 25
Number of electrons in the reaction (n) 2
Series equivalent resistance at impedance at 10 kHz
(Rohmcell), Ohm
0.05
the anode, causing the pH to decrease, the electrode potential to
increase, and the cell voltage to decrease.
2. At the anode, we use a solution of AQS at concentrations close
to the solubility limit (0.2 M). Consequently, at a low state-of-
charge when the oxidized form of AQS at the negative electrode
is present in high concentrations in the bulk of the solution, the
high rates of discharge would cause the solubility limits to be
exceeded at the surface of the negative electrode. This would
result in the precipitation of redox materials at the surface of the
electrode and with a significant reduction of the current. To avoid
such an abrupt drop in cell voltage at high current densities and
low states-of-charge, the solubility of the redox materials must be
high. Additionally, reducing the thickness of the diffusion layer
by using a flow-through electrode will increase the “saturation-
limited” current density.
The analysis helps us to quantify the variations in performance
that can result from changes to local mass-transport conditions at any
state-of-charge. The observed differences between the experimental
data and the predictions of simple analysis of the cell performance
also help us to identify the phenomena that are important to consider
for further design and modeling of redox flow cells.
When an aqueous solution of 0.2 M AQDS was used on the negative
side of the flow battery, the tests showed charge-discharge cycling
Figure 7. Simulation of cell voltage as a func-
tion of discharge current density (as per Eq. A5)
at various states-of-charge, as indicated on the
curves using parameters in Table III.

Figure 8. a) 12 charge and discharge curves with a 25 cm2 redox flow cell,
0.2 M BQDS, 0.2 M AQDS, 1 M sulfuric acid, charge and discharge at
200 mA, flow rate of 1 liters min−1 on a peristaltic pump. b) Cell voltage-
current density curves as a function of state-of-charge (as indicated by percent
stated, 5% difference each run) in a 25 cm2 redox flow cell, 0.2 M BQDS,
0.2 M AQDS, 1 M sulfuric acid.
stability similar to AQS. By operating at a higher pumping speed,
the cell voltages and capacity for the BQDS/AQDS cell could be
increased, consistent with the reduction of the voltage losses from
mass transport limitations (Figure 8a). The cell voltage and current
density did not drop off as quickly with state-of-charge as in the case
of AQS. The aqueous solubility limit of AQDS is about 0.5 M while
that of AQS is about 0.2 M. Consequently, even at very low states-of-
charge, the solubility limit was less likely to be exceeded with AQDS
than with AQS. Thus, higher solubility allows the cell voltage to be
maintained at a higher value with AQDS compared to AQS especially
at low states-of-charge. This difference in performance of AQDS and
AQS highlights the role of solubility limits on the rate capability
at various states-of-charge. Such findings motivate us to investigate
higher concentrations and temperatures with AQDS in future studies.
Conclusions
For the first time, we have demonstrated the feasibility of operating
an aqueous redox flow cell with reversible water-soluble organic redox
couples (we have termed ORBAT). This type of metal-free flow bat-
tery opens up a new area of research for realizing inexpensive and ro-
bust electrochemical systems for large-scale energy storage. The cells
were successfully operated with 1,2-benzoquinone disulfonic acid at
the cathode and anthraquinone-2-sulfonic acid or anthraquinone-2,6-
disulfonic acid at the anode. The cell used a membrane-electrode
assembly configuration similar to that used in the direct methanol
fuel cell.39
We have shown that no precious metal catalyst is needed
because these redox couples undergo fast proton-coupled electron
transfer.
We have determined the critical electrochemical parameters and
various other factors governing the performance of the cells. The stan-
dard reduction potentials calculated using density functional theory
were consistent with the experimentally determined values. This type
of agreement suggested that quantum mechanical methods for predic-
tion of the reduction potentials could be used reliably for screening
various redox compounds. The experimental values of the diffusion
coefficients of the various quinones in aqueous sulfuric acid sug-
gested that strong interaction of the ionized quinones with water
resulted in lower diffusion coefficients compared to those in non-
aqueous media. Further, we found that significant stabilization by
intra-molecular hydrogen bonding occurred with the sulfonic acid
substituted molecules. These differences will be important to con-
sider in interpreting the changes in the rate of proton-coupled electron
transfer in these molecules.
Our experiments also demonstrated that the organic redox flow
cells could be charged and discharged multiple times at high faradaic
efficiency without any sign of degradation. Our analysis of cell per-
formance shows that the mass transport of reactants and products and
their solubilities are critical to achieving high current densities.
Further testing and analysis is underway to understand the behavior
of other redox couples in the quinone family, assess the impact of
solubility on full cell performance, and optimize the structure of the
membrane-electrode assemblies. Solubility is still a challenge for this
type of redox flow battery. Choosing a substituent such as sulfonic
acid to modify both positive and negative electrode materials appears
to be the most promising approach at this time to meet the challenge of
solubility in water. However, understanding the effect of substituent
group type and placement on the standard reduction potential and
kinetic reversibility are also important areas for further study.
Acknowledgment
We acknowledge the financial support for this research from
ARPA-E Open-FOA program (DE-AR0000337), the University of
Southern California, and the Loker Hydrocarbon Research Insti-
tute. Both authors Bo Yang and Lena Hoober-Burkhardt contributed
equally to this work.
Appendix
Analysis of Current-voltage Curves at Various States of Charge
The following analysis is aimed at deriving a relationship between cell voltage and
discharge current for a flow battery using electrochemically reversible redox couples. For
any cell undergoing disharge, the cell voltage and current are related by Eq. A1.
Vcell = Ec − Ea − Idisch Rohm [A1]
Where Vcell is the cell voltage, Ec is the electrode potential of the cathode, Ea is the
electrode potential of the anode, I is the discharge current, and Rohm is the series equivalent
resistance of the electrolyte and current collectors.
For electrodes with reactions involving with rapid charge transfer kinetics relative to
mass transport processes (usually termed as kinetically reversible reactions), the electrode
potentials are given by the Nernst equation.
Ec = Eo
c + (RT/nF) ln (Co,c/Cr,c) [A2]
Ea = Eo
a + (RT/nF)ln(Co,a/Cr,a) [A3]
In Eqs. A2 and A3, and in subsequent equations, the subscripts c and a refer to the cathode
and anode, respectively. Co and Cr are the concentrations of the oxidized and reduced
species at the surface of the electrode.
Let Co,c*, Cr,c*, Co,a*, Cr,a* be the concentration of the species in the bulk of the
solution at the respective electrodes, and the asterisk will always refer to the values in
the bulk of the solution.When the flow rate is high relative to the consumption rate at the
electrodes, we expect the concentrations to be uniform in the plane of the electrode. Under
these conditions, we may describe the variation in concentration due to mass transport
limitations only in one-dimension, i.e., perpendicular to the surface of the electrode.
In the steady-state, the discharge current Idisch, can be related to the diffusion coeffi-
cient, D, concentration of reactants and products, and the diffusion layer thickness, δ by

Fick’s first law.
Idisch = nF AD(C∗
o,c − Co,c)/δ = nF Amt (C∗
o,c − Co,c) [A4]
Idisch = nF AD(Cr,c − C∗
r,c)/δ = nF Amt (Cr,c − C∗)
r,c [A5]
We also define the mass transfer coefficient, mt, as is equal to D/δ for the oxidized and
reduced species. The values of mt for the cathode and anode will be assumed to be same
in the analysis because we are considering a case of symmetrical electrode structures and
same flow rates, and molecules with very similar diffusion coefficients. A is the area of
the electrode.
Re-arranging Eqs. A4 and A5 we obtain,
Co,c = C∗
o,c − Idisch /nF Amt ; Cr,c = Idisch /nF Amt + C∗
r,c [A6]
Similarly, at the anode, the discharge current and surface concentration are given by,
Co,c = Idisch /nF Amt + C∗
o,c; Cr,a = C∗
r,c − Idisch /nF Amt [A7]
The change in concentration resulting from changes at the anode and cathode is given by
dCo,c = −dIdisch /nF Amt [A8a]
dCr,c = dIdisch /nF Amt [A8b]
dCo,c = −dIdisch /nF Amt [A9a]
dCr,a = −dIdisch /nF Amt [A9b]
From the Nernst equation, we can obtain the change in potential resulting from such
changes in concentration.
dEc = (RT/nF){(1/Co,c)dCo,c − (1/Cr,c)dCr,c} [A10a]
dEa = (RT/nF){(1/Co,a)dCo,a − (1/Cr,a)dCr,a } [A10b]
From Eqs. A8, A9 and A10 we obtain,
d(Ec − Ea) = −(RT/n2
F2
Amt )(1/Co,c + 1/Cr,c + 1/Co,a + 1/Cr,a)dIdisch [A11]
Let Cinitial be the concentration in both reservoirs of equal volume at the start of the
experiment. Then, even at various states of charge, the sum of the bulk concentrations of
the oxidized and reduced species on both sides will equal Cinitial.
Cinitial = Co∗
+ Cr∗
[A12]
Since Cinitial can now be used to define the state-of-charge, Q, as follows.
C∗
o,c = QCinitial [A13]
C∗
r,c = (1 − Q)Cinitial [A14]
For the anode,
C∗
r,a = QCinitial [A15]
C∗
o,a = (1 − Q)Cinitial [A16]
Combining Eqs. A6, A7, A11, and A12–A15, we obtain
d(Ec − Ea)/dIdisch = −2(RT/nF){(1/(nF Amt QCinitial − Idisch )
+1/(nF Amt (1 − Q)Cinitial + Idisch )
At any state of charge, the electrode potential is related to the state of charge by:
Ec,eq = Ecnot + RT/nF ln Q/(1 − Q) [A17]
Ea,eq = Eanot + RT/nF ln(1 − Q)/Q [A18]
The cell voltage in Eq. A1 can then be expressed as,
Vcell = Ec,eq − Ea,eq − Idisch {d(Ec − Ea)/dIdisch } − IRohmcell [A19]
Substituting for the various terms in Eq. A19 from above, we have an expression for
the cell voltage as a function of discharge current and state of charge, when the initial
concentration of the active materials are the same on both sides.
Vcell = Ecnot − Eanot + (RT/nF) ln Q2
/(1 − Q)2
− 2Idisch (RT/nF)[(1/(nFmt AQCinitial − Idisch )
+ 1/(nFmt A(1 − Q)Cinitial + Idisch )] − Idisch Rohmcell [A20]
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