pubs.acs.org/CR

Review

Electrolyte Lifetime in Aqueous Organic Redox Flow Batteries: A
Critical Review
David G. Kwabi, Yunlong Ji, and Michael J. Aziz*
Cite This: https://dx.doi.org/10.1021/acs.chemrev.9b00599

Downloaded via HARVARD UNIV on February 13, 2020 at 19:54:18 (UTC).
See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

ACCESS

Read Online

Metrics & More

Article Recommendations

ABSTRACT: Aqueous organic redox ﬂow batteries (RFBs) could enable widespreadintegration of renewable energy, but only if costs are suﬃciently low. Because the levelized
cost of storage for an RFB is a function of electrolyte lifetime, understanding and improving
the chemical stability of active reactants in RFBs is a critical research challenge. We review
known or hypothesized molecular decomposition mechanisms for all ﬁve classes of aqueous
redox-active organics and organometallics for which cycling lifetime results have beenreported: quinones, viologens, aza-aromatics, iron coordination complexes, and nitroxide
radicals. We collect, analyze, and compare capacity fade rates from all aqueous organic
electrolytes that have been utilized in the capacity-limiting side of ﬂow or hybrid ﬂow/
nonﬂow cells, noting also their redox potentials and demonstrated concentrations oftransferrable electrons. We categorize capacity fade rates as being “high” (>1%/day),“moderate” (0.1−1%/day), “low” (0.02−0.1%/day), and “extremely low” (≤0.02%/day)and discuss the degree to which the fade rates have been linked to decomposition mechanisms. Capacity fade is observed to be timedenominated rather than cycle-denominated, with a temporal rate that can depend on molecular concentrations and electrolyte stateof charge through, e.g., bimolecular decomposition mechanisms. We then review measurement methods for capacity fade rate andﬁnd that simple galvanostatic charge−discharge cycling is inadequate for assessing capacity fade when fade rates are low or extremely
low and recommend reﬁning methods to include potential holds for accurately assessing molecular lifetimes under suchcircumstances. We consider separately symmetric cell cycling results, the interpretation of which is simpliﬁed by the absence of adiﬀerent counter-electrolyte. We point out the chemistries with low or extremely low established fade rates that also exhibit open
circuit potentials of 1.0 V or higher and transferrable electron concentrations of 1.0 M or higher, which are promising performance
characteristics for RFB commercialization. We point out important directions for future research.

CONTENTS

References

1. Introduction
1.1. Flow Batteries and Storage of Renewable
Energy1.2. Levelized Cost and Calendar Lifetime
2. Chemical Decomposition and Capacity Fade2.1. Active Reactant Decomposition: Expectations from Theory and Chemical Studies
2.2. Electrochemical Assessments of Full Cell
Capacity Fade
2.3. Symmetric Cell Cycling: A Technique for
Coupling Capacity Fade to Chemical Decomposition
3. Summary
4. Conclusions
Author Information
Corresponding Author
Authors
Notes
Biographies
Acknowledgments© XXXX American Chemical Society

SA

1. INTRODUCTION

A
B
C

1.1. Flow Batteries and Storage of Renewable Energy

Increasing global energy needs, coupled with the harmful eﬀects
of CO2 emissions on the climate, have spurred interest in thedevelopment of carbon-free energy sources.1−3 The inherently
intermittent availability of solar and wind power, however,makes their deep penetration into the energy economy, andparticularly into the electricity grid, diﬃcult. This problem canbe solved by eﬃcient, inexpensive and scalable electrical energy
storage (EES) options.4,5 Whereas several storage options suchas capacitors, ﬂywheels, and enclosed batteries exist for shortduration grid services such as grid power quality and frequency

C
F

M
Q
Q
R
R
R
R
R
S

Special Issue: Beyond Li Battery Chemistry
Received: September 22, 2019

Ahttps://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

Figure 1. (a) Schematic of an RFB during battery charging. (b) Cyclic voltammogram curves showing redox reactions involving 2,6dihydroxyanthraquinone in the negolyte and ferricyanide in the posolyte, with vertical dotted lines indicating associated reduction potentials. The
dotted baseline represents the background cyclic voltammogram of a graphite foil electrode in 1 M KOH. Reprinted with permission from ref 11.
Copyright 2015 The American Association for the Advancement of Science.

regulation, the main option for storage with the long-duration
discharge needed to regulate wind or solar sources is pumped
hydroelectric energy storage. Pumped hydroelectric storage,
however, has a large areal footprint and is geographically limited
by requiring two large reservoirs at signiﬁcantly diﬀerent
elevations. Compressed air energy storage in undergroundcaverns requires natural gas combustion during discharge;
compressed air implemented adiabatically aboveground would
not require combustion but is in its infancy.
Given its lack of geographic constraints, EES in the form of
secondary (i.e., rechargeable) batteries can potentially enable
the storage of intermittent renewable energy at scale. Moreover,the large variety of chemistries and electrode architectures inwhich it can be embodied translates to a wide range of speciﬁc
energy and power capabilities in secondary batteries. Despite
this high operational ﬂexibility, the only electrochemical EES
system that has attracted extensive commercial attention is therechargeable lithium-ion battery (LIB), which has facilitated theadoption of portable electronic devices in daily life. Considerable research and industrial eﬀort has also been directed
toward the incorporation of LIBs into electric vehicles.
However, the prospects for their integration into long-duration
grid scale storage are challenged by projected issues withscarcity, geopolitical concentration of their raw materials,6 andﬁre hazards7,8 from the organic solvents they contain.
Aqueous redox ﬂow batteries (RFBs) are a potentially safe andcost-eﬀective technology for grid scale storage. In an RFB,
chemical energy is stored in the form of two redox-active speciesof disparate reduction potentials, which are dissolved orsuspended in an electrolyte; the electrolyte containing thespecies with the lower reduction potential is referred to as thenegative electrolyte (“negolyte”) whereas the other electrolyte is
the positive electrolyte (“posolyte”). Interconversion between
stored chemical and electrical energy during battery charging ordischarging is facilitated by pumping the electrolytes through anelectrochemical cell (Figure 1a) where electron transfer
reactions take place at a pair of porous electrodes, typicallycarbon-based, separated by a separator that is often an ionselective membrane. With this design, energy storage capacity
scales with the volume of the electrolyte reservoirs, concentrations of redox-active species therein, and the diﬀerence in
reduction potentials of the energy-storing species (Figure 1b);whereas the power scales primarily with the active area in theelectrochemical cell stack. This enables energy and power to bescaled independently, unlike in an enclosed battery. Con-

sequently, at large energy-to-power ratios (i.e., long discharge
durations at rated power), the system cost per kWh of stored
energy is largely dependent on the cost of the electrolytes, rather
than the cost of the electrochemical cell.9 Although the energy
density (energy stored per unit volume) and speciﬁc energy(energy stored per unit mass) of aqueous electrolytes are anorder of magnitude lower than for LIBs, these metrics are ofsecondary importance for stationary storage. Techno-economic
analysis estimates that an aqueous RFB with an open-circuit
voltage of 1.5 V, total electrolyte cost of $5 per kg of active
material and 150 g per mole of transferrable electrons should
cost about $120/kWh at the system level, with room for further
cost reductions with cheaper electrolytes and large enoughproduction volumes.10
1.2. Levelized Cost and Calendar Lifetime

Minimizing electrolyte cost is critically important to enablingcommercialization of RFBs. Chemistries based on vanadium inhighly acidic electrolytes have attracted the most research andindustrial attention,12−14 featuring the following negolyte andposolyte redox reactions and associated reduction potentials(E0) versus the standard hydrogen electrode (SHE):
V3 + + e− ↔ V2 +(E 0 = −0.255V)

VO2+ + 2H+ + e− ↔ VO2 + + H 2O

(E 0 = 0.991V)

respectively. The diﬀerence between these reduction potentials
yields a full-cell voltage of 1.25 V for a proton concentration of 1
M, which increases with decreasing pH because only the second
reaction is proton-coupled and thus has a higher potential atlower pH. Considerable strides have been made toward thedevelopment and commercialization of all-vanadium RFBs, withthe largest installation, rated at 200 MW/800 MWh, currently
under construction by Rongke Power in Dalian, China.
However, although redox-active vanadium species are stableindeﬁnitely, the price of vanadium is too volatile and too high on
average for globally ubiquitous deployment in large-scale RFB
installations. To wit, the United States’ Department of Energy
estimates that energy storage systems for utility-scale delivery ofelectricity should cost not more than $150/kWh in order to becommercially feasible,15 and vanadium RFBs are expected tohave higher system costs.16
As a result, inorganic RFB chemistries based on more Earthabundant raw materials, are being developed, includinghydrogen−bromine,17,18 metal-hydride,19 sulfur−oxygen,9 andB

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

slurry-based all-iron20−22 conﬁgurations. Despite the generally
low projected chemical costs of these systems, all of them facecritical technical challenges that currently prevent them from
replacing vanadium RFBs. For instance, the all-iron and sulfur−
oxygen chemistries have projected chemical costs of ∼6.59 and 1$/kWh,9 respectively. However, for the all-iron slurry system,
spatially inhomogeneous plating of elemental iron in thenegolyte may result in low capacity utilization, and the highviscosity of the slurry exacerbates mass transport and pumping
losses; the sulfur−oxygen chemistry faces low power densities(<10 mW/cm2) due to high membrane resistances and thesluggish kinetics of the oxygen reduction and evolution
reactions. Hybrid zinc-based ﬂow/nonﬂow conﬁgurations suchas zinc-halide (zinc-bromide23 or zinc-iodide24,25), zinc−iron,26
and zinc−cerium27 systems have also received researchattention, but lose the advantage of decoupling of energy andpower because zinc electroplating capacity scales with cell area,rather than electrolyte reservoir volume.
An alternative pathway toward low-cost RFBs is the use oforganic and organometallic redox-active reactants. Although
they are expected to yield lower full-cell energy densities than
the all-vanadium system,28,29 the high abundance of their raw
material precursors potentially enables much lower electrolyte
costs than vanadium, and their thermochemical properties−
chieﬂy solubility and reduction potential−can be tailored via
chemical structure modiﬁcation and functionalization within awide parameter space. Their molecular size can also be tailoredto suppress the rate of undesirable permeation throughmembranes and separators without debilitating increases inviscosity.30,31 It is worth noting that aqueous RFB chemistries
are generally limited by a full-cell voltage of roughly 1.5−1.8 V,
given the practical potential limits for the electrolysis of water.
This has the eﬀect of limiting their expected energy densities.
Nonaqueous RFBs may reach substantially higher voltages andenergy densities; however, these are based on low-conductivity,
ﬂammable, and expensive solvents.
A key deﬁciency of most organic chemistries that have beenreported is capacity fade rates exceeding 0.1%/day,32−35 chieﬂycaused by the decomposition of their organic constituents,
although permeation of active material through the membranealso plays a nontrivial role in many studies. Such high rates ofcapacity fade are utterly impractical for RFB installations withproject lifetimes anticipated to be decades long. Decadalchemical lifetime is not, however, strictly necessary forcommercially interesting aqueous organic RFB conﬁgurations.
Rather, a more appropriate goal is to achieve suﬃcient stability
at a reasonable levelized cost of the storage chemistry, given thepossibility of periodically replacing the lost electrolyte. Levelized
cost in this case refers to the present value of the initial andreplacement costs amortized over the energy delivered out ofstorage over the lifetime of the project. In particular, if thepresent value of a series of future replacement costs for an
organic RFB system is lower than the capital cost savings fromusing an organic system rather than a vanadium system, then thelower-cost organic system may be the more economically
preferable option (Figure 2).36
This trade-oﬀ is expressed quantitatively by the replacement
cost ratio, deﬁned as
replacement cost ratio ≡Review

Figure 2. Breakeven value of replacement cost ratio vs interest rate fordiscounting, assuming project lifetimes as indicated. Adapted withpermission from ref 36. Copyright 2017 John Wiley and Sons.

Because the annual replacement cost is inversely proportional tothe calendar lifetime and thereby plays such a major role in thistechno-economic reasoning, comprehensively understanding
and mitigating temporal capacity fade in aqueous organic RFB
research is a compelling and urgent research task. This review isdedicated to the status of our understanding in this respect. We
particularly emphasize the importance of judicious choice ofexperimental method in rationalizing capacity fade in terms ofspeciﬁc reactant decay or loss mechanisms, as well as the
temporal capacity fade rate as a more useful indicator of RFB
chemical lifetime than the more often reported cycledenominated fade rate. In addition to reactant decay, reactant
permeation through the ion-selective membrane may also cause
capacity loss,37 and where appropriate, we discuss theimportance of distinguishing between the two pathways forparticular chemistries. We note, however, that decomposition inaqueous organic ﬂow cells is typically the more dominantcapacity fade mechanism, and that crossover should become
considerable for chemistries exhibiting extremely low fade rates
or inorganic redox-active species.38−40 Other aspects of aqueousorganic RFB performance, as well as issues related to aqueous
metal-ion-based and nonaqueous RFB reactant chemistries, arebeyond the scope of the review; interested readers are directed
to surveys of studies in those ﬁelds encompassing work
published in the past decade.16,29,41−47

2. CHEMICAL DECOMPOSITION AND CAPACITY FADE
2.1. Active Reactant Decomposition: Expectations from
Theory and Chemical Studies

The chemical decomposition of organic molecules is a wide andexpansive topic, however, we restrict the present discussion to
the set of redox-active reactants currently under investigation inaqueous organic RFB research, which we group here into thefollowing classes: (i) quinones, (ii) organometallic coordination
complexes, (iii) nitroxide radical derivatives, (iv) viologens, and(v) all other aromatic heterocycles. Thus, in this review, we
consider only those mechanisms to which species from theseclasses have been hypothesized or shown to be susceptible,
which would lead to loss or diminution of redox activity, such as(Scheme 1): (i) nucleophilic addition or substitution such as
gem-diol formation, Michael addition and hydrolysis, (ii)disproportionation to redox-inactive centers, (iii) dimerizationor polymerization, and (iv) tautomerization.

annual replacement cost
upfront capital cost savingsChttps://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

Scheme 1. Illustration of the Main Chemical Decomposition Mechanisms for Redox-Active Molecules Considered in This Review

hydrophilicity of the functional groups attached to it.52 For
example, anthraquinone sulfonic acid (AQS) exhibits a ∼20 mVlower reduction potential than anthraquinone disulfonic acid
(AQDS), due to the removal of one of the electron-withdrawing
sulfonic acid groups. Dihydroxyanthraquinone disulfonic acid
(DHAQDS) in turn has an ∼80 mV lower reduction potential
than AQS, owing to the installation of two electron-donatinghydroxyl groups next to one of the ketone groups. Distancing theelectron-withdrawing sulfonic acid groups from the redox-active
center, as in dihydroxyanthraquinone dimethylsulfonic acid
(DHAQDMS), lowers the reduction potential of the quinone byanother ∼100 mV.
The inﬂuence of substituent hydrophilicity and location on
solubility follows trends suggested by theory but is notquantitatively understood. In agreement with the study by Er
et al.,48 4,4′-((9,10-anthraquinone-2,6-diyl)dioxy)dibutyrate(2,6-DBEAQ) has been shown to have a signiﬁcantly higher
solubility53 than 2,6-dihydroxyanthraquinone (2,6-DHAQ)11
under alkaline conditions (>1 M vs 0.6 M at pH 14; 0.6 M vs 0.1
M at pH 12), ostensibly due to the carboxylate solubilizing
groups of the former being located farther away from theanthraquinone core. Replacing the carboxylate functionalgroups with more hydrophilic phosphonate groups in 2,6di(3-phosphonic acid)propyl ether anthraquinone (2,6DPPEAQ)54 boosts the solubility even further (0.75 M vs 35
mM at pH 9). However, depending on the molecule class, otherstrategies for modifying reactant solubility based on otherchemical principles are possible. For instance, based oncomputational work suggesting that asymmetric charge
distribution improves organic molecule solubility,55 Hollas etal.34 synthesized dihydroxyphenazine sulfonate (DHPS), which
has a considerably higher solubility than the parent phenazine
molecule under strongly alkaline conditions (1.80 M vs nearzero at pH 14). Luo et al. have also shown that replacing the

Because the number of possible combinations of RFB reactant
solubility, reduction potential and decomposition mechanisms
is practically inﬁnite, rapid computational screening of large
databases of particular redox-active molecule types and their
thermochemical properties has attracted attention recently.
With regard to quinones, Er et al. introduced a high-throughput
computational screening protocol48 to generate a virtual libraryof 10 000 molecules. Using density functional theory (DFT),
they examined the inﬂuence of selected functional groups andtheir spatial positioning on the quinones’ reduction potentials
and solvation energies and sought to identify structure−property trends that could inform how solubility and reduction
potential may be rationally tuned. Analysis of the results yielded
two important conclusions: (i) at least 31 and 28 synthesizablequinones had reduction potentials below 0.2 V and above 0.9 V
vs SHE, respectively, and (ii) placement of long alkyl chains withterminal hydrophilic functional groups increased quinone
solubility without aﬀecting reduction potential, whereas
attaching electron-withdrawing or-donating groups close to
the ketone group had a large impact on the quinone’s reduction
potential. This latter point was buttressed by recent combined
experimental-DFT studies by Huynh et al.49 and Kim et al.50,51
In the report by Huynh, it was found that a quinone’s reduction
potential scales linearly with its attached substituents’ eﬀective
Hammett constant, which is an estimate of the substituents’electron-withdrawing ability. The study, however, notedsigniﬁcant deviations from the scaling relationship for quinones
featuring intramolecular hydrogen bonding, as well as
halogenated, charged, or sterically bulky substituents.
There has been some experimental support of these proposed
relationships among functional group chemistry and positioning, and reactant reduction potential. Gerhardt et al. showed that
it is possible to tune the reduction potential of an anthraquinone
derivative through manipulating the number, position andD

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Reviewgroup leads to the loss of trimethylamine as a leaving group.67,68
As water is a weaker nucleophile than OH−, an important lesson
from this work is that aqueous RFB reactants based on nitroxide
radicals should ideally be deployed in neutral pH electrolytes,
and those based on Fc or ferricyanide should ideally be deployedin neutral-to-weakly alkaline electrolytes.
It is worth noting that some low-potential redox couples may
also be prone to decomposition by nucleophilic attack inaddition to other mechanisms that may delimit a range ofpreferred operating conditions. This is the case for methyl
viologen (MV), although it has low ﬁrst and second electron
reduction potentials at −0.45 and −0.7 V vs SHE, respectively.
The ammonium functional groups in either the fully oxidized(MV2+) or singly reduced (MV•+) redox states are vulnerable toOH− attack, whereas the doubly reduced MV° species has poor
water solubility and is prone to protonation under acidic
conditions.69 Even at pH 7, however, dimerization of two
monocationic MV+ radicals can result in the formation of astable internal charge transfer complex70 that can thendisproportionate into MV2+ and MV°,71,72 resulting in loss ofredox activity.
A subset of nucleophilic reactions initiated by water to which
aqueous organic RFB reactants are prone is the breakdown, orhydrolysis, of substituents attached to redox-active molecules.
Functional groups such as esters, ethers, and amides73 areparticularly vulnerable to this attack, and in some cases,
hydrolysis can lead to the collapse of the backbone of theoxidized form of the redox-active molecule, resulting in acomplete loss of redox activity. Such collapse can be catalyzed by
H+ or OH− in acid or alkaline electrolytes, respectively,
depending on the nature of the substituent. Hydroxide-initiatedhydrolysis of alloxazines such as lumichrome,74,75 and theirtautomer isoalloxazines such as riboﬂavin,76−78 has been widely
studied. Koziol et al. have shown that under alkaline conditions
the uracil ring in lumichrome can be opened through thehydrolysis of amide, leading to the formation of 6,7-dimethy1-3ureidoquinoxaline-2-carboxylic acid.74 Riboﬂavin likewise canundergo ring-opening upon hydrolysis, yielding urea and l,2dihydro-6,7-dimethyl-2-keto-l-D-ribityl-3-quinoxaline carboxylic acid.76−78 Although high pH values catalyze amidehydrolysis,77 the low pKa values at the amide nitrogen (8.4
and 11.4 for lumichrome) result in the accumulation of two
negative charges in the imidic conjugate system, which lowers
the electrophilicity of the carbonyl groups and hinders the attack
by OH−.74 Introducing electron-donating substituents toorganic molecules prone to hydrolysis by OH− is therefore
expected to lower their reactivity, and a computational study
suggests that attaching such groups to alloxazines can indeed
decrease the electrophilicity of the carbonyl groups and protect
against hydrolysis.79
The disproportionation of some redox couples can initiatechemical decomposition. Disproportionation is said to occur
when two species with identical oxidation states undergo a redox
reaction in which one is reduced and the other is oxidized,
resulting in two diﬀerent products, of lower and higher oxidation
states, respectively. Capacity fade in an RFB results if either ofthe products is redox-inactive or is of suﬃciently low solubility to
precipitate. Moreover, if either product exhibits a more extreme
redox potential than the original couple or a similar redox
potential with more sluggish kinetics, an apparent capacity fadewill be observed for given cycling potential limits. Quinones are
susceptible to this mechanism: the disproportionation ofanthrahydroquinone (2-electron reduced anthraquinone) tocation in potassium ferrocyanide (K4Fe(CN)6) or sodium
ferrocyanide with the more hydrophilic ammonium cation
roughly doubles the Fe(CN)64− solubility at pH 7 from 0.76 or
0.56 M, respectively, to 1.56 M;56 however, generally reliable
quantitative structure−property relationships for organic
molecule solubility do not exist.
Nucleophilic addition or substitution is one of the mostcommon modes of chemical decomposition for aqueous RFB
reactants, and a number of computational studies have examined
their thermodynamic susceptibility to such attack. It occurs
when a nucleophilic species (such as water, hydroxyl/hydroperoxide anions, or reduced redox-active species) reacts with anelectron-deﬁcient redox-active species, typically the oxidized
form of the organic redox couple. In this reaction the attacking
nucleophile may either irreversibly replace a good leaving
substituent attached to the molecule57 or attach to an
unsubstituted position,58 resulting in either loss of conjugation(and hence complete loss of redox activity) or a large decrease inreactant reduction potential and/or solubility.
Chemical principles suggest that nucleophilic attack is more
likely to occur as either the Lewis basicity of the nucleophile oracidity of the substrate, increases. To assess this, Tabor et al.59
conducted a computational study of a virtual library of 140 000
benzene, naphthalene, anthracene, phenanthrene, and tetracene-based quinone-hydroquinone couples, in which theyexamined the susceptibility of each quinone to a Michael
addition or substitution, as well as reversible gem-diol formation.
The main conclusion was that as the quinone reduction
potentials (and thus electron deﬁciency or Lewis acidity)
increased above 0.9 V vs SHE, the quinones became severely
susceptible to irreversible Michael addition/substitution reactions with water. Indeed, they found that almost no quinones
exhibited both a reduction potential above 0.9 V, andthermodynamic resistance to Michael addition and/or gemdiol formation. As will be discussed in more detail in the next
section, this insight has been strongly supported by theexperimental literature, which features signiﬁcantly morechemically stable quinone-based negolyte reactants (withreduction potentials <0.2 V vs SHE)53,54,60 than posolyte
reactants (with reduction potentials >0.8 V), the most stable ofwhich decomposes at ∼0.5%/day.61
Chemical studies have established that nucleophilic reactions
occur in other nonquinone-related high-potential substrates as
well, such as nitroxide radicals,62,63 which have reduction
potentials around 0.8 V vs SHE, and organometallic complexes
based on ferrocene (Fc) or ferrocyanide, which have reduction
potentials around ∼0.4−0.6 V. As an example of the former case,electrochemical oxidation of 4-hydroxy-(2,2,6,6 tetramethylpiperidin-1-yl)oxyl (4-HO-TEMPO) results in the formation of 4HO-TEMPO + , which, in alkaline electrolyte, is highly
susceptible to nucleophilic addition of OH−, which causes
progressive acidiﬁcation of the electrolyte. Acidic conditions
lead to 4-HO-TEMPO disproportionation to 4-HO-TEMPO+
and 1,4-dihydroxy-2,2,6,6-tetramethylpiperidine, the latter ofwhich is redox-inactive. Similarly, electrochemical oxidation of
Fc and ferrocyanide results in the formation of Fc+ andferricyanide, respectively, both of which are prone tonucleophilic attack by OH−.64,65 In the case of Fc+, nucleophilic
attack by water may still occur to some extent, even when
protective alkyl substituents are used,66 particularly if thosegroups are close to the Fc redox center. An example of this isfurnished by (ferrocenylmethyl)-trimethylammonium salts,
where nucleophilic substitution of the quarternary ammonium
E

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

concentration of redox-active reactants, and repeated reductionoxidation cycling is a means to assessing average RFB
chemical lifetime. The fraction of negolyte (posolyte) active
species in the reduced (oxidized) state is often referred to as thestate of the charge (SOC) of the electrolyte.
In most studies, ﬂow cells comprising two electrolytes
separated by an ion-exchange membrane are used for capacity
fade analysis. The reactant of interest is typically dissolved in thecapacity-limiting side (CLS) or electrolyte, whereas thenoncapacity-limiting side (NCLS) contains a counter electrolyte with a higher initial Coulombic capacity. It is important that
the NCLS have excess capacity for discharging as well as forcharging, so that small amounts of side reactions such ashydrogen evolution or oxygen evolution, as well as degradationof the active species itself, do not cause the NCLS to reach its
capacity limit. Where the NCLS comprises a diﬀerent redox
couple than that in the CLS, the setup (hereafter referred to as
the “full cell” setup) mimics the basic unit of an actual RFB(Figure 1a). The current eﬃciency of a full cell is often evaluatedas the ratio of the discharge duration to the immediately
preceding charging duration in a galvanostatic charge−discharge
cycle. Any nonrecoverable capacity losses originate fromdecomposition of redox-active material in the CLS, provided
electrolyte leakage and permeation across the membrane, thelatter of which increases with concentration,37 do not occur toany appreciable extent.88 Limiting unwanted current eﬃciency
and capacity losses via ingress of atmosphere, electrolyte leakage,or electrolyte permeation through the membrane by using
properly sealed plumbing and/or testing the ﬂow cell in a
glovebox, as well as judicious choice of membrane37 aretherefore critical to proper evaluation of the stability of RFB
chemistries.
The inﬂuence of cycling protocol on the measured capacity
fade rate, and how it may diﬀer from the actual fade in
Coulombic capacity of the electrolyte, is an underappreciatedissue. Constant-current (galvanostatic) cycling, where eachreduction/oxidation half-cycle ends when a prescribed cell
overpotential is reached, is performed in most studies. However,
because of high mass transport overpotentials at extreme SOC
values, as well as the time lag between the attainment of aparticular SOC value in the cell and its attainment in thereservoirs, less than 100% of the actual capacity is accessed
before the switching occurs,33,89 with this apparent capacity
dropping with increasing current density. Consequently,
transient changes in membrane resistance or surface exchange
kinetics, caused by temperature variations, membrane aging, orother time-varying environmental factors, can result incorresponding changes to apparent capacity, although the actual
electrolyte capacity may remain unchanged and accessible via
other cycling regimens. These spurious changes in apparent
capacity are typically a negligible fraction of the actual
electrolyte capacity but become more important for chemicallystable electrolytes where capacity fade from moleculardecomposition is comparably small. A more rigorous techniquefor interrogating actual capacity is constant-potential, or
potentiostatic, cycling, where each half-cycle ends only after
the reduction/oxidation current has become negligible, at which
point close to 100% of the available charge is expected to havebeen extracted. This is strictly true if the chosen potential limits
provide large enough overpotentials to access essentially 100%of the electrolyte capacity and side reactions at those limits are
negligible. Equally eﬀective is galvanostatic cycling in which eachhalf-cycle is followed by a potential hold until the current has

anthrone and anthraquinone was observed during electrochemical reduction of unsubstituted anthraquinone underhighly acidic condition.80 Shi et al. also obtained monomers
and dimers of 1,8-dimethoxy-10-anthrone upon chemical
reduction of 1,8-dimethoxyanthraquinone with sodium borohydride and triﬂuoroacetic acid.81 In agreement with these studies,
Goulet and Tong et al.32 and Jin and Jing et al.82 havedemonstrated the formation of anthrone and its dimer inalkaline and neutral media, respectively, their results are
discussed in more detail in the next section.
Tautomerization occurs where there is a formal migration of ahydrogen atom along with an exchange between a single bondand an adjacent double bond without changing the carbon
skeleton.83 Tautomers generally have diﬀerent thermochemical
properties than the original molecule, and if they have
suﬃciently more extreme reduction potentials or lower
solubilities as well, the formation of any such species during
electrochemical cycling will result in capacity fade and/or areduction in cell energy density. Aza-substituted conjugated πsystems based on quinones, viologens, and phenazines areusually susceptible to this transformation during electrochemical
cycling.84 This is because redox reactions of these organic
molecules involve the reorganization of conjugated double
bonds, and in the course of such rearrangement, a tautomer may
be a metastable intermediate. This is not the case for simpler
aromatic molecules. Phenols, for example, exist exclusively in theenolic form, as the energy decrease by rearrangement to thetautomeric keto-structure is oﬀset by the simultaneous large
decrease in resonance energy.85 Tautomeric ketonic structures
become particularly energetically favorable in naphthoquinones,86 polyhydric phenols and especially in the hydroxyl
derivatives of polycyclic aromatic hydrocarbons.
Understanding the susceptibilities of various classes ofaqueous organic species to decomposition mechanisms
discussed above is critical to devising molecular engineering
strategies for enhancing RFB reactant stability. This is also
important because decay rates may increase with reactant
concentration for certain chemistries, such as those subject tobimolecular decomposition pathways (e.g., disproportionationand dimerization). Because practical RFBs would ideally operate
at high reactant concentrations in order to achieve high energy
densities, any trade-oﬀ between energy density and stability
must be well understood. In general, quantitative correlation of
capacity fade measured at the cell level to particulardecomposition mechanisms is an important research task and,
where it has occurred, is highlighted in this review.
2.2. Electrochemical Assessments of Full Cell Capacity Fade

Electrochemical measurements involving redox reactions at anelectrode are a natural way to probe the chemical stability of aredox couple in both its oxidized and reduced forms. A classic
means of doing this is via cyclic voltammetry (CV), in which thereversibility of the current passed at the electrode during cyclic
potential sweeps is a direct indication of the chemical stability ofthe oxidized and reduced forms of the redox couple involved.87
Given the sensitivity of CV measurements, the longest timescales for chemical decay for which CV-based analysis of bulk
concentrations can be used are, however, much shorter (on theorder of minutes) than typical time scales for aqueous RFB
reactant stability (hours to days). Thus, bulk electrolysis in ﬂow
cells is by far the most common technique used to assess
chemical stabilityhere, Coulombic capacity upon oxidation
and reduction of the entire electrolyte is a proxy forF

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

Table 1. Summary of Results from Full Cells Containing Quinone-Based Reactants in the CLSa
capacity fade rate
name
DMBQ/ K4Fe(CN)6DMOBQ/K4Fe(CN)6
DHBQ/K4Fe(CN)6
DHAQ/K4Fe(CN)6
bislawsone/K4Fe(CN)6
2,6-DBEAQ/K4Fe(CN)6
2,3-HCNQ/K4Fe(CN)6
2,6-DPPEAQ/K4Fe(CN)6
PEGAQ/K4Fe(CN)6
ARS/BQDS
AQS/HBr
AQDS/HBr
Zn(OH)42−/FQH2
AQDS/DHDMBS

author/year
35

Sun/2019

Yang and Tong/201836
Lin/201511
Goulet and Tong/201932
Tong/201998
Kwabi and Lin/201853
Wang/2018100
Ji/201854
Jin and Jing/201982
Zhang/2015102
Gerhardt/201752
Huskinson/201460
Chen/201688
Park/2019103
Hoober-Burkhardt/201796%/cycle%/day

electron conc. (M)redox potential (V vs SHE)

pH

0.11
0.17
0.24
0.1

22
35
9
8
0.14
0.74
0.04
3.4
0.014
0.5
16
19
0.19
1.2
2.5
3.5

0.2
0.2
1
1
0.18
2
1
0.5
1
3
0.1
2
2
2
0.2
2

−0.75/0.50−0.70/0.50−0.72/0.50−0.68/0.50−0.68/0.50−0.55/0.50−0.54/0.50−0.53/0.50−0.47/0.50−0.43/0.50
0.08/0.91
0.19/1.06
0.21/1.06
0.21/1.06−1.30/0.70
0.21/0.82

14
14
14
14
14
14
12
14
9−13
7
0
0
0
0
-0.7
0

0.038
0.001
0.053
0.0004
0.04
2
1
0.014
0.1
0.09
1

a
Capacity fade results established by potentiostatic cycling or galvanostatic cycling with potentiostatic ﬁnishes, are marked in bold text, whereas
apparent capacity fade results from purely galvanostatic cycling are left in plain text. In the “name” ﬁeld, the active reactant in the negolyte is on theleft and that for the posolyte is on the right. The CLS is marked out in bold, whereas the NCLS is in plain text. “electron conc.” refers to moles oftransferrable electrons per liter of electrolyte. “redox potential” refers to redox couple in the CLS (bold), and that of the NCLS (plain text).
DMOBQ = 2,5-dimethoxy-3,6-dihydroxy-1,4-benzoquinone, DMBQ = 2,5-dimethyl-3,6-dihydroxy-1,4-benzoquinone, DHBQ = dihydroxybenzoquinone, DHAQ = 2,6-dihydroxanthraquinone, 2,6-DBEAQ = 4,4′-((9,10-anthraquinone-2,6-diyl)dioxy)dibutyrate, 2,3-HCNQ = 2-hydroxy-3carboxy-1,4-naphthoquinone, 2,6-DPPEAQ = (((9,10-dioxo-9,10-dihydroanthracene-2,6-diyl) bis(oxy))bis(propane-3,1-diyl))bis(phosphonicacid), PEGAQ = 1,8-bis(2-(2-(2-hydroxyethoxy)ethoxy)ethoxy)anthracene-9,10-dione, ARS = 3,4-dihydroxy-9,10-anthraquinone-2-sulfonic acid,
AQS = anthraquinone monosulfonic acid, AQDS = anthraquinone disulfonic acid, FQH2 = 2,3,5,6-tetrakis((dimethylamino)methyl)hydroquinone,
DHDMBS = 3,6-dihydroxy-2,4-dimethylbenzenesulfonic acid.

depends only on the passage of time, although the rate candepend on molecular concentrations, pH, and SOC.32,33,54,82,90
It is also worth noting that galvanostatic cycling was used toassess capacity retention in all but ﬁve32,53,54,98,99 of thesestudies. However, given that the vast majority of reported faderates are above 0.1%/day (which extrapolates to 36.5%/year
assuming a constant fade rate),35,36,52,60,61,82,96,100 it is likely thatmuch of that apparent capacity fade is a consequence ofmolecular decomposition rather than of the artifacts to whichgalvanostatic cycling is vulnerable.
Understanding how capacity fade rates can be quantitatively
attributed to speciﬁc decomposition mechanisms can informmitigation strategies. A notable example of such an exercise is thework of Narayanan and co-workers regarding the developmentof a benzoquinone -based posolyte for aqueous organic
RFBs.61,95,96 They ﬁrst demonstrated the feasibility of operating
an aqueous organic RFB with 4,5-dihydroxybenzene-1,3disulfonic acid (BQDS), also known as tiron, in the posolyte
and either anthraquinone-2-sulfonic acid (AQS)101 or anthraquinone-2,6-disulfonic (AQDS)95 acid in the negolyte.
The hydrophilic, electron-withdrawing sulfonic acid groups
on BQDS were found to confer on it a high solubility (∼4 M)
and a relatively high reduction potential (0.85 V vs SHE),
yielding full-cell potentials of 0.65 and 0.80 V when paired
against AQDS and AQS, respectively. Galvanostatic cycling of acapacity-balanced AQDS/BQDS cell yielded a very highapparent capacity fade rate (50% capacity loss after the ﬁrst
cycle), after which the authors conducted chemical analysis ofcycled electrolytes using proton nuclear magnetic resonance(NMR) spectroscopy.95 Signiﬁcant molecular changes wereobserved only in the posolyte, which led to the hypothesis, inagreement with previous work,91 that the ortho-quinone
resulting from electrochemical oxidation of BQDS was subject

become negligible or in which capacity is assessed only inoccasionally inserted potentiostatic cycles.33,90 To demonstrate
the vulnerability of galvanostatic cycling results to misinterpretation, Goulet and Aziz33 inserted and removed a series
resistance equivalent to a 25% increase in cell resistance into a
cycling experiment, demonstrating artiﬁcial increases anddecreases in apparent capacity from galvanostatic cycling
measurements, and demonstrating the absence of such an
artifact in potentiostatic or combined potentiostatic-galvanostatic cycling. Henceforth, we use the term “apparent capacity”
to describe the result of galvanostatic cycling without potential
holds.
Table 1 summarizes temporal and cycle-based capacity faderates of some recently reported aqueous organic full cells where
the CLS contains a quinone-based redox couple. High-potential
redox couples based on benzoquinone were among the ﬁrst91−94
reactants studied for aqueous ﬂow batteries, and in recent years,
considerable eﬀort has been put into developing chemicallystable low-53,54 and high-potential61,95−97 quinone-based redox
couples. There is no correlation between cyclic and temporalfade rates in Table 1. This indicates that widely diﬀering cycle
periods (which are functions of the applied current, reactant
concentration and electrolyte volume) have been used in theliterature, and that full cell chemistries with low fade rates per
cycle will not necessarily have a long RFB calendar lifetime,
particularly in practical deployment, where cycle periods arelikely to be much longer than in published research. It is
important to note that, for all aqueous organic electrolytes forwhich we have observed suﬃcient experimentation to
distinguish between cycle-denominated fade and time-denominated fade, it has been found that the amount of lost capacity isessentially independent of the number of charge−discharge
cycles to which a species is subjected: the amount of lost capacityG

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

Figure 3. (a) Schematic of the transformation of BQDS during charging; (b) schematic of the Michael reaction showing the nucleophilic addition ofwater to the ortho-quinone resulting from BQDS oxidation. Reprinted with permission from ref 95. Copyright 2016 Electrochemical Society.

to chemical reduction via the Michael addition of water. Two
successive electrochemical oxidation and Michael addition
reactions eventually result in the formation of a fully substituted
para-benzoquinone with two electron-donating hydroxyl
groups,104 and a correspondingly lower redox potential (∼0.65
V vs SHE) than the original redox couple (Figure 3).
These transformations led to the negolyte becoming capacitylimiting, as three equivalents of AQDS were required tocompletely convert every equivalent of BQDS to its ﬁnal form.
Moreover, the electron-donating hydroxyl substituents introduced by Michael reactions decreased the reduction potential ofthe molecule, cutting the overall cell voltage to less than 0.5 V.
To mitigate the Michael reaction, Narayanan and co-workers
designed a new redox couple based on 3,6-dihydroxy-2,4dimethylbenzenesulfonic acid (DHDMBS).96 This moleculefeatures two moderately electron-donating methyl groups (asopposed to strongly electron-donating hydroxyl groups) and
one electron-withdrawing sulfonic acid group and had arelatively high redox potential (0.82 V vs SHE) and solubility(2 M in 1 M H2SO4). However, because the methyl groups
lower the electrophilicity of the only unsubstituted position,
which is itself adjacent to one of the bulky methyl groups,
DHDMBS was hypothesized to be less sterically and energetically prone to Michael reaction compared to BQDS uponelectrochemical oxidation. Indeed, in an AQDS/DHDMBS ﬂowcell, DHDMBS did not show any noticeable chemical
decomposition for 25 cycles, although there was signiﬁcantapparent capacity fade, at a rate of 5.5%/day. This was attributed
to a combination of DHDMBS crossover through the Naﬁon
membrane and its slow desulfonation by hydrolysis of thesulfonic acid, the latter of which causes capacity fade via removalof the solubilizing sulfonic acid group on DHDMBS, and thus
loss of redox-active material from solution through precipitation.
In a subsequent publication,61 crossover and proto-desulfonation were disambiguated using a combination of a lowpermeability sulfonated polyether ether ketone-based membrane and a symmetric cell protocol (discussed in more detail in
the next section), in which both electrolytes comprised

DHDMBS and AQDS. This protocol substantially reduced
DHDMBS crossover, and allowed direct interrogation of theinﬂuence of desulfonation on the apparent capacity fade rate,
which turned out to be about 0.48%/day in the absence ofsulfuric acid.
It is notable that DHDMBS, as a rationally designed posolyte
reactant, is the only quinone-based species with a redox potential
above 0.8 V vs SHE that exhibits moderate stability (between 0.1
and 1%/day) or better upon electrochemical cycling. That highpotential benzoquinone-based reactants have generally notexhibited high long-term chemical stability91,92,103,105,106
amounts to experimental validation of previously discussed
theoretical calculations59 suggesting that quinones become
more susceptible to Michael addition/substitution reactions as
their reduction potentials increase. Indeed, particularly direct
experimental evidence has been furnished for this conclusion ina recent study by Preger et al.,107 in which the chemical stabilityof several benzoquinone-derived molecules was evaluated, andthey found a strong inverse correlation between quinone half-lifeand reduction potential.
A consequence of this trend is that considerably more successhas come from the development of quinone-based negolytes,
which are often paired with high-potential inorganic redox
couples (chieﬂy based on bromine or ferrocyanide in acidic orbasic conditions, respectively) in full cells. The ﬁrst of suchnegolyte reactants was anthraquinone-2,7-disulfonic acid
(AQDS),60 which exhibited an apparent capacity fade rate of
0.2%/day in a full cell with a hydrobromic acid posolyte. There issome ambiguity regarding the mechanism of full redox access ofthe theoretical 2-electron Coulombic capacity of AQDS. On thebasis of nuclear magnetic resonance (NMR), CV, and bulk
electrolysis measurements, Carney108 and Wiberg109 and coworkers have suggested that reversible intermolecular dimerization of AQDS occurs at solution concentrations greater than 10
mM, which inhibits full 2-electron electrochemical reduction or
reduction using chemical oxidants. Carney et al.108 speculated
that CO2 is bound in a 1:1 anthraquinone-CO2 adduct in bulk
AQDS powder. Upon dissolution in acidic electrolyte, theH

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

Figure 4. Current eﬃciency (circles), energy eﬃciency (squares), charge (upward-pointing triangles) capacity, and discharge (downward-pointing
triangles) capacity versus time and cycle for a full cell comprising 5 mL of 0.5 M 2,6-DPPEAQ at pH 9 in the negolyte, and 80 mL of 0.4 M potassium
ferrocyanide and 0.1 M potassium ferricyanide at pH 9 in the posolyte. The cell was cycled galvanostatically at 100 mA/cm2 between 1.5 and 0.5 V andeach half-cycle ended with a potentiostatic hold until the magnitude of the current density fell below 2 mA/cm2. Inset: capacity versus cell voltage at the
1st, 10th, 100th, and 480th cycle. Adapted with permission from ref 54. Copyright 2019 John Wiley and Sons.

the formation of an anthrone dimer, which is redox-inactive.
Interestingly, in the case of 2,6-dihydroxyanthrone (DHA)
formation at pH 14, capacity recovery is possible because DHA
can be reoxidized to DHAQ under aerobic conditions. Even in
the presence of oxygen, however, such recovery is only partial,
because DHA irreversibly converts to its dimer, (DHA)2. The
ﬁrst demonstration of the recovery eﬀect, by Goulet and Tong etal.,32 showed a recovery of 70% of the lost capacity. Because(DHA)2 has a higher reduction potential, has a 2-fold lower
speciﬁc capacity than DHAQ, and is unstable at pH 14, some
capacity loss is unavoidable. It was hypothesized that minimizing
anthrone formation by avoiding a high concentration of reduced
species, or high SOC, during cycling would reduce the capacity
fade rate of a DHAQ-based cell. Consistent with this hypothesis,
restricting the maximum SOC from ∼100% to 88% cut thecapacity fade rate by almost 2 orders of magnitude, from 5.6%/
day to 0.14%/day without utilizing the recovery eﬀect.
Computational chemistry indicates that the propensity foranthrahydroquinones to disproportionate into the corresponding anthrone and anthraquinone correlates inversely to the
reduction potential of the anthraquinone.32 This result is
consistent with a recent study showing that a full cell featuring anegolyte with polyethylene glycol-functionalized anthraquinone(PEGAQ), with a higher reduction potential than DHAQ,
exhibits a slower capacity fade rate from anthrone and anthrone
dimer formation.82 Speciﬁcally, PEGAQ at pH 7 has a 250 mV
higher reduction potential than DHAQ at pH 14, but a capacity
fade rate about 10 times smaller, at 0.5%/day.82 It is interesting
to note that a number of full cells with benzoquinone
reactants35,36 of lower reduction potentials than DHAQ exhibit
accordingly greater capacity fade rates, up to 35%/day in one
case. However, the degree to which these higher capacity faderates reﬂect higher rates of formation of anthrone-like species is
an open question, requiring quantitative study of the correlation
between capacity fade and chemical decay and controlling for
pH.

AQDS-CO2 adduct would be hydrolyzed to form semianthraquinone, which would then dimerize into semianthraquinone−anthraquinone units with an overall Coulombic capacityof 1.5 electrons per molecule of AQDS. In neutral to basic pH,
they expect that hydroxyanthraquinone−anthraquinone dimers
would be formed, resulting in an overall capacity of one electron
per molecule of AQDS. The universality of this capacityrestriction mechanism is however in doubt, as Chen et al.
accessed over 95% of the theoretical capacity of AQDS in anacidic AQDS/Br cell.88 Such a high capacity is consistent with anumber of other reports of acidic,110 neutral,82 and alkaline33
AQDS or anthraquinone-based chemistries in which electrochemical access of ≥90% of the full 2-electron Coulombic
capacity has been demonstrated.
Several studies have explored the use of AQDS and its
derivatives52,110 in full cells as a platform for mechanistic analysis
of current−voltage relationships in RFBs,88 partitioning voltage
losses within the full cell stack,89 and demonstrating thefeasibility of RFBs charged by solar cells.111−113 However,beyond the contribution of bromine crossover,88,114 the origin of
capacity fade in AQDS/Br cells has not been investigated indetail, although hydroxylation,110 and disproportionation108 of
the reduced anthrahydroquinone to anthraquinone and redoxinactive anthrone, have been hypothesized.
Owing in part to potential safety issues raised by the use ofbromine in acidic AQDS-Br full cells, and the facile crossover of
uncharged Br2 species across most ion-exchange membranes,
2,6-dihydroxyanthraquinone (DHAQ) was developed as analternative negolyte reactant at pH 14 and demonstrated in a fullcell against a posolyte comprising potassium ferrocyanide.11,115
A capacity fade rate of 8%/day was observed, and was attributed
to molecular decomposition in a subsequent study.33 By jointly
using electrochemical cycling and chemical analysis, Goulet andTong et al.32 quantitatively demonstrated that progressive loss of
capacity during DHAQ cycling is primarily due to thedisproportionation of 2-electron-reduced DHAQ followed byI

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

Table 2. Summary of Results from Full Cells Containing Viologen, Nitroxide Radical and Iron Coordination Complex-Based
Reactantsa
capacity fade rate
nameauthor/year%/cycle%/day

electron conc. (M)redox potential (V vs SHE)

pH

Diquat 5/FcNCl
MV/4-HO-TEMPO(NPr)2TTZ/NMe-TEMPO
MV/TEMPO polymer
MV/TEMPTMA(SPr)2V/KI
BTMAP-Vi/BTMAP-Fc
BTMAP-Vi/BTMAP-Fc(NPr)2V/FcNCl(NPr)2V/N-Me-TEMPO[(Me)(NPr)V]Cl3/FcNCl
MV/FcNCl
Phenazine-TEMPO combimolecule
Zn2+/g+-TEMPOBTMAP-Vi/TMAP-TEMPO

Huang/2018128
Liu/201562
Luo/2018122
Janoschka/2015120
Janoschka/2016121
DeBruler/2018127
Beh/201790
DeBruler/2017129
Hu/2018130
DeBruler/2017129
Hu/2017125
Winsberg/2016119
Chang/2017117
Liu/2019118

Zn2+/SO4-TEMPO
Zn2+/Im-TEMPO

Winsberg/2017124
Chang/2019123

0.2
0.1
0.03
0.25
0.026
0.01
0.00570.0011
0.01
0.005
0.18
0.01
0.011
0.046
0.007
0.015
0.0058
0.11

0.8
27.5
2.25
7.5
0.27
0.45
0.1
0.033
0.5
0.2
8.8
0.22
0.62
0.7
0.6
0.55
1.38
21.2

0.5
0.5
0.1
0.37
2
0.5
1.3
1.0
0.5
0.5
0.5
0.5
0.01
0.2
0.1
1.5
0.035
0.4

−0.51/0.61−0.45/0.80−0.44/1.00−0.45/0.90−0.45/1.00−0.43/0.57−0.36/0.39−0.36/0.39(−0.35,−0.72)/0.61−0.38/1.00(−0.39,-0.78)/0.61−0.45/0.61−0.39/0.81−0.69/0.81−0.38/0.81−0.38/0.81−0.87/0.82−0.69/0.95

7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7

Capacity fade results established by potentiostatic cycling or galvanostatic cycling with potentiostatic ﬁnishes are marked in bold text, whereas
apparent capacity fade results from purely galvanostatic cycling are left in plain text. In the “name” ﬁeld, the active reactant in the negolyte is on theleft and that for the posolyte is on the right. Where there is a CLS, the reactant it contains is marked out in bold, whereas in capacity-balanced
conﬁgurations, both reactants are in plain text. “electron conc.” refers to moles of transferrable electrons per liter in the CLS. “redox potential”
refers to redox couple(s) in the negolyte (left), and that of the posolyte (right), with the CLS marked out in bold. We include hybrid ﬂow/non-ﬂow
batteries with zinc electroplating when the ﬂowing aqueous organic is the capacity-limiting side or is believed to be the capacity-limiting side after
capacity fade begins. MV = methyl viologen, 4-HO-TEMPO = 4-hydroxy-2,2,6,6-tetramethylpiperidin-1-oxyl, (NPr)2TTZ = 4,4′-(thiazolo[5,4d]thiazole-2,5-diyl)bis(1-(3-(trimethylammonio)propyl)pyridin-1-ium) tetrachloride, NMe-TEMPO = 4-trimethylammonium-TEMPO, (SPr)2V =1,1′-bis(3-sulfonatopropyl)-4,4′-bipyridinium, FcNCl = ferrocenylmethyl)trimethylammonium chloride, BTMAP-Vi = bis(3-trimethylammonio)propyl viologen tetrachloride, BTMAP-Fc = bis((3-trimethylammonio)propyl)ferrocene dichloride, (NPr)2V = 1,10-bis[3-(trimethylammonio)propyl]-4,40-bipyridinium tetrabromide, (Me)(NPr)V = 1-methyl-10-[3-(trimethylammonio)propyl]-4,40-bipyridi-nium trichloride, g+-TEMPO =glycidyltrimethylammonium cation-grafted TEMPO, TMAP-TEMPO = 4-[3-(trimethylammonio)propoxy]-2,2,6,6-tetramethylpiperidine-1-oxyl,
SO4-TEMPO = TEMPO-4-sulfate potassium salt, TEMPTMA = N,N,N-2,2,6,6-heptamethylpiperidinyl oxy-4-ammonium chloride, Im-TEMPO =imidazolium-grafted TEMPO.
a

A recent detailed study by Tong et al.98 of the dimerizednaphthoquinone bislawsone in an alkaline cell has demonstrated
that tautomerization, rather than disproportionation, of thehydroquinone chieﬂy accounts for molecular decompositioninduced loss in cell capacity. The correlation between
tautomerization of the hydroquinone to insoluble and/or
redox-inactive di- and tetra-ketone forms and capacity fade
was supported in similarly quantitative fashion to an earlier study
on DHAQ decomposition: by comprehensive electrochemicalcell cycling and post mortem chemical analysis of the electrolyte
via NMR and high-resolution high-performance liquid chromatography (HPLC). Theoretical calculations suggest thatnaphthoquinone susceptibility to tautomerization generally
increases with reduction potential but is also highly sensitive
to functionalization−in particular, the calculations suggest that
oxy-alkyl substitutions impart greater protection against
tautomerization than the hydroxyl groups in bislawsone.
Although it is possible that tautomerization is universally
responsible for capacity fade in naphthoquinones, corroboration
is diﬃcult given the paucity of publications in this area.100,104
Other hypotheses have been proposed, such as nucleophilic
attack either directly by OH−, or after quinone epoxidation fromhydroperoxides created in the presence of trace amounts ofoxygen.104
Notwithstanding these uncertainties, alkaline ﬂow cells
utilizing negolytes featuring quinones with reduction potentials

at least 100 mV higher than that of DHAQ and oxy-alkyl
substitutions (i.e., to limit both anthrone formation andtautomerization) have exhibited among the lowest capacity
fade rates reported in the literature.53,54 Indeed, the current
record for the lowest demonstrated capacity fade rate for any
ﬂow battery, organic or inorganic, in the absence of electrolyte
rebalancing processes belongs to one with a phosphonatefunctionalized anthraquinone (2,6-DPPEAQ) in the negolyte
and a K4Fe(CN)6 posolyte, at 0.014%/day according togalvanostatic cycling with potential hold ﬁnishes (Figure 4).54
A full cell based on a carboxylate-functionalized analogue of
2,6-DPPEAQ (2,6-DBEAQ) exhibited a higher capacity faderate, at 0.04%/day.53 At such low fade rates (≤0.1%/day), themeasured fade rate is often associated with a high relative
experimental uncertainty, as it is highly sensitive to smallchanges in the rates of negolyte leakage and crossover throughpinholes in the membrane. Purely galvanostatic cycling is
unlikely to provide an assessment of the actual capacity fade rate,
as the apparent capacity would be highly sensitive to ﬂuctuations
caused by diurnal temperature changes and drifts in membrane
resistance. Additionally, discerning the origins of capacity fadebecomes challenging, as several weeks to months of testing atroom temperature would be required to accrue enoughdecomposed material for meaningful chemical analysis (e.g.,using NMR and HPLC). Consequently, chemical analysis of
2,6-DBEAQ and 2,6-DPPEAQ after accelerated chemicalJ

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Reviewcell, presumably because the two-electron-reduced viologen
core in the former is more chemically unstable than the oneelectron version.33,70
The extent to which this hypothesis explains capacity fadecannot be discerned from these results, because the cell wascapacity-balanced and no chemical analysis was performed on
cycled electrolytes. Replacing the TEMPO posolyte with one
containing an ammonium-substituted Fc complex((ferrocenylmethyl)trimethylammonium chloride, FcNCl)129
resulted in a drastic reduction in the apparent capacity fade rate,
to 0.5%/day. A reasonable inference from this result is that theredox couple associated with NMe-TEMPO is more chemicallyunstable than that of (NPr)2V. However, this is apparently
contradicted by a subsequent publication in which Hu et al.130
show that accessing only the ﬁrst reduction of (NPr)2V in acapacity-balanced (NPr)2V/NMe-TEMPO full cell results in aneven lower apparent capacity fade rate of 0.2%/day, whereas itshould be at least 50%/day if NMe-TEMPO were the capacitylimiting reactant. These ambiguities underscore the need forestablishing quantitatively rigorous relationships between
capacity fade and molecular decay or crossover through themembrane, where applicable.
A signiﬁcant step toward this goal was taken by Beh et al.,90
who demonstrated extremely low capacity fade rates (≤0.1%/
day) in ammonium-functionalized viologen (Vi)/Fc full cells
and provided evidence that capacity fade is attributable tobimolecular decomposition reactions. In their conﬁguration,
both Vi and Fc were functionalized with positively charged
quaternary ammonium groups, yielding bis(3-trimethylammonio)propyl viologen (BTMAP-Vi), which is identical to(NPr)2V, and bis((3-trimethylammonio)propyl)ferrocene(BTMAP-Fc), respectively. Conventional galvanostatic cycling
indicated a higher apparent capacity fade rate than that frompotentiostatic cycling; this observation was attributed to a slow
increase in membrane resistance. The apparent capacity also
exhibited diurnal, cyclic variations due to changes in ambient
temperature. Both of these observations underscore theincreasing unreliability of purely galvanostatic cycling as a
guide to measuring the actual cell capacity fade rate as it
approaches 0.1%/day from above.
Potentiostatic cycling of a BTMAP-Vi/BTMAP-Fc cell
yielded a temporal capacity fade rate of 0.1%/day for acapacity-balanced cell with reactant concentrations of 1.3 M.
This temporal rate was found to be essentially independent ofcycle period, indicating that capacity fade was primarily timedenominated (i.e., driven by molecular decomposition or
crossover) rather than cycle-denominated. A record low faderate of 0.033%/day for a BTMAP-Fc-limited cell with BTMAPFc concentration of 1.0 M was also found and, based on ex situ
membrane permeability measurements, about half of the overall
capacity fade was attributed to reactant crossover. That thetemporal capacity fade rate increases with concentration both in
this study and earlier reports on Vi/Fc or TEMPO-based
chemistries62,125 provides support for the interpretation thatcapacity fade originates in large part from bimolecular reactions
between reactants. Taken together, these data are consistent
with the notion that capacity fade occurs at a signiﬁcantly slower
rate with ammonium-functionalized reagents owing to increased
Coulombic repulsion aﬀorded by the charged substituents.
Further corroboration in the case of BTMAP reagents waschallenging because, given the fade rates and experimental run
times in the study, total capacity fade is on the order of ∼0.1−1%
only, which approaches the sensitivity of conventional chemical

decomposition by storage at elevated temperatures in both
redox states was used to determine decomposition pathways at
diﬀerent pH values. The suggestion from these studies was that
whereas nucleophilic substitution by OH− was the dominant
decay mechanism for both quinones at pH 14, an intramolecular
reaction with carboxylate acting as the nucleophile dominates inthe decomposition of 2,6-DBEAQ and no analogous intramolecular reaction of 2,6-DPPEAQ was observed at pH 12 atelevated temperatures due to a much weaker nucleophilicity ofthe bulky phosphonate group relative to carboxylate.116
In comparison to much of the RFB literature involving
quinones, the details of molecular decomposition-induced
capacity fade in most full ﬂow cell reports with nitroxide
radical,62,117−124 iron coordination complex,90,125,126 andviologen-based62,90,122,127−130 reactants are less understood.
This is in large part because most full cell studies featuring these
species (Table 2) do not report a CLS. Thus, because initial
Coulombic capacities of the negolyte and posolyte are nominallyidentical, it is impossible, in the absence of other information
about the electrolyte chemistries, to pinpoint the decomposing
electrolyte. In principle, one may avoid this shortcoming via
chemical analysis of both cycled electrolytes; however, suchanalysis is valid only if molecule crossover and electrolyte
leakage are properly accounted for.
Because viologens and ferrocene/nitroxide-radical based
reactants tend to have low (∼−0.4 V vs SHE) and high(∼0.4−1.0 V vs SHE) redox potentials, respectively, they areoften paired as negolyte/posolyte combinations in full cells, withmaximum full cell voltages (1.2−1.4 V) approaching thethermodynamic stability window for water.62,118,120−122,127,130
This pairing is particularly appropriate, as these redox couples
are highly chemically unstable under corrosive conditions, i.e.,
vulnerable, in the oxidized state, to nucleophilic attack by OH−
under alkaline conditions,63,65 and, in the reduced state, toprotonation in acid; consequently in organic RFBs they are used
only under neutral conditions. Nevertheless, moderate to highcapacity fade rates have been reported in the literature. In theﬁrst such study of a viologen/nitroxide-radical full cell in water,
Janoschka et al.120 deployed polymerized MV and TEMPO in a
ﬂow cell. Because of the macromolecular nature of the reactants
used, the viscosity restricted the concentration to well below thesolubility limit, but they were able to use a low-cost microporous
separator rather than an ion-selective membrane. The ﬂow cell
exhibited an apparent capacity fade rate of 0.75%/day.
The vast majority of other full cell reports based on this motif
have used monomeric reactants. Liu et al.62 reported a waterbased methyl viologen/4-HO-TEMPO ﬂow cell with an anionexchange membrane, and an apparent capacity fade rateequivalent to ∼28%/day; however, the cell was capacitybalanced, and no chemical analysis on cycled electrolytes was
performed. DeBruler et al.129 subsequently reported an apparent
capacity fade rate of ∼50%/day for a full cell in which both MV
and TEMPO were functionalized with ammonium substituents,
yielding 1,10-bis[3-(trimethylammonio)propyl]-4,40-bipyridinium ((NPr)2V) and 4-trimethylammonium-TEMPO (NMeTEMPO), respectively. During the cycling, they accessed thesecond reduction of (NPr)2V at −0.75 V vs SHE, resulting in ahigher speciﬁc capacity and average full cell voltage than that of
the one-electron MV/TEMPO ﬂow cell. The use of charged
substituents was meant to prevent dimerization and precipitation of the singly and doubly reduced (NPr)2V, respectively.
However, a higher apparent capacity fade rate was exhibited bythe (NPr)2V/NMe-TEMPO cell than the MV/4-HO-TEMPO
K

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

Figure 5. RFB design of (a) anion-exchange viologen/TEMPO and viologen/Fc chemistries and (b) cation-exchange (SPr)2V/iodide chemistry.
Reprinted with permission from ref 127. Copyright 2018 American Chemical Society.

Table 3. Summary of Full Cell Results with Aza-Aromatic Reactantsa
capacity fade rate
nameauthor/year%/cycle%/day

electron conc. (M)redox potential (V vs SHE)pH

DHPS/Fe(CN)6
ACA/Fe(CN)6
FMN/ Fe(CN)6
V3+/MB

Hollas/201834
Lin/201679
Orita/2016139
Zhang/2019145

0.02
0.013
0.02
0.074

0.68
1.2
0.6
0.76

2.8
1.0
0.48
3.0

−0.81/0.50−0.62/0.50−0.5/0.50−0.26/0.57

14
14
14
−0.5

All apparent capacity fade results were obtained by purely galvanostatic cycling. In the “name” ﬁeld, the active reactant in the negolyte is on the left
and that for the posolyte is on the right. The CLS is marked out in bold and the NCLS is in plain text. “electron conc.” refers to moles oftransferrable electrons per liter of electrolyte. “redox potential” refers to the redox couple in the negolyte (left), and that of the posolyte (right),
with the CLS marked out in bold. DHPS = dihydroxyphenazine sulfonic acid, ACA = alloxazine 7/8-carboxylic acid, FMN = Flavinmononucleotide, MB = methylene blue.
anation complexes with negolyte/posolyte redox potentials
dictated by appropriate choice of ligand.126
Aza-aromatics, which contain nitrogen atoms in the aryl rings,
have received attention as candidate RFB reactants.34,79,138−143
More comprehensive assessments of capacity fade have thepotential to accelerate the discovery and development of longlived aza-aromatic chemistries. Quinoxaline, which comprises a
fused benzene-pyrazine ring, was ﬁrst proposed as a promisingnegolyte reactant by Milshtein et al.138 owing to its highsolubility (∼4.5 M in water, 0.5 M in 1 M KOH) and low
reduction potential of −0.78 V vs SHE at pH 14. A subsequent
patent application144 reports quinoxaline full cells with BQDS
and ferrocyanide posolytes that both exhibited rapid capacity
decay, equivalent to greater than 100%/day. A proof-of-concept
negolyte-limited quinoxaline/oxygen ﬂow cell142 exhibited ahigh apparent capacity fade rate as well, equivalent to ∼10%/
day. However, this ﬁgure was not partitioned into contributions
from molecular decomposition and crossover, the latter of whichis expected to be moderately high for most charge-exclusionbased membranes, given that unsubstituted quinoxaline isuncharged.
Aza-aromatics with charged, hydrophilic side chains havedemonstrated considerably lower capacity fade rates in full cells.
This is the case for bioinspired, quinoxaline-like heterocycles,
based in particular on phenazine34 and phenothiazine (i.e.,comprising either pyrazine or thiazine, respectively, sandwiched
between two benzene rings), but also others based onalloxazine79,139 (Table 3). Hollas et al.34 synthesized dihydroxyphenazine sulfonic acid (DHPS) for use as a negolyte reactant
in an alkaline ﬂow cell, which has the best performance so far
among aza-aromatic negolytes. At pH 14, DHPS has the lowest
reduction potential of any organic negolyte species (−0.81 V vsSHE), a high solubility (1.8 M) and is essentially tied for thelowest apparent capacity fade rate among aza-aromatics in a celland electrochemical assay techniques. Indeed, NMR analysis of
the cycled electrolytes detected no crossed over reactants and noevidence of decomposed BTMAP-Fc.
Developing new or tailored techniques capable of rationalizing full cell capacity fade in terms of speciﬁc mechanisms ofmolecular loss, particularly when the total capacity fade at theend of an experimental run at room temperature is ≤1%, iscritical. It would greatly accelerate the assessment of novel
strategies for molecular engineering of RFB chemistries withimproved cell performance characteristics (e.g., open-circuit
voltage, peak galvanic power density, and volumetric energy
density) that still possess the long lifetimes required forcommercialization. This need is illustrated by a recent study
by DeBruler et al.127 of a new RFB chemistry in which asulfonate-functionalized viologen (1,1′-bis(3-sulfonatopropyl)4,4′-bipyridinium, (SPr)2V) and I−/I3− were used in thenegolyte and posolyte, respectively (Figure 5). This system
features negatively charged redox-active moieties and cationiccounterionsand thus potentially enables both higher maximum power densities and more durable RFBs than previous
cationic MV/TEMPO or Fc conﬁgurations with anioniccounterionsvia the use of higher conductivity, longer-lifetime
cation-exchange membranes.131 The lowest apparent capacity
fade rate attained for a negolyte-limited (SPr)2V/KI cell was
equivalent to 0.45%/day in galvanostatic cycling; however, witha cycling duration ∼2 days long, this amounts to a total capacity
fade of only 1%, and NMR and CV analysis of cycled electrolytes
failed to detect any decomposed or crossed over products. As aresult, the promise of this potentially signiﬁcant innovation is yet
to be demonstrated or fully appreciated. Similar concerns apply
to other advanced organic RFB concepts, such as those based onmolecules comprising combined MV-Fc132 or phenazineTEMPO119 moieties, diﬀerential pH MV-air systems,133 hybrid
metal−organic polymer conﬁgurations,134−137 and iron coordiL

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

unrelated to OH−-initiated hydrolysis products. Rather, they
attributed the new CV peaks to the FMN dimer147 which, theyargued, should have a lower redox potential than thecorresponding monomers owing to energetically favorable
dipole−dipole interactions. This conclusion was supported inpart by UV−vis absorption measurements that showedattenuated π → π* transitions from FMN monomers overtime. CVs of the cycled electrolyte failed to detect evidence forreactant crossover, indicating that molecular decomposition or
diminution of redox activity is the main mechanism of active
reactant loss. The relative contributions of irreversible FMN
dimer formation and hydrolysis to capacity fade were notexplored in detail; however, this ﬁnding of irreversible FMN
dimerization raises the question of whether both mechanisms
might be generally operative in alkaline alloxazine-based cells.
There is some support for this notion from (i) CVs of ACA andcharging proﬁles of ACA/Fe(CN)6 full cells79 displaying
features similar to what was observed for FMN, and consistent
with ACA dimer formation and (ii) the capacity fade rate for the
ACA/Fe(CN)6 cell being twice that of FMN/Fe(CN)6, for aconcentration of ACA that is twice that of FMN − this isconsistent with capacity fade originating from a mixture of ﬁrst
order (OH−-initiated hydrolysis) and second-order (dimerization) decay mechanisms.

against a ferrocyanide posolyte (0.68%/day). Capacity loss wasdeemed unlikely to be accounted for by DHPS decomposition,
given that an ex situ NMR study showed no decomposition of
the oxidized form after 3 weeks of storage at pH 14; however, asimilar study of the reduced form, which is more diﬃcult due tothe need to protect from oxidation by atmospheric O2 invasion,
was not performed. Visual inspection of the ion-exchange
membrane after cycling showed intense discoloration, which led
the authors to suggest that DHPS interaction with or crossover
through the ion-exchange membrane was the more likely culprit.
This hypothesis was supported by a concentration-independent
temporal capacity fade rate, which is consistent with amechanism for molecule loss with a ﬁrst order dependence inconcentration.
Replacing one of the nitrogen atoms in phenazine with a more
electron-withdrawing sulfur atom yields phenothiazine, andZhang et al. have recently shown that methylene blue (MB), anammonium-functionalized phenothiazine, is a chemically stable
derivative with a high enough redox potential under acidic
conditions (0.57 V vs SHE) for use in the posolyte of a fullcell.145 DFT calculations suggested that the solvation of MB in
acetic acid should be greater than in water. This enabled an∼10× higher MB solubility in a mixture of acetic acid andsulfuric acid (1.8 M) than in water alone (<0.2 M). A full cellwith an MB posolyte as the CLS and a negolyte based on the
V2+/V3+ redox couple displayed moderate apparent capacity
fade rates of 0.52 and 0.76%/day for MB concentrations of 1.2
and 1.5 M, respectively. Similar to the case of phenazine, thesefade rates were inconsistent with the apparently extremely highchemical stability of MB (established in this case by symmetriccell cycling, which we discuss in the next section); thus, capacity
loss was attributed to continuous vanadium ion crossover
through the anion-exchange membrane, resulting in thenegolyte becoming the CLS. Based on these results and thoseof other studies involving pyridinic and phenazine-based
reactants,140,143 it appears that a wide variety of aza-aromatic
molecules may be particularly promising for use in future RFBs,
if their chemical stability can be more deﬁnitively establishedand compatible membranes and counter-electrolytes can befound.
Alloxazine-based alkaline full cells exhibit similar temporalcapacity fade rates; however, molecular decay appears to play amore signiﬁcant role than crossover. Lin et al.79 and Orita etal.139 reported cell chemistries with negolytes comprising,respectively, alloxazine 7/8-carboxylic acid (ACA), and ﬂavinmononucleotide (FMN) mixed with nicotinamide (NA) toboost its water solubility. Both ACA and FMN derive from theisoalloxazine backbones of vitamin B2, and the corresponding
full cells with ferrocyanide posolytes exhibited temporalapparent capacity fade rates of 1.2 and 0.6%/day for negolyte
reactant concentrations of 0.5 and 0.24 M, respectively. Lin et al.
attributed capacity fade in ACA to OH−-initiated hydrolysis of
the amidic carbonyls followed by a ring-opening reaction andfurther hydrolysis into redox-inactive species, which, aspreviously discussed, is a mechanism strongly supported byprior chemical studies76−78,146 and an NMR study of thechemical stability of ACA in pH 14 solution.
Orita et al. conducted CV measurements of FMN in 1 M
KOH at 363 K, and observed that a more irreversible redox
couple ∼300 mV lower than that of FMN evolved and grew overthe course of 100 h. Given that alloxazine hydrolysis products
are not redox-active, and the new redox feature was present at
pH 5.5 and 8.6, they concluded that the new CV features were

2.3. Symmetric Cell Cycling: A Technique for Coupling
Capacity Fade to Chemical Decomposition

In full cells, both CLS active reactant chemical decompositionand crossover across the membrane37,88 contribute to capacity
fade. Thus, where reactant permeability is non-negligible,
quantifying the role of molecular decomposition alone to cell
capacity fade is diﬃcult. One potentially simple way of
decoupling molecular decomposition from crossover is byusing a cell in which both the negolyte and posolyte have thesame redox couple at the same concentration; this way, reactant
diﬀusion through the membrane is virtually eliminated becauseof similar chemical compositions on either side of it. Darling andPerry148 implemented the ﬁrst embodiment of this idea in theform of a “double half-cell” conﬁguration in which a 50% SOC
V2+/V3+ electrolyte from a single reservoir was circulated acrossa cell stack that was held at a constant potential such thatelectrolyte oxidation occurred at one electrode and reduction at
the other. In principle, if there are no side reactions, an opencircuit potential of 0 V and constant SOC would be maintainedin the reservoir, diﬀusion-induced reactant and solvent crossoverwould be eliminated, and the measured current would be
proportional to the concentration of vanadium ions. Although
this technique is useful for evaluating the performance of a cellstack, one could not use the measured current as a proxy forreactant concentration because it would be sensitive to changes
in membrane resistance, electron-transfer kinetics and ambient
temperature, in addition to concentration changes frommolecular decomposition.
A route to avoiding these limitations is the symmetric cell
conﬁguration, in which two separate electrolyte reservoirs
contain the same redox couple (Figure 6a), and their Coulombic
capacity is measured with constant-current (Figure 6b) or
constant-voltage cycling. This approach, ﬁrst reported by
Milshtein et al.149,150 for nonaqueous RFB chemistries, enables
the assessment of capacity fade because, in the absence ofsigniﬁcant mass transport overpotentials, capacity is directly
related to reactant concentration. Goulet and Aziz33 presented avariation called the “unbalanced, compositionally symmetric cellMhttps://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

capacity fade rate of the CLS correlates quantitatively withchemical instability of the redox couple contained therein. They
also demonstrated that introducing pauses in cell cycling at
diﬀerent SOC values of the CLS could establish relative
susceptibilities of the diﬀerent redox states of the molecule ofinterest to chemical decomposition (Figure 7c). Symmetric-cell
cycling is therefore a robust but simple means of evaluating
molecular decomposition rates and of informing molecularengineering strategies for developing long-lived aqueous RFB
reactants.
Symmetric cell cycling is relatively rare in the molecular RFB
literature, but it has shed light on chemical stability of and cell
results particularly involving quinone,33,53,54,61,100,151 viologen,33 aza-aromatic,145 ferrocyanide,33,56,64 and TEMPObased118 molecules (Table 4). The single-electrolyte capacity
fade rate revealed by these measurements is expected to provide
a lower limit on the fade rate of a full cell pairing that electrolyte
against another electrolyte. We have found only two exceptionsin the literature54,100 and can hypothesize an explanation in one
case: Ji et al.54 performed symmetric-cell cycling at a diﬀerent
pH than the full-cell cycling reported in the same work. An
extremely low (0.02%/day and lower) full cell capacity fade rate
has been established for 2,6-DPPEAQ/K4Fe(CN)6,54 demonstrating that organic molecules can possess the lifetime
properties necessary for RFB commercialization. Comparably
low apparent fade rates of other species oﬀer the possibility ofmore such demonstrations.56,145,152 Where temporal fade rates
are moderate (0.1−1%/day) or worse, symmetric cell cycling is aparticularly powerful diagnostic tool. The lowest reported fullcell apparent fade rate for AQDS, for instance, is 0.2%/day for aconcentration of 0.5 M,153 which is reasonably consistent withthe symmetric cell fade rate of 0.08 ± 0.02%/day obtained in
potentiostatic cycling for a concentration of 0.1 M.33 Pausing thesymmetric cell at 100% SOC resulted in capacity loss equivalent
to 0.2%/day, whereas a similar hold at 0% SOC resulted in no

Figure 6. (a) Schematic of a symmetric ﬂow cell during charging of a
TEMPO-based molecule in nonaqueous electrolyte. (b) Potential vscharge for selected cycles of symmetric cell cycling of 50 mM total
active species in 0.5 M LiBF4/PC. Reprinted with permission from ref
149. Copyright 2016 Elsevier.

method”, in which the electrolytes were compositionallysymmetric but volumetrically unbalanced, so that one side is
the CLS during both charge and discharge, and the other the
NCLS (Figure 7a). By implementing constant-potential cycling
with this setup (Figure 7b), they demonstrated that the temporal

Figure 7. (a) Schematic of unbalanced, compositionally symmetric ﬂow cell. (b) Example of symmetric cell potentiostatic cycling protocol. (c)Semilog plot of unbalanced, compositionally symmetric cell cycling of 50% SOC 0.1 M DHAQ in 1 M KOH with cycling pauses at diﬀerent SOCs.
Temporal capacity fade rates from slopes of each cycling segment are denoted in black text. Capacity loss and SOC of CLS during 24-h pauses aredenoted in blue text. Reprinted with permission from ref 33. Copyright 2018 Electrochemical Society.
N

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

Table 4. Summary of Symmetric Cell Results, Arranged by Class of Moleculea
capacity fade rate
namequinones
DHAQ
2,3-HCNQ
2,6-DBEAQ

author/year%/cycle%/day

electron conc. (M)redox potential (V vs SHE)

pH

Goulet/201833
Wang/2018100
Kwabi and Lin/201853

0.14
0.12
0.001
0.001
0.00020.00035
0.001
0.03
0.7

6
7.3
0.008
0.0075
0.007
0.02
0.08
0.48
70

0.2
0.2
0.2
1.3
0.2
0.2
0.2
1.1
0.05

−0.68
−0.53
−0.52
−0.52
−0.52
−0.47
0.20
0.85
0.89

14
14
14
14
12
13
0
0
7

0.000014
0.015

0.0016
1.5

0.1
0.1

−0.36
−0.45

7
7

0.00077

0.016

2

0.57

−0.5

Luo/201856

0.24
0.026
0.008
0.00027

8
21
2.3
0.7
0.017

0.1
0.2
0.2
0.2
0.5

0.50
0.50
N/A
0.39
0.39

14
14
10
7.2
7

Liu/2019118

0.0076

0.55

0.1

0.81

7

2,6-DPPEAQ
Ji/201954
AQDS
Goulet/201833
DHDMBS
Murali/201861
TMHQ
Drazevic/2019151
viologens
BTMAP-Vi
Goulet/201833
Methyl-Vi
aza-aromatics
MB
Zhang/2019145
iron coordination complexes
K4Fe(CN)6
Goulet/201833
Luo/201764(NH4)4Fe(CN)6
nitroxide radicals
TMAP-TEMPO

Capacity fade results established by potentiostatic cycling or galvanostatic cycling with potentiostatic ﬁnishes are marked in bold text, whereas
purely galvanostatic results for apparent capacity fade are left in plain text. “electron conc.” refers to moles of transferrable electrons per liter ofelectrolyte. TMHQ = tetramorpholinohydroquinone.

aand sensitivity to SOC of these and other hypothesized decay
mechanisms, such as dimerization and hydrolysis of alloxazines,79,139,147 and gem-diol formation and nucleophilic
Michael reactions in quinones.59
As with full cell cycling, however, deploying the appropriate
symmetric cell cycling protocol is critical if one intends to ensurethat measured capacities most closely reﬂect actual amounts ofactive reactants. First, the two electrolytes should ideally not be
capacity-balanced, in order that only one electrolyte (the CLS)
experiences the excursions in SOC prescribed during electrochemical cycling. This permits direct attribution of capacity fademeasured at the cell level to a single electrolyte without the needfor further partitioning, and eases ex situ chemical analysis
because, under ideal circumstances, decomposition of moleculesin the NCLS will not contribute to capacity fade. Second, cycling
with potential holds to access the maximum capacity (i.e., either
potentiostatic, or galvanostatic with potential hold ﬁnishes) ismore accurate than purely galvanostatic cycling, as the latter isvulnerable to spurious changes in apparent capacity caused bytemporal changes in mass transport, ambient temperatureﬂuctuations, and drifts in membrane resistance and electrodekinetics.33,90 Caution must be exercised in interpreting results
obtained under such circumstances. Indeed, extremely low
apparent capacity fade in purely galvanostatic symmetric cell
cycling tests56,145 may be the result of a fortuitous combinationof progressively lower cell resistance and signiﬁcant actual
capacity loss, rather than high chemical stability of the redox
couple in question.
Capacity retention results from hybrid full-symmetric cells, inwhich the two electrolytes are identical, each comprising both
high and low-potential redox couples,56 cannot necessarily beinterpreted directly in terms of reactant stability. Much like a fullperceptible change in capacity, leading to the conclusion that the
reduced form of AQDS is more susceptible to decomposition
than the oxidized form. This conclusion is corroborated by an
NMR study showing a 40% reduction in AQDS concentration
after 2 weeks of storage at 45 °C in the reduced form, while theoxidized form stayed intact under identical treatment. Regarding
DHAQ, combinations of full and symmetric cell experiments
together with NMR and HPLC analysis32,33 have shown boththat the 5−8%/day fade rate observed in full cell cycling isquantitatively attributable to chemical decomposition and thatthe disproportionation of reduced DHAQ to its corresponding
quinone and anthrone forms, and the dimerization of the latter isthe capacity-loss mechanism.
It remains an intriguing open question whether symmetric cell
cycling combined with chemical analysis can discern the link
between capacity fade and molecular decomposition in othermolecules and molecule classes. For instance, although Tong etal. provided evidence that tautomerization was the main mode ofbislawsone decomposition,98 they did so with a full rather thansymmetric cell, and about 46% of the cell’s capacity loss could
not be accounted for by chemical decomposition. Wang et al.,100
on the other hand, showed that a 2-hydroxy-3-carboxy-1,4naphthoquinone (2,3-HCNQ) symmetric cell loses capacity
faster at 100% SOC (13.4%/day) than 0% SOC (0.6%/day) but
did not perform follow-up chemical analysis, and thus did notverify that tautomerization of the reduced form was the main
capacity-loss mechanism. Symmetric cell cycling has been
employed in attempting to elucidate the proto-desulfonation
reaction in DHDMBS,61 the dimerization and protonation of
reduced MV and BTMAP-Vi, respectively,33 and nucleophilic
attack by water in TEMPO-based chemistries.118 In principle, it
can be applied to great eﬀect to precisely elucidate the kinetics
O

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review

Figure 8. Plot showing full cell capacity fade rate in%/day vs redox potential (vs SHE) for viologen (red), aza-aromatic (green), iron coordination
complex (orange), TEMPO (blue), and quinone-based (black) chemistries. Posolytes are indicated using diamonds, and negolytes in circles, whereas
points depicting chemistries with electron concentrations ≥1.0 M are larger than those <1.0 M. Points depicting capacity fade results of chemistries
that have been tested using potentiostatic or combined potentiostatic-galvanostatic cycling have been ﬁlled, whereas those depicting apparent capacity
fade results obtained with only galvanostatic cycling have been left empty. Capacity-balanced cells where both sides contain organics/organometallics
are marked out with a horizontal line connecting negolyte and posolyte chemistries.

ferricyanide solution at pH 14 and the carbon electrode evenunder open-circuit conditions resulted in the spontaneous
reduction of the former to ferrocyanide. This had the eﬀect ofmaking both electrolytes capacity-limiting after the passage ofenough time in the symmetric cell. The precise mechanistic
details of ferricyanide reduction are yet to be understood,
however there is evidence that it occurs faster with increasing
concentration of OH− and light exposure. Given that both theconcentration of OH− and light exposure have been implicatedin accelerating the replacement of the cyanide ligand in Fe(CN)6
with hydroxl,64,154,155 determining the interplay between these
two mechanisms and their relative contributions to capacity fadeis critical. Understanding the implications of this interplay isparticularly urgent for full cell cycling, because most posolytes inalkaline organic ﬂow cells comprise the Fe(CN)63‑/4− redox
couple (see Table 1). Drazevic et al.151 speculated that a similar
carbon-mediated reduction mechanism may be at play insymmetric cell cycling of TMHQ, based on the observation thatex situ LC-MS analysis of cycled electrolytes cannot rationalize
total capacity fade in terms of irreversible tautomerization andhydrolysis of oxidized TMHQ. Future aqueous organic RFB
research should therefore focus on the development of methods
that can comprehensively map out the landscape of chemicaland electrochemical decomposition in these more complex
chemistries.
It is important to note that, in practical deployment, RFBs will
operate on variable power-based proﬁles rather than cyclic
galvanostatic/potentiostatic proﬁles in which the average SOCof the CLS is close to 50%. Consequently, understanding how
capacity fade in a ﬂow cell varies with SOC by introducing
pauses at diﬀerent SOC values33 or restricting the SOC range
experienced by the CLS32 may be useful starting points inpredicting RFB longevity in real-world applications. In any case,cell, this method features a nonzero average cell potential during
charge/discharge. One oxidizes/reduces the high-potential
redox couple in one electrolyte while reducing/oxidizing the
low-potential couple in the other, with the complementary redox
couples acting as “dormant” spectators. Low apparent capacity
fade rates in this case are not suﬃcient to establish high chemical
stability of the redox couples involved because at low or high
SOC values there could be asymmetric crossover33 of dormantmaterial, in such a way as to compensate for decomposed species
in the side that is or becomes capacity-limiting; this replenishment would mask the actual capacity fade. Additionally, capacity
fade could comprise fractional decomposition among four redox
species, each of which is experiencing potential oscillations overtwo distinct ranges, making useful interpretation diﬃcult. This isparticularly exacerbated when the cell polarity is switched after anumber of cycles and previously dormant active material isaccessed electrochemically. Luo et al.56 argued that switching
the polarity midway through a cycling regimen demonstrates
that species have not been lost through crossover; however,extremely small errors in exactly balancing the original volumes
on the two sides could vitiate this argument.
In cases where there are deleterious interactions between theelectrolyte-contacting materials and redox-active molecules,
symmetric cell cycling may yield an artiﬁcially high capacity faderate that underestimates the actual chemical stability of theredox couple under consideration. This is the case for cycling of
the ferro-/ferricyanide redox couple on activated carbon paper
electrodes at high pH,33 and may also apply to tetramorpholinoquinone (TMHQ).151 In the case of unbalanced compositionally symmetric cell cycling of Fe(CN)6 at pH 14, Goulet andAziz observed an acceleration in the instantaneous capacity faderate, from <10%/day up to 45−67%/day.33 UV−vis and CV
measurements jointly established that contact between
Phttps://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CRtranslating capacity fade measured at the laboratory-scale to apractical RFB is a complex undertaking, and will require further
study as load proﬁle standards for grid-scale storage devices are
adopted.

Review

technique has been successful for cell chemistries that lose
capacity at >0.1%/day, it is limited in doing the same forlonger lasting chemistries because, for more stablechemistries, it becomes increasingly diﬃcult to run
experiments for suﬃcient durations to reach thesensitivity limits of most chemical analysis tools.
• As shown in Tables 1−3, there is a poor correlation
between time- and cycle-denominated capacity fade rates.
We therefore expect the cycle-denominated capacity faderate to be a poor predictor of RFB reactant chemical
lifetime under conditions of practical deployment unlessproven otherwise.
• A signiﬁcant number of chemical stability studies of
proposed RFB reactant are conducted using galvanostatic
cycling, which is inadequate for quantifying low or
extremely low fade rates due to possible artifacts
introduced by temporal changes in mass transport,
ambient temperature ﬂuctuations, and drifts in membrane
resistance or electrode kinetics. Far more reliable for thispurpose are potentiostatic cycling or galvanostatic cycling
with potential holds at the end of charging and ofdischarging of every n th cycle, where only the capacities
measured potentiostatically are taken seriously.
• A technique that more tightly couples capacity fade tochemical decomposition is potentiostatic cycling ofvolumetrically unbalanced, compositionally symmetriccells, along with holds at various SOC to evaluate
temporal fade rate vs SOC. It has been used to elucidate
decomposition mechanisms in a number of quinonebased chemistries32,33,98,151 and should be extended toother molecular classes.

3. SUMMARY
We list and elaborate upon key ﬁndings of this review below:• As shown in Figure 8, which shows the full cell capacity
fade rate vs redox potential for chemistries reported inTables 1−3, quinone and viologen-based chemistries are
the most studied for use in aqueous organic negolytes, andfull cells featuring them as the CLS exhibit a wide range offade rates. In most cases, capacity fade results in large partfrom chemical decomposition of active reactants, themitigation of which through synthetic modiﬁcation hasreceived intense research attention.
• For viologens, functionalization with charged ammonium
groups considerably slows down viologen dimerization,
enabling a low capacity fade rate in the case of a BTMAPVi/BTMAP-Fc cell. Quinones are susceptible to a large
number of decay mechanisms (nucleophilic addition/
substitution, disproportionation, dimerization, and tautomerization); however, oxy-alkyl substitutions, as in 2,6DPPEAQ and 2,6-DBEAQ, appear to suppress tautomerization and disproportionation under alkaline conditions,and aﬀord the only currently demonstrated pathway to
extremely low fade rates in full cells.
• Negolytes based on aza-aromatic molecules have received
less attention, but are potentially promising RFB reactantsin terms of chemical stability. Capacity fade in phenazine
and phenothiazine-based full cells appears to originate inlarge part from crossover and membrane incompatibility
issues; however, independent assessments of reactant
chemical stability have not been thorough enough toestablish this conclusion.
• Aqueous organic or organometallic posolyte development
lags far behind negolyte development: only one capacitylimiting organic posolyte, BTMAP-Fc,90 which has arelatively low redox potential, has demonstrated a low
temporal capacity fade rate (less than 0.1%/day). In thecase of quinones, this observation is consistent withtheoretical calculations suggesting that quinones with
reduction potentials greater than 0.9 V vs SHE are highly
susceptible to irreversible nucleophilic attack by water.
For acidic negolyte chemistries, the inorganic bromine/bromide redox couple is thus frequently used in the
posolyte, whereas Fe(CN)63‑/4− is nearly ubiquitous
against organic negolytes in alkaline conditions.
• TEMPO derivatives, used most frequently in neutral pH,
generally exhibit moderate-to-high fade rates in full cells,
presumably owing to nucleophilic attack by water. In
contrast to Fc, functionalization with ammonium
groups118,121 does not drastically reduce fade rates tolow or extremely low levels. Both of these statements are
consistent with a recent report118 demonstrating that thetemporal fade rate of ﬂow cells containing an ammoniumfunctionalized TEMPO derivative is independent ofconcentration (see Table 2).
• Combining electrochemical full cell cycling with ex situ
chemical analysis of cycled electrolytes has been acommon means of illuminating the relationship between
capacity fade and chemical decomposition. Whereas this

4. CONCLUSIONSThe potentially long discharge duration at rated power that is
attainable with RFBs makes them promising for surpassing
lithium ion batteries for safe, low-cost, large-scale storage fordischarge durations of several hours or longer. Aqueous organic
RFBs have potentially lower cost than their vanadium-based
counterparts; however, this depends on ﬁnding negolyte andposolyte chemistries that achieve low levelized cost through a
combination of long chemical and crossover lifetime, low
synthetic cost at scale, and adequate safety, energy density, andpower density. Because there is a reasonably well-known ﬂoor tothe mass-production cost of organic chemicals, attaining anddemonstrating long chemical lifetime is the greatest present
challenge for aqueous organic RFBs.
A signiﬁcant amount of research attention needs to be paid tounderstanding the mechanisms responsible for capacity fade, asa basis for rational engineering of robust RFBs. Understanding
chemical decomposition, in particular, is important because
organic and organometallic redox couples tend to be less stable
than purely inorganic ones. This task is challenging, because themain molecules are susceptible to several decomposition
mechanisms in water that incur loss or diminution of redox
activity, such as nucleophilic reactions, disproportionation,
dimerization and tautomerization. In ﬂow cells, these pathways
may be simultaneously operative, with overlapping time scales
and unknown rates and reaction orders.
Based on our survey of the literature, we surmise thefollowing:• For all aqueous organic electrolytes for which we haveobserved suﬃcient experimentation to distinguishQ

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review“low” or “extremely low” fade rates are 2,6-DBEAQ/
Fe(CN)6 and 2,6-DPPEAQ/Fe(CN)6, both of whichwere negolyte-limited. K4Fe(CN)6 is challenging as an
RFB posolyte because it has limited solubility (∼0.8 M at
pH 756,156 and lower at higher pH), its crossover andchemical decomposition rates can be signiﬁcant, and its
mass-production cost, though well below that of
vanadium, is a concern. Bromine is likewise problematicas an acidic posolyte because it has a high crossover rateand high cost of safe containment. This highlights theneed for further investigation into low-cost, high-potential
redox couples possessing high chemical stability.
• Because increasing reactant concentration reduces thelevelized cost of RFB storage for a given chemistry,
research into highly soluble organics is important. All fullcell chemistries for which low or extremely low capacity
fade rates have been established have featured CLS
reactant concentrations equivalent to ≥1 M of transferrable electrons. Other chemistries have this attribute but
exhibit moderate fade rates (AQDS, DHPS, PEGAQ,
bislawsone, and MB).

between cycle-denominated fade and time-denominatedfade, it has been found that the amount of lost capacity isessentially independent of the number of charge−
discharge cycles to which a species is subjected. Rather,
capacity fade is primarily time-denominated. The
temporal fade rate can depend on molecular concentrations, pH and SOC. Thus, the capacity fade for new
electrolytes should be presumed to be primarily timedenominated unless a cycle-denominated component isproven to be signiﬁcant.
• Irreversible capacity fade can arise from reactant crossoveror decomposition, or electrolyte leakage. Apparent
capacity fade that is recoverable, which can arise from avariety of other eﬀects such as a side reaction putting thetwo sides out of balance, can prevent direct attribution of
apparent capacity fade measured at the cell level to theseirreversible capacity fade mechanisms.
• Deﬁnitively establishing molecular decomposition rates
less than about 0.1%/day requires careful experimentation and simple experimental designs in which interference from competing eﬀects such as crossover are
suppressed. Results are more convincing if close
correlation of capacity fade to chemical decompositionis established by chemical analysis, but this becomesincreasingly diﬃcult with lower fade rates. Symmetric-cell
cycling can be interpreted more readily than full-cell
cycling due to the suppression of contamination issues
associated with reactant crossover. Volumetrically unbalanced compositionally symmetric cell cycling (Figure7) is recommended because the fade rate can be
attributed to the SOC of only the CLS.
• Inferences about molecular decomposition mechanisms
can be made by holding both oxidized and reduced formsof reactants in the presence or absence of electrolytecontacting materials, or at various pH values, durations
and temperatures, followed by chemical analysis using
such techniques as cyclic voltammetry, NMR, and mass
spectrometry. Accelerated ex situ chemical decomposition
at elevated temperatures may permit elucidation of slower
capacity fade rates at lower operating temperatures, but
this is valid only if the dominant degradation mechanism
does not change within the chosen temperature range.
• The vast majority of aqueous-soluble redox-active organicand organometallic compounds reported to date havehigh (>1%/day) or moderate (0.1−1%/day) fade rates(see Tables 1−4), which are too high for RFB project
lifetimes anticipated to last decades. Only a few have low(0.02−0.1%/day) fade rates, and very few (2,6-DPPEAQ,
2,6-DBEAQ, BTMAP-Fc, BTMAP-Vi, (NH4)4Fe(CN)6and MB) have extremely low (≤0.02%/day) establishedor apparent capacity fade rates. Only 2,6-DPPEAQ, 2,6DBEAQ, BTMAP-Vi, and BTMAP-Fc have had their faderates established by potentiostatic cycling. Some of these(2,6-DPPEAQ, 2,6-DBEAQ, and BTMAP-Vi) have been
established in symmetric-cell cycling and two (2,6DPPEAQ and BTMAP-Fc) have been established as the
CLS in full-cell cycling.
• Posolyte development deserves increased attention. No
posolyte-limited full cell conﬁgurations with an opencircuit voltage greater than or equal to 1.0 V and atemporal fade rate less than 0.1%/day have been reported.
The only full cells with open-circuit voltages ≥1.0 V with

AUTHOR INFORMATION
Corresponding Author

Michael J. Aziz − Harvard School of Engineering and Applied
Sciences, Cambridge, Massachusetts 02138, United States;
orcid.org/0000-0001-9657-9456; Email: maziz@
harvard.edu
Authors

David G. Kwabi − Department of Mechanical Engineering,
University of Michigan, Ann Arbor, Michigan 48109, UnitedStates; orcid.org/0000-0003-3663-8789
Yunlong Ji − Department of Chemistry and Chemical Biology,
Harvard University, Cambridge, Massachusetts 02138, UnitedStates; orcid.org/0000-0002-1124-4209
Complete contact information is available at:https://pubs.acs.org/10.1021/acs.chemrev.9b00599
Notes

The authors declare no competing ﬁnancial interest.
Biographies
David G. Kwabi obtained a B.S.E in Mechanical Engineering and acertiﬁcate in Materials Science and Engineering at Princeton University
in 2010. He carried out M.S. and Ph.D. studies in MechanicalEngineering at the Massachusetts Institute of Technology from 2010
until 2016, under the supervision of Prof. Yang Shao-Horn, in the ﬁeldof Li−O2 batteries. He subsequently joined the group of Prof. Michael J.
Aziz at Harvard University as a postdoctoral scholar, where he worked
on aqueous organic redox-ﬂow batteries. In January 2019, he joined the
Mechanical Engineering department at the University of Michigan as an
assistant professor. His research is in the investigation of fundamental
reaction mechanisms and morphological transformations in advanced
electrochemical energy storage systems such as conversion and ﬂow
batteries.
Yunlong Ji obtained his B.A. degree in Chemistry from Central South
University in 2010. He then pursued his graduate studies in organic
synthesis and ﬂuorine chemistry at Shanghai Institute of Organic
Chemistry, Chinese Academy of Sciences under the supervision of Prof.
Jichang Xiao. In 2015, he started his postdoctoral research in organic
semiconductor materials with Prof. Jian Li at Arizona State University.
Rhttps://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review(8) Wang, Q.; Ping, P.; Zhao, X.; Chu, G.; Sun, J.; Chen, C. Thermal
runaway caused fire and explosion of lithium ion battery. J. Power
Sources 2012, 208, 210−224.
(9) Li, Z.; Pan, M. S.; Su, L.; Tsai, P.-C.; Badel, A. F.; Valle, J. M.; Eiler,
S. L.; Xiang, K.; Brushett, F. R.; Chiang, Y.-M. Air-Breathing Aqueous
Sulfur Flow Battery for Ultralow-Cost Long-Duration Electrical
Storage. Joule 2017, 1, 306−327.
(10) Darling, R. M.; Gallagher, K. G.; Kowalski, J. A.; Ha, S.; Brushett,
F. R. Pathways to low-cost electrochemical energy storage: acomparison of aqueous and nonaqueous flow batteries. Energy Environ.
Sci. 2014, 7, 3459−3477.
(11) Lin, K.; Chen, Q.; Gerhardt, M. R.; Tong, L.; Kim, S. B.;
Eisenach, L.; Valle, A. W.; Hardee, D.; Gordon, R. G.; Aziz, M. J.;
Marshak, M. P. Alkaline quinone flow battery. Science 2015, 349, 1529−
1532.
(12) Skyllas-Kazacos, M.; Chakrabarti, M. H.; Hajimolana, S. A.;
Mjalli, F. S.; Saleem, M. Progress in Flow Battery Research andDevelopment. J. Electrochem. Soc. 2011, 158, R55−R79.
(13) Skyllas-Kazacos, M.; Kazacos, G.; Poon, G.; Verseema, H. Recent
advances with UNSW vanadium-based redox flow batteries. Int. J.
Energy Res. 2010, 34, 182−189.
(14) Skyllas-Kazacos, M.; Kasherman, D.; Hong, D.R.; Kazacos, M.
Characteristics and performance of 1 kW UNSW vanadium redox
battery. J. Power Sources 1991, 35, 399−404.
(15) Moseley, P. T.; Garche, J. Electrochemical Energy Storage for
Renewable Sources and Grid Balancing, 1st ed.; Elsevier: Amsterdam,
2014.
(16) Soloveichik, G. L. Flow Batteries: Current Status and Trends.
Chem. Rev. 2015, 115, 11533−11558.
(17) Kreutzer, H.; Yarlagadda, V.; Van Nguyen, T. Performance
Evaluation of a Regenerative Hydrogen-Bromine Fuel Cell. J.
Electrochem. Soc. 2012, 159, F331−F337.
(18) Braff, W. A.; Bazant, M. Z.; Buie, C. R. Membrane-less hydrogen
bromine flow battery. Nat. Commun. 2013, 4, 2346.
(19) Andrews, J.; Seif Mohammadi, S. Towards a ‘proton flow battery:
Investigation of a reversible PEM fuel cell with integrated metal-hydridehydrogen storage. Int. J. Hydrogen Energy 2014, 39, 1740−1751.
(20) Savinell, R. F.; Wainright, J. S. Iron ﬂow batteries. U.S. Patent
9,559,375, 2017.
(21) Petek, T. J.; Hoyt, N. C.; Savinell, R. F.; Wainright, J. S. Slurryelectrodes for iron plating in an all-iron flow battery. J. Power Sources
2015, 294, 620−626.
(22) Tucker, M. C.; Phillips, A.; Weber, A. Z. All-Iron Redox Flow
Battery Tailored for Off-Grid Portable Applications. ChemSusChem
2015, 8, 3996−4004.
(23) Rajarathnam, G. P.; Vassallo, A. M. The Zinc/Bromine Flow
Battery: Materials Challenges and Practical Solutions for Technology
Advancement; Springer: Singapore, 2016.
(24) Li, B.; Nie, Z.; Vijayakumar, M.; Li, G.; Liu, J.; Sprenkle, V.;
Wang, W. Ambipolar zinc-polyiodide electrolyte for a high-energy
density aqueous redox flow battery. Nat. Commun. 2015, 6, 6303.
(25) Weng, G.-M.; Li, Z.; Cong, G.; Zhou, Y.; Lu, Y.-C. Unlocking thecapacity of iodide for high-energy-density zinc/polyiodide and lithium/polyiodide redox flow batteries. Energy Environ. Sci. 2017, 10, 735−741.
(26) Gong, K.; Ma, X.; Conforti, K. M.; Kuttler, K. J.; Grunewald, J. B.;
Yeager, K. L.; Bazant, M. Z.; Gu, S.; Yan, Y. A zinc-iron redox-flowbattery under $100 per kWh of system capital cost. Energy Environ. Sci.
2015, 8, 2941−2945.
(27) Leung, P. K.; Ponce-de-León, C.; Low, C. T. J.; Shah, A. A.;
Walsh, F. C. Characterization of a zinc-cerium flow battery. J. Power
Sources 2011, 196, 5174−5185.
(28) Chen, H.; Cong, G.; Lu, Y.-C. Recent progress in organic redox
flow batteries: Active materials, electrolytes and membranes. J. Energy
Chem. 2018, 27, 1304−1325.
(29) Park, M.; Ryu, J.; Wang, W.; Cho, J. Material design andengineering of next-generation flow-battery technologies. Nat. Rev.
Mater. 2017, 2, 16080.
(30) Hendriks, K. H.; Robinson, S. G.; Braten, M. N.; Sevov, C. S.;
Helms, B. A.; Sigman, M. S.; Minteer, S. D.; Sanford, M. S. High-He was a postdoctoral associate with Professor Roy G. Gordon andProfessor Michael J. Aziz at Harvard University from 2017 to 2019,
where his work focused on aqueous organic redox ﬂow batteries. He
received his faculty position at Zhejiang Sci-Tech University and began
his independent career in January 2020. His research interests include
the design, synthesis, and characterization of organic materials forenergy storage and conversion applications.
Michael J. Aziz received a B.S. from Caltech in 1978 and a Ph.D. in
Applied Physics from Harvard in 1983. He spent two years at Oak Ridge
National Laboratory as Eugene P. Wigner Postdoctoral Fellow. He has
been a member of the faculty at what is now the Harvard John A.
Paulson School of Engineering and Applied Sciences since he joined in1986 and is now Gene and Tracy Sykes Professor of Materials andEnergy Technologies. He has made signiﬁcant contributions to anumber of ﬁelds in applied physics and materials science including the
kinetics of rapid solidiﬁcation, pressure and stress eﬀects on kinetics ofdiﬀusion and growth, nonequilibrium surface pattern formation during
ion irradiation, and applications of materials synthesized utilizing
nonequilibrium kinetics. His recent interests involve energy storage andCO2 capture. He has directed a multi-investigator research program on
stationary electrical energy storage since 2012 and is coinventor of theaqueous organic redox ﬂow battery, with multiple patents that havebeen licensed for commercialization. He is corecipient of the 2019
Energy Frontiers Prize from Eni for these developments. He served as
the Faculty Coordinator for Harvard’s University-Wide Graduate
Consortium on Energy and Environment from 2009 to 2018, for which
he developed a quantitative course on Energy Technology for a groupof students in diverse disciplines, and is currently authoring a textbook
on energy technology.

ACKNOWLEDGMENTS
Research at Harvard has been supported by U.S. DOE Award
DE-AC05-76RL01830 through PNNL Subcontract 428977, by
U.S. DOE ARPA-E Award DE-AR-0000767, by Innovation
Fund Denmark via the Grand Solutions project “ORBATS” ﬁleNo. 7046-00018B, by the U.S. National Science Foundation
through Grant No. NSF CBET-1914543, and by the
Massachusetts Clean Energy Technology Center. We thank
Sanat Modak and Swayam Patnaik for assistance in compilingand plotting the results from the literature and Eric Fell and Dr.
Yan Jing for helpful discussions.
REFERENCES(1) IPCC Climate Change 2014: Synthesis Report; IPCC, 2014.
(2) Raupach, M. R.; Marland, G.; Ciais, P.; Le Quere, C.; Canadell, J.
G.; Klepper, G.; Field, C. B. Global and regional drivers of accelerating
CO2 emissions. Proc. Natl. Acad. Sci. U. S. A. 2007, 104, 10288−10293.
(3) Pacala, S.; Socolow, R. Stabilization Wedges: Solving the Climate
Problem for the Next 50 Years with Current Technologies. Science
2004, 305, 968−971.
(4) Ferrara, M.; Chiang, Y. M.; Deutch, J. M. Demonstrating NearCarbon-Free Electricity Generation from Renewables and Storage.
Joule 2019, 3, 2585−2588.
(5) Ziegler, M. S.; Mueller, J. M.; Pereira, G. D.; Song, J.; Ferrara, M.;
Chiang, Y.-M.; Trancik, J. E. Storage Requirements and Costs of
Shaping Renewable Energy Toward Grid Decarbonization. Joule 2019,3, 2134−2153.
(6) Olivetti, E. A.; Ceder, G.; Gaustad, G. G.; Fu, X. Lithium-Ion
Battery Supply Chain Considerations: Analysis of Potential Bottlenecks
in Critical Metals. Joule 2017, 1, 229−243.
(7) Lisbona, D.; Snee, T. A review of hazards associated with primarylithium and lithium-ion batteries. Process Saf. Environ. Prot. 2011, 89,
434−442.
Shttps://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Performance Oligomeric Catholytes for Effective Macromolecular
Separation in Nonaqueous Redox Flow Batteries. ACS Cent. Sci. 2018,
4, 189−196.
(31) Doris, S. E.; Ward, A. L.; Baskin, A.; Frischmann, P. D.;
Gavvalapalli, N.; Chénard, E.; Sevov, C. S.; Prendergast, D.; Moore, J.
S.; Helms, B. A. Macromolecular Design Strategies for Preventing
Active-Material Crossover in Non-Aqueous All-Organic Redox-Flow
Batteries. Angew. Chem., Int. Ed. 2017, 56, 1595−1599.
(32) Goulet, M. A.; Tong, L.; Pollack, D. A.; Tabor, D. P.; Odom, S.
A.; Aspuru-Guzik, A.; Kwan, E. E.; Gordon, R. G.; Aziz, M. J. Extending
the Lifetime of Organic Flow Batteries via Redox State Management. J.
Am. Chem. Soc. 2019, 141, 8014−8019.
(33) Goulet, M.-A.; Aziz, M. J. Flow Battery Molecular Reactant
Stability Determined by Symmetric Cell Cycling Methods. J. Electrochem. Soc. 2018, 165, A1466−A1477.
(34) Hollas, A.; Wei, X.; Murugesan, V.; Nie, Z.; Li, B.; Reed, D.; Liu,
J.; Sprenkle, V.; Wang, W. A biomimetic high-capacity phenazine-based
anolyte for aqueous organic redox flow batteries. Nat. Energy 2018, 3,
508−514.
(35) Sun, P.; Liu, Y.; Li, Y.; Shehzad, M. A.; Liu, Y.; Zuo, P.; Chen, Q.;
Yang, Z.; Xu, T. 110th Anniversary: Unleashing the Full Potential of
Quinones for High Performance Aqueous Organic Flow Battery. Ind.
Eng. Chem. Res. 2019, 58, 3994−3999.
(36) Yang, Z. J.; Tong, L. C.; Tabor, D. P.; Beh, E. S.; Goulet, M. A.;
De Porcellinis, D.; Aspuru-Guzik, A.; Gordon, R. G.; Aziz, M. J. Alkaline
Benzoquinone Aqueous Flow Battery for Large-Scale Storage ofElectrical Energy. Adv. Energy Mater. 2018, 8, 1702056.
(37) Darling, R.; Gallagher, K.; Xie, W.; Su, L.; Brushett, F. Transport
Property Requirements for Flow Battery Separators. J. Electrochem. Soc.
2016, 163, A5029−A5040.
(38) Darling, R. M.; Weber, A. Z.; Tucker, M. C.; Perry, M. L. The
Influence of Electric Field on Crossover in Redox-Flow Batteries. J.
Electrochem. Soc. 2016, 163, A5014−A5022.
(39) Ashraf Gandomi, Y.; Aaron, D. S.; Mench, M. M. Coupled
Membrane Transport Parameters for Ionic Species in All-Vanadium
Redox Flow Batteries. Electrochim. Acta 2016, 218, 174−190.
(40) Luo, Q.; Li, L.; Nie, Z.; Wang, W.; Wei, X.; Li, B.; Chen, B.; Yang,
Z. In-situ investigation of vanadium ion transport in redox flow battery.
J. Power Sources 2012, 218, 15−20.
(41) Chen, H.; Cong, G.; Lu, Y.-C. Recent progress in organic redox
flow batteries: Active materials, electrolytes and membranes. J. Energy
Chem. 2018, 27, 1304−1325.
(42) Wei, X. L.; Pan, W. X.; Duan, W. T.; Hollas, A.; Yang, Z.; Li, B.;
Nie, Z. M.; Liu, J.; Reed, D.; Wang, W.; Sprenkle, V. Materials andSystems for Organic Redox Flow Batteries: Status and Challenges. ACS
Energy Lett. 2017, 2, 2187−2204.
(43) Leung, P.; Shah, A. A.; Sanz, L.; Flox, C.; Morante, J. R.; Xu, Q.;
Mohamed, M. R.; Ponce de León, C.; Walsh, F. C. Recent
developments in organic redox flow batteries: A critical review. J.
Power Sources 2017, 360, 243−283.
(44) Leung, P.; Li, X.; Ponce de León, C.; Berlouis, L.; Low, C. T. J.;
Walsh, F. C. Progress in redox flow batteries, remaining challenges andtheir applications in energy storage. RSC Adv. 2012, 2, 10125.
(45) Chen, R. Toward High-Voltage, Energy-Dense, and Durable
Aqueous Organic Redox Flow Batteries: Role of the Supporting
Electrolytes. ChemElectroChem 2019, 6, 603−612.
(46) Singh, V.; Kim, S.; Kang, J.; Byon, H. R. Aqueous organic redox
flow batteries. Nano Res. 2019, 12, 1988−2001.
(47) Luo, J. A.; Hu, B.; Hu, M. W.; Zhao, Y.; Liu, T. L. Status andProspects of Organic Redox Flow Batteries toward Sustainable Energy
Storage. ACS Energy Lett. 2019, 4, 2220−2240.
(48) Er, S.; Suh, C.; Marshak, M. P.; Aspuru-Guzik, A. Computational
Design of Molecules for an All-Quinone Redox Flow Battery. Chem. Sci.
2015, 6, 885−893.
(49) Huynh, M. T.; Anson, C. W.; Cavell, A. C.; Stahl, S. S.; HammesSchiffer, S. Quinone 1 e- and 2 e-/2 H+ Reduction Potentials:
Identification and Analysis of Deviations from Systematic Scaling
Relationships. J. Am. Chem. Soc. 2016, 138, 15903−15910.

Review(50) Kim, H.; Goodson, T.; Zimmerman, P. M. Achieving Accurate
Reduction Potential Predictions for Anthraquinones in Water andAprotic Solvents: Effects of Inter- and Intramolecular H-Bonding and
Ion Pairing. J. Phys. Chem. C 2016, 120, 22235−22247.
(51) Lim, H.-D.; Lee, B.; Zheng, Y.; Hong, J.; Kim, J.; Gwon, H.; Ko,Y.; Lee, M.; Cho, K.; Kang, K. Rational design of redox mediators foradvanced Li-O2 batteries. Nat. Energy 2016, 1, 16066.
(52) Gerhardt, M. R.; Tong, L. C.; Gomez-Bombarelli, R.; Chen, Q.;
Marshak, M. P.; Galvin, C. J.; Aspuru-Guzik, A.; Gordon, R. G.; Aziz, M.
J. Anthraquinone Derivatives in Aqueous Flow Batteries. Adv. Energy
Mater. 2017, 7, 1601488.
(53) Kwabi, D. G.; Lin, K.; Ji, Y.; Kerr, E. F.; Goulet, M.-A.; De
Porcellinis, D.; Tabor, D. P.; Pollack, D. A.; Aspuru-Guzik, A.; Gordon,
R. G.; Aziz, M. J. Alkaline quinone flow battery with long lifetime at pH
12. Joule 2018, 2, 1907.
(54) Ji, Y.; Goulet, M.-A.; Pollack, D. A.; Kwabi, D. G.; Jin, S.;
Porcellinis, D.; Kerr, E. F.; Gordon, R. G.; Aziz, M. J. A PhosphonateFunctionalized Quinone Redox Flow Battery at Near-Neutral pH with
Record Capacity Retention Rate. Adv. Energy Mater. 2019, 9, 1900039.
(55) Han, K. S.; Rajput, N. N.; Vijayakumar, M.; Wei, X.; Wang, W.;
Hu, J.; Persson, K. A.; Mueller, K. T. Preferential Solvation of an
Asymmetric Redox Molecule. J. Phys. Chem. C 2016, 120, 27834−
27839.
(56) Luo, J.; Hu, B.; Debruler, C.; Bi, Y.; Zhao, Y.; Yuan, B.; Hu, M.;
Wu, W.; Liu, T. L. Unprecedented capacity and stability of ammonium
ferrocyanide catholyte in pH neutral aqueous redox flow batteries. Joule
2019, 3, 149−163.
(57) Hamlin, T. A.; Swart, M.; Bickelhaupt, F. M. Nucleophilic
Substitution (SN2): Dependence on Nucleophile, Leaving Group,
Central Atom, Substituents, and Solvent. ChemPhysChem 2018, 19,1315−1330.
(58) Smith, M. B.; March, J. March’s Advanced Organic Chemistry:
Reactions, Mechanisms, and Structure, 6th ed.; Wiley: Hoboken, NJ,
2007.
(59) Tabor, D. P.; Gómez-Bombarelli, R.; Tong, L.; Gordon, R. G.;
Aziz, M. J.; Aspuru-Guzik, A. Mapping the frontiers of quinone stability
in aqueous media: implications for organic aqueous redox flow
batteries. J. Mater. Chem. A 2019, 7, 12833−12841.
(60) Huskinson, B.; Marshak, M. P.; Suh, C.; Er, S.; Gerhardt, M. R.;
Galvin, C. J.; Chen, X.; Aspuru-Guzik, A.; Gordon, R. G.; Aziz, M. J. A
metal-free organic-inorganic aqueous flow battery. Nature 2014, 505,
195−198.
(61) Murali, A.; Nirmalchandar, A.; Krishnamoorthy, S.; HooberBurkhardt, L.; Yang, B.; Soloveichik, G.; Prakash, G. K. S.; Narayanan,
S. R. Understanding and mitigating capacity fade in aqueous organic
redox flow batteries. J. Electrochem. Soc. 2018, 165, A1193.
(62) Liu, T.; Wei, X.; Nie, Z.; Sprenkle, V.; Wang, W. A Total Organic
Aqueous Redox Flow Battery Employing a Low Cost and Sustainable
Methyl Viologen Anolyte and 4-HO-TEMPO Catholyte. Adv. Energy
Mater. 2016, 6, 1501449.
(63) Orita, A.; Verde, M. G.; Sakai, M.; Meng, Y. S. The impact of pH
on side reactions for aqueous redox flow batteries based on nitroxyl
radical compounds. J. Power Sources 2016, 321, 126−134.
(64) Luo, J.; Sam, A.; Hu, B.; DeBruler, C.; Wei, X.; Wang, W.; Liu, T.
L. Unraveling pH Dependent Cycling Stability of Ferricyanide/
Ferrocyanide in Redox Flow Batteries. Nano Energy 2017, 42, 215−221.
(65) Kumpan, N.; Poonsawat, T.; Chaicharoenwimolkul, L.;
Pornsuwan, S.; Somsook, E. Ferrocenated nanocatalysts derived fromthe decomposition of ferrocenium in basic solution and their aerobic
activities for the rapid decolorization of methylene blue and the facile
oxidation of phenylboronic acid. RSC Adv. 2017, 7, 5759−5763.
(66) Tabbì, G.; Cassino, C.; Cavigiolio, G.; Colangelo, D.; Ghiglia, A.;
Viano, I.; Osella, D. Water Stability and Cytotoxic Activity Relationshipof a Series of Ferrocenium Derivatives. ESR Insights on the Radical
Production during the Degradation Process. J. Med. Chem. 2002, 45,5786−5796.
(67) Thomas, J.-L.; Howarth, J.; Hanlon, K.; McGuirk, D. Ferrocenyl
imidazolium salts as a new class of anion receptors with C-H&Xhydrogen bonding. Tetrahedron Lett. 2000, 41, 413−416.
Thttps://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR(68) Gasser, G.; Fischmann, A. J.; Forsyth, C. M.; Spiccia, L. Productsof hydrolysis of (ferrocenylmethyl)trimethylammonium iodide: Synthesis of hydroxymethylferrocene and bis(ferrocenylmethyl) ether. J.
Organomet. Chem. 2007, 692, 3835−3840.
(69) Bird, C. L.; Kuhn, A. T. Electrochemistry of the viologens. Chem.
Soc. Rev. 1981, 10, 49−82.
(70) Murugavel, K. Benzylic viologen dendrimers: a review of theirsynthesis, properties and applications. Polym. Chem. 2014, 5, 5873−
5884.
(71) Bard, A. J.; Ledwith, A.; Shine, H. J. Formation, Properties andReactions of Cation Radicals in Solution. Adv. Phys. Org. Chem. 1976,
13, 155−278.
(72) Venturi, M.; Mulazzani, Q. G.; Hoffman, M. Z. Radiolyticallyinduced one-electron reduction of methyl viologen in aqueous solution.
Radiat. Phys. Chem. 1984, 23, 229−236.
(73) O’Connor, C. Acidic and basic amide hydrolysis. Q. Rev., Chem.
Soc. 1970, 24, 553−564.
(74) Koziol, J.; Tyrakowska, B.; Muller, F. The Structure of Covalent
Hydrates) of Alloxazines - a Reinvestigation. Helv. Chim. Acta 1981, 64,
1812−1817.
(75) Koziol, J.; Metzler, D. E. Formation and Possible Structure of
Covalent Hydrates of Alloxazines. Z. Naturforsch., B: J. Chem. Sci. 1972,
27, 1027−1029.
(76) Smith, S. B.; Bruice, T. C. Mechanisms of isoalloxazine (flavine)hydrolysis. J. Am. Chem. Soc. 1975, 97, 2875−2881.
(77) Surrey, A. R.; Nachod, F. C. Alkaline Hydrolysis of Riboflavin. J.
Am. Chem. Soc. 1951, 73, 2336−2338.
(78) Harayama, T.; Tezuka, Y.; Taga, T.; Yoneda, F. Hydrolysis
Products of Flavins (Isoalloxazines). J. Chem. Soc., Perkin Trans. 1 1987,
75−83.
(79) Lin, K.; Gómez-Bombarelli, R.; Beh, E. S.; Tong, L.; Chen, Q.;
Valle, A.; Aspuru-Guzik, A.; Aziz, M. J.; Gordon, R. G. A redox-flowbattery with an alloxazine-based organic electrolyte. Nat. Energy 2016,
1, 16102.
(80) Beck, F.; Heydecke, G. On the Mechanism of the Cathodic
Reduction of Anthraquinone to Anthrone. Ber. Bunsenges. Phys. Chem.
1987, 91, 37−43.
(81) Shi, Z.-W.; Li, Y.-Z.; Li, Y.; Lu, G.-Y.; Liu, S.-H. 1,1′,8,8′Tetramethoxy-10,10′-bianthrone. Acta Crystallogr., Sect. E: Struct. Rep.
Online 2004, 60, No. o2275.
(82) Jin, S.; Jing, Y.; Kwabi, D. G.; Ji, Y.; Tong, L.; De Porcellinis, D.;
Goulet, M.-A.; Pollack, D. A.; Gordon, R. G.; Aziz, M. J. A WaterMiscible Quinone Flow Battery with High Volumetric Capacity andEnergy Density. ACS Energy Lett. 2019, 4, 1342−1348.
(83) Antonov, L. Tautomerism: Methods and Theories; Wiley-VCH:
Weinheim, Germany, 2013.
(84) Deuchert, K.; Hünig, S. Multistage Organic Redox SystemsA
General Structural Principle. Angew. Chem., Int. Ed. Engl. 1978, 17,875−886.
(85) Thomson, R. H. Phenol tautomerism. Q. Rev., Chem. Soc. 1956,
10, 27−43.
(86) Thomson, R. H. 352. The structure of β-hydrojuglone andrelated compounds. Keto-;enols of the naphthalene series. J. Chem. Soc.
1950, 0, 1737−1742.
(87) Saveant, J.-M. Elements of Molecular and Biomolecular Electrochemistry; Wiley: Hoboken, NJ, 2006.
(88) Chen, Q.; Eisenach, L.; Aziz, M. J. Cycling Analysis of a QuinoneBromide Redox Flow Battery. J. Electrochem. Soc. 2016, 163, A5057−
A5063.
(89) Chen, Q.; Gerhardt, M. R.; Aziz, M. J. Dissection of the voltage
losses of an acidic quinone redox flow battery. J. Electrochem. Soc. 2017,
164, A1126−A1132.
(90) Beh, E. S.; De Porcellinis, D.; Gracia, R. L.; Xia, K. T.; Gordon, R.
G.; Aziz, M. J. A neutral pH aqueous organic-organometallic redox flow
battery with extremely high capacity retention. ACS Energy Lett. 2017,
2, 639−644.
(91) Xu, Y.; Wen, Y. H.; Cheng, J.; Cao, G. P.; Yang, Y. S. A study oftiron in aqueous solutions for redox flow battery application.
Electrochim. Acta 2010, 55, 715−720.

Review(92) Xu, Y.; Wen, Y.; Cheng, J.; Cao, G.; Yang, Y. Study on a single
flow acid Cd-chloranil battery. Electrochem. Commun. 2009, 11, 1422−
1424.
(93) Nawar, S.; Huskinson, B.; Aziz, M. Benzoquinone-Hydroquinone Couple for Flow Battery. MRS Proceedings 2013, 1491, 1491.
(94) Huskinson, B.; Nawar, S.; Gerhardt, M. R.; Aziz, M. J. Novel
Quinone-Based Couples for Flow Batteries. ECS Trans. 2013, 53, 101−105.
(95) Yang, B.; Hoober-Burkhardt, L.; Krishnamoorthy, S.; Murali, A.;
Prakash, G. K. S.; Narayanan, S. R. High-Performance Aqueous Organic
Flow Battery with Quinone-Based Redox Couples at Both Electrodes. J.
Electrochem. Soc. 2016, 163, A1442−A1449.
(96) Hoober-Burkhardt, L.; Krishnamoorthy, S.; Yang, B.; Murali, A.;
Nirmalchandar, A.; Prakash, G. K. S.; Narayanan, S. R. A new Michaelreaction-resistant benzoquinone for aqueous organic redox flow
batteries. J. Electrochem. Soc. 2017, 164, A600−A607.
(97) Yang, B.; Prakash, G. K. S.; Aniszfeld, R.; Narayan, S. R.; HooberBurkhardt, L.; Krishnamoorthy, S.; Murali, A.; Nirmalchandar, A.
Stable Positive Side Material for All-Organic Flow Battery. U.S. Patent
Application 2019/0115594 A1, 2019.
(98) Tong, L.; Goulet, M.-A.; Tabor, D. P.; Kerr, E. F.; De Porcellinis,
D.; Fell, E. M.; Aspuru-Guzik, A.; Gordon, R. G.; Aziz, M. J. MolecularEngineering of an Alkaline Naphthoquinone Flow Battery. ACS Energy
Lett. 2019, 4, 1880−1887.
(99) Jin, S.; Jing, Y.; Kwabi, D. G.; Ji, Y.; Tong, L.; De Porcellinis, D.;
Goulet, M.-A.; Pollack, D. A.; Gordon, R. G.; Aziz, M. J. A WaterMiscible Quinone Flow Battery with High Volumetric Capacity andEnergy Density. ACS Energy Letters 2019, 4, 1342−1348.
(100) Wang, C.; Yang, Z.; Wang, Y.; Zhao, P.; Yan, W.; Zhu, G.; Ma,
L.; Yu, B.; Wang, L.; Li, G.; Liu, J.; Jin, Z. High-performance alkalineorganic redox flow batteries based on 2-hydroxy-3-carboxy-1,4naphthoquinone. ACS Energy Lett. 2018, 3, 2404−2409.
(101) Yang, B.; Hoober-Burkhardt, L.; Wang, F.; Surya Prakash, G. K.;
Narayanan, S. R. An inexpensive aqueous flow battery for large-scale
electrical energy storage based on water-soluble organic redox couples.
J. Electrochem. Soc. 2014, 161, A1371−A1380.
(102) Zhang, S.; Li, X.; Chu, D. An Organic Electroactive Material for
Flow Batteries. Electrochim. Acta 2016, 190, 737−743.
(103) Park, M.; Beh, E. S.; Fell, E. M.; Jing, Y.; Kerr, E. F.; Porcellinis,
D.; Goulet, M.-A.; Ryu, J.; Wong, A. A.; Gordon, R. G.; Cho, J.; Aziz, M.
J. A high voltage aqueous zinc-organic hybrid flow battery. Adv. Energy
Mater. 2019, 9, 1900694.
(104) Wedege, K.; Dražević, E.; Konya, D.; Bentien, A. Organic Redox
Species in Aqueous Flow Batteries: Redox Potentials, Chemical
Stability and Solubility. Sci. Rep. 2016, 6, 39101.
(105) Leung, P. K.; Martin, T.; Shah, A. A.; Anderson, M. A.; Palma, J.
Membrane-less organic-inorganic aqueous flow batteries with improved cell potential. Chem. Commun. 2016, 52, 14270−14273.
(106) Xu, Y.; Zheng, Y.; Wang, C.; Chen, Q. An All-Organic Aqueous
Battery Powered by Adsorbed Quinone. ACS Appl. Mater. Interfaces
2019, 11, 23222−23228.
(107) Preger, Y.; Gerken, J. B.; Biswas, S.; Anson, C. W.; Johnson, M.
R.; Root, T. W.; Stahl, S. S. Quinone-Mediated Electrochemical O2
Reduction Accessing High Power Density with an Off-Electrode CoN/C Catalyst. Joule 2018, 2, 2722−2731.
(108) Carney, T. J.; Collins, S. J.; Moore, J. S.; Brushett, F. R.
Concentration-Dependent Dimerization of Anthraquinone Disulfonic
Acid and Its Impact on Charge Storage. Chem. Mater. 2017, 29, 4801−
4810.
(109) Wiberg, C.; Carney, T. J.; Brushett, F.; Ahlberg, E.; Wang, E.
Dimerization of 9,10-anthraquinone-2,7-Disulfonic acid (AQDS).
Electrochim. Acta 2019, 317, 478−485.
(110) Gerhardt, M. R.; Beh, E. S.; Tong, L.; Gordon, R. G.; Aziz, M. J.
Comparison of Capacity Retention Rates During Cycling of QuinoneBromide Flow Batteries. MRS Adv. 2017, 2, 431−438.
(111) Wedege, K.; Azevedo, J.; Khataee, A.; Bentien, A.; Mendes, A.
Direct Solar Charging of an Organic-Inorganic, Stable, and Aqueous
Alkaline Redox Flow Battery with a Hematite Photoanode. Angew.
Chem., Int. Ed. 2016, 55, 7142−7147.
U

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR(112) Li, W.; Fu, H. C.; Li, L.; Caban-Acevedo, M.; He, J. H.; Jin, S.
Integrated Photoelectrochemical Solar Energy Conversion and Organic
Redox Flow Battery Devices. Angew. Chem., Int. Ed. 2016, 55, 13104−
13108.
(113) McKone, J. R.; DiSalvo, F. J.; Abruna, H. D. Solar energy
conversion, storage, and release using an integrated solar-driven redox
flow battery. J. Mater. Chem. A 2017, 5, 5362−5372.
(114) Kwabi, D. G.; Wong, A. A.; Aziz, M. J. Rational Evaluation and
Cycle Life Improvement of Quinone-Based Aqueous Flow Batteries
Guided by In-Line Optical Spectrophotometry. J. Electrochem. Soc.
2018, 165, A1770−A1776.
(115) Xie, W.; Darling, R. M.; Perry, M. L. Processing andPretreatment Effects on Vanadium Transport in Nafion Membranes.
J. Electrochem. Soc. 2016, 163, A5084−A5089.
(116) Smith, M. B. March’s Advanced Organic Chemistry: Reactions,
Mechanisms, and Structure, 5th ed.; Wiley-VCH: Weinheim, Germany,
2001.
(117) Chang, Z.; Henkensmeier, D.; Chen, R. One-Step CationicGrafting of 4-Hydroxy-TEMPO and its Application in a Hybrid Redox
Flow Battery with a Crosslinked PBI Membrane. ChemSusChem 2017,10, 3193−3197.
(118) Liu, Y.; Goulet, M.-A.; Tong, L.; Liu, Y.; Ji, Y.; Wu, L.; Gordon,
R. G.; Aziz, M. J.; Yang, Z.; Xu, T. A Long-Lifetime All-Organic
Aqueous Flow Battery Utilizing TMAP-TEMPO Radical. Chem. 2019,5, 1861−1870.
(119) Winsberg, J.; Stolze, C.; Muench, S.; Liedl, F.; Hager, M. D.;
Schubert, U. S. TEMPO/phenazine combi-molecule: a redox-active
material for symmetric aqueous redox-flow batteries. ACS Energy Lett.
2016, 1, 976−980.
(120) Janoschka, T.; Martin, N.; Martin, U.; Friebe, C.; Morgenstern,S.; Hiller, H.; Hager, M. D.; Schubert, U. S. An aqueous, polymer-based
redox-flow battery using non-corrosive, safe, and low-cost materials.
Nature 2015, 527, 78−81.
(121) Janoschka, T.; Martin, N.; Hager, M. D.; Schubert, U. S. An
aqueous redox-flow battery with high capacity and power: the
TEMPTMA/MV system. Angew. Chem., Int. Ed. 2016, 55, 14427−
14430.
(122) Luo, J.; Hu, B.; DeBruler, C.; Liu, T. A “π-Conjugation
Extended Viologen” as Novel Two-Electron Storage Anolyte for Total
Organic Aqueous Redox Flow Battery. Angew. Chem., Int. Ed. 2018, 57,
231−235.
(123) Chang, Z.; Henkensmeier, D.; Chen, R. Shifting redox potentialof nitroxyl radical by introducing an imidazolium substituent and its use
in aqueous flow batteries. J. Power Sources 2019, 418, 11−16.
(124) Winsberg, J.; Stolze, C.; Schwenke, A.; Muench, S.; Hager, M.
D.; Schubert, U. S. Aqueous 2,2,6,6-Tetramethylpiperidine-N-oxyl
Catholytes for a High-Capacity and High Current Density OxygenInsensitive Hybrid-Flow Battery. ACS Energy Lett. 2017, 2, 411−416.
(125) Hu, B.; DeBruler, C.; Rhodes, Z.; Liu, T. L. Long-cycling
aqueous organic redox flow battery (AORFB) toward sustainable andsafe energy storage. J. Am. Chem. Soc. 2017, 139, 1207−1214.
(126) Gong, K.; Xu, F.; Grunewald, J. B.; Ma, X.; Zhao, Y.; Gu, S.; Yan,
Y. All-Soluble All-Iron Aqueous Redox-Flow Battery. ACS Energy Lett.
2016, 1, 89−93.
(127) DeBruler, C.; Hu, B.; Moss, J.; Luo, J.; Liu, T. L. A sulfonatefunctionalized viologen enabling neutral cation exchange, aqueousorganic redox flow batteries toward renewable energy storage. ACS
Energy Lett. 2018, 3, 663−668.
(128) Huang, J. H.; Yang, Z.; Murugesan, V.; Walter, E.; Hollas, A.;
Pan, B. F.; Assary, R. S.; Shkrob, I. A.; Wei, X. L.; Zhang, Z. C. SpatiallyConstrained Organic Diquat Anolyte for Stable Aqueous Flow
Batteries. ACS Energy Lett. 2018, 3, 2533−2538.
(129) DeBruler, C.; Hu, B.; Moss, J.; Liu, X. A.; Luo, J. A.; Sun, Y. J.;
Liu, T. L. Designer Two-Electron Storage Viologen Anolyte Materials
for Neutral Aqueous Organic Redox Flow Batteries. Chem. 2017, 3,
961−978.
(130) Hu, B.; Tang, Y.; Luo, J.; Grove, G.; Guo, Y.; Liu, T. Improved
Radical Stability of Viologen Anolyte in Aqueous Organic Redox Flow
Battery. Chem. Commun. 2018, 54, 6871−6874.

Review(131) Varcoe, J. R.; Atanassov, P.; Dekel, D. R.; Herring, A. M.;
Hickner, M. A.; Kohl, P. A.; Kucernak, A. R.; Mustain, W. E.; Nijmeijer,
K.; Scott, K.; Xu, T.; Zhuang, L. Anion-exchange membranes inelectrochemical energy systems. Energy Environ. Sci. 2014, 7, 3135−
3191.
(132) Zhu, Y.; Yang, F.; Niu, Z.; Wu, H.; He, Y.; Zhu, H.; Ye, J.; Zhao,
Y.; Zhang, X. Enhanced cyclability of organic redox flow batteries
enabled by an artificial bipolar molecule in neutral aqueous electrolyte.
J. Power Sources 2019, 417, 83−89.
(133) Yu, X. W.; Manthiram, A. A strategically managed rechargeable
battery system with a neutral methyl viologen anolyte and an acidic aircathode enabled by a mediator-ion solid electrolyte. Sustainable Energy& Fuels 2018, 2, 1452−1457.
(134) Winsberg, J.; Hagemann, T.; Muench, S.; Friebe, C.; Häupler,B.; Janoschka, T.; Morgenstern, S.; Hager, M. D.; Schubert, U. S.
Poly(boron-dipyrromethene)A Redox-Active Polymer Class for
Polymer Redox-Flow Batteries. Chem. Mater. 2016, 28, 3401−3405.
(135) Winsberg, J.; Janoschka, T.; Morgenstern, S.; Hagemann, T.;
Muench, S.; Hauffman, G.; Gohy, J. F.; Hager, M. D.; Schubert, U. S.
Poly(TEMPO)/Zinc Hybrid-Flow Battery: A Novel, “Green,” High
Voltage, and Safe Energy Storage System. Adv. Mater. 2016, 28, 2238−
2243.
(136) Winsberg, J.; Muench, S.; Hagemann, T.; Morgenstern, S.;
Janoschka, T.; Billing, M.; Schacher, F. H.; Hauffman, G.; Gohy, J.-F.;
Hoeppener, S.; Hager, M. D.; Schubert, U. S. Polymer/zinc hybrid-flowbattery using block copolymer micelles featuring a TEMPO corona as
catholyte. Polym. Chem. 2016, 7, 1711−1718.
(137) Häupler, B.; Rössel, C.; Schwenke, A. M.; Winsberg, J.; Schmidt,
D.; Wild, A.; Schubert, U. S. Aqueous zinc-organic polymer battery witha high rate performance and long lifetime. NPG Asia Mater. 2016, 8,
No. e283.
(138) Milshtein, J. D.; Su, L.; Liou, C.; Badel, A. F.; Brushett, F. R.
Voltammetry study of quinoxaline in aqueous electrolytes. Electrochim.
Acta 2015, 180, 695−704.
(139) Orita, A.; Verde, M. G.; Sakai, M.; Meng, Y. S. A biomimetic
redox flow battery based on flavin mononucleotide. Nat. Commun.
2016, 7, 13230.
(140) Sevov, C. S.; Hendriks, K. H.; Sanford, M. S. Low-Potential
Pyridinium Anolyte for Aqueous Redox Flow Batteries. J. Phys. Chem. C
2017, 121, 24376−24380.
(141) Huang, Z. F.; Zhang, P.; Gao, X. P.; Henkensmeier, D.;
Passerini, S.; Chen, R. Y. Unlocking Simultaneously the Temperatureand Electrochemical Windows of Aqueous Phthalocyanine Electrolytes.
Acs Applied Energy Materials 2019, 2, 3773−3779.
(142) Leung, P.; Aili, D.; Xu, Q.; Rodchanarowan, A.; Shah, A. A.
Rechargeable organic-air redox flow batteries. Sustain. Energy Fuels
2018, 2, 2252−2259.
(143) Hong, J.; Kim, K. Neutral Red and Ferroin as Reversible and
Rapid Redox Materials for Redox Flow Batteries. ChemSusChem 2018,11, 1866−1872.
(144) Brushett, F.; Jansen, A. N.; Vaughey, J. T.; Su, L.; Milshtein, J. D.
Materials for use with aqueous redox ﬂow batteries and related methods
and systems. U.S. Patent 14,625,417, 2015.
(145) Zhang, C.; Niu, Z.; Peng, S.; Ding, Y.; Zhang, L.; Guo, X.; Zhao,Y.; Yu, G. Phenothiazine-Based Organic Catholyte for High-Capacity
and Long-Life Aqueous Redox Flow Batteries. Adv. Mater. 2019, 31,
1901052.
(146) McCormick, D. B.; Fory, W. Base-Catalyzed Hydrolysis of
Flavins - Effect of Amines and Particular Role of Position 3. J. Pharm.
Sci. 1968, 57, 841−844.
(147) Grajek, H.; Zurkowska, G.; Drabent, R.; Bojarski, C. The
structure of the flavomononucleotide dimer. Biochim. Biophys. Acta,
Gen. Subj. 1986, 881, 241−247.
(148) Darling, R. M.; Perry, M. L. Half-Cell, Steady-State FlowBattery Experiments. ECS Trans. 2013, 53, 31−38.
(149) Milshtein, J. D.; Barton, J. L.; Darling, R. M.; Brushett, F. R. 4acetamido-2,2,6,6-tetramethylpiperidine-1-oxyl as a model organic
redox active compound for nonaqueous flow batteries. J. Power Sources
2016, 327, 151−159.
V

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

pubs.acs.org/CR

Review(150) Milshtein, J. D.; Kaur, A. P.; Casselman, M. D.; Kowalski, J. A.;
Modekrutti, S.; Zhang, P. L.; Harsha Attanayake, N.; Elliott, C. F.;
Parkin, S. R.; Risko, C.; Brushett, F. R.; Odom, S. A. High current
density, long duration cycling of soluble organic active species for nonaqueous redox flow batteries. Energy Environ. Sci. 2016, 9, 3531−3543.
(151) Drazevic, E.; Szabo, C.; Konya, D.; Lund, T.; Wedege, K.;
Bentien, A. Investigation of Tetramorpholinohydroquinone as a
Potential Catholyte in a Flow Battery. ACS Appl. Energy Mater. 2019,2, 4745−4754.
(152) Hu, B.; Luo, J.; Hu, M.; Yuan, B.; Liu, T. L. A pH-Neutral,
Metal-Free Aqueous Organic Redox Flow Battery Employing an
Ammonium Anthraquinone Anolyte. Angew. Chem., Int. Ed. 2019, 58,
16629−16636.
(153) Huskinson, B.; Marshak, M. P.; Gerhardt, M. R.; Aziz, M. J.
Cycling of a Quinone-Bromide Flow Battery for Large-Scale Electrochemical Energy Storage. ECS Trans. 2014, 61, 27−30.
(154) Asperger, S. Kinetics of the decomposition of potassium
ferrocyanide in ultra-violet light. Trans. Faraday Soc. 1952, 48, 617−
624.
(155) Arellano, C. A. P.; Martínez, S. S. Effects of pH on the
degradation of aqueous ferricyanide by photolysis and photocatalysis
under solar radiation. Sol. Energy Mater. Sol. Cells 2010, 94, 327−332.
(156) Esswein, A. J.; Goeltz, J.; Amadeo, D. High Solubility Iron
Hexacyanides. U.S. Patent 13,887,461, 2014.

W

https://dx.doi.org/10.1021/acs.chemrev.9b00599
Chem. Rev. XXXX, XXX, XXX−XXX