A1466

Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)

Flow Battery Molecular Reactant Stability Determined by
Symmetric Cell Cycling Methods
Marc-Antoni Goulet

and Michael J. Aziz∗,z

Harvard John A. Paulson School of Engineering and Applied Sciences, Cambridge, Massachusetts 02138, USA
We present an unbalanced compositionally-symmetric flow cell method for revealing and quantifying different mechanisms forcapacity fade in redox flow batteries that are based on molecular energy storage. We utilize it, accompanied in some cases bya corresponding static-cell cycling method, to study capacity fade in cells comprising anthraquinone di-sulfonate, di-hydroxy
anthraquinone, iron hexacyanide, methyl viologen, and bis-trimethylammoniopropyl viologen. In all cases the cycling capacity
decay is reasonably consistent with exponential in time and is independent of the number of charge-discharge cycles imposed. By
introducing pauses at various states of charge of the capacity-limiting side during cycling, we show that in some cases the temporal fadetime constant is dependent on the state of charge. These observations suggest that molecular lifetime is dominated by chemical rather
than electrochemical mechanisms. These mechanisms include irrecoverable chemical decomposition and recoverable interactions
with cell materials. We conclude with recommendations for cell cycling protocols for evaluating stability of single electrolytes.
© The Author(s) 2018. Published by ECS. This is an open access article distributed under the terms of the Creative Commons
Attribution Non-Commercial No Derivatives 4.0 License (CC BY-NC-ND, http://creativecommons.org/licenses/by-nc-nd/4.0/),
which permits non-commercial reuse, distribution, and reproduction in any medium, provided the original work is not changed in any
way and is properly cited. For permission for commercial reuse, please email: oa@electrochem.org. [DOI: 10.1149/2.0891807jes]

Manuscript submitted March 26, 2018; revised manuscript received May 1, 2018. Published May 15, 2018.

Growing interest in redox flow batteries (RFBs) for grid-scale energy storage has prompted development of new redox-active organicand organometallic molecules synthesized from commonly available
raw materials containing earth-abundant elements such as carbon,
nitrogen and oxygen. Flow batteries based upon these synthetic compounds might achieve lower electrolyte capital costs than current
vanadium-based technologies without the scarcity risks associatedwith finite resources.1 To become a cost-effective solution for matching of intermittent renewable energy supply to demand, these flow
batteries should retain most of their capacity over decadal time scales.
As opposed to the elements of the periodic table, e.g. transition metals,
traditionally used for the redox-active components of flow batteries,
these new organic and organometallic compounds may be susceptible
to molecular decomposition, leading to an additional mechanism forflow battery charge storage capacity fade.2 Typically, capacity fadein conventional flow batteries can occur through various mechanisms
such as precipitation of reactant species or crossover of these species
through the membrane.3,4 Some of these sources of capacity loss
are theoretically reversible while others are not. Some are related todevice engineering, such as electrolyte leakage, whereas others are
intrinsic to the electrolyte chemistry. RFBs that rely upon electrolytes
with different dissolved redox-active species on either side, such as
the chromium/iron cell, suffer from irreversible crossover of these
species unless their electrolytes are balanced by putting both species
on both sides, thereby doubling the cost of active materials.5 a The
all-vanadium RFB was developed to avoid these problems by having
a single common redox-active species on both sides of the cell suchthat any asymmetric crossover of active species could be reversed byrebalancing the electrolytes by simple remixing.6 Other examples ofcapacity fade that can be recovered by electrolyte rebalancing methods
include those due to heterogeneous side reactions, such as hydrogen
evolution, or from homogeneous side reactions, such as species oxidation from surrounding air. Both of these side reactions can eventuallylimit the accessible state of charge (SOC) of one side, thereby limiting
the capacity of the battery. For redox flow batteries with dissimilar
active species in their two electrolytes, remixing is not an option;
hence these systems require alternative rebalancing methods such as
passing one of the electrolytes through a separate electrolysis cell tocorrect the SOC drift and recover lost capacity.7,8 Another form of
capacity fade that may affect all types of RFBs is the precipitation of∗ Electrochemical Society Member.
zE-mail: maziz@harvard.eduaThis is only feasible if the species do not interact chemically with each other and if all
species are stable at the potentials experienced on both sides of the cell.

dissolved species.b This is a common problem for the vanadium flow
battery due to the precipitation of vanadium pentoxide in the posolyte(positive electrolyte) at moderately high temperatures.9
All the capacity fade mechanisms mentioned above can also occur
in RFBs with molecular active species. In addition to those mechanisms, however, these compounds can also decompose or convert
into redox-inactive molecules, leading to irreversible capacity loss
over time. To aid the development of improved compounds for flow
batteries it is critical to refine experimental methods to differentiate
between these capacity fade mechanisms and adopt metrics that makeit possible to adequately compare electrolytes. The vast majority offlow battery studies published to date report capacity fade rates inpercent per cycle; this rate has been used to suggest the expected
lifetime of a hypothetical system cycled daily to provide grid scalestorage. This metric may be appropriate for solid state deposition orintercalation battery systems because their capacity fade can be related to cycling-related mechanisms such as dendrite formation ormechanical electrode damage, respectively.10–12 In contrast, if loss offlow battery capacity is primarily due to decomposition of chemicallyunstable redox active species, then a temporal metric is more appropriate such as the capacity fade per day, introduced by Beh et al. in2017.13 Moreover, lifetime comparisons between electrolytes are notwell captured by a cycle-denominated metric unless several factors
other than cycling regimen, such as concentration, reservoir size, flow
velocity and membrane area, are standardized between the electrolyte
chemistries being evaluated.c
For flow batteries to reach the decadal lifetimes that are ideal for
grid storage, it is critical to separate and quantify the contributionsof each source of capacity fade. To this end, methods for controllingcertain contributions while measuring others are essential. As a phenomenon common to all redox flow batteries with unlike dissolved
redox-active species on the two sides, crossover can be reduced oreliminated as a source of fade by symmetric cell testing.14 By using
two separate reservoirs with identical redox active species concentrations, Milshtein et al. further demonstrated how this method could be
used to study the capacity retention of a single electrolyte and pointed
out the advantages of the method over bulk electrolysis such as improved mass transfer and elimination of cross contamination by otherb
Although it may be theoretically possible to filter out and re-dissolve the precipitate,some of it may remain trapped in the porous electrodes leading to poor mass transportand cell failure.
cMore appropriately, capacity fade per cycle has been heavily utilized for evaluating membranes for imperviousness to reactant crossover: by keeping other parameters constant andstudying only one electrolyte combination, capacity fade per cycle can be a meaningful
way of comparing membranes for this purpose.

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Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)

A1467

species.15,16 In the present study, we further develop the method withexperimental variations that provide more specific information about
possible causes of capacity fade.
We utilize reservoir volume differences to impose a capacity limiting side (CLS) and a non-capacity-limiting side (NCLS), such thatthe instantaneous fade rate can be uniquely attributed to the state of
the CLS. We call this the unbalanced compositionally-symmetric cellmethod to reflect the fact that the volumes are unbalanced whereas
the compositions are symmetric. Validation of the method is performed using a simple and robust case for which capacity fade is
not expected, namely Fe2+ /Fe3+ . The method is then expanded toinclude cycling pauses at various SOC of the CLS to separately determine the stability of oxidized and reduced forms of the redoxactive species. These pauses suggest that electrolytes containing either 2,7-anthraquinone disulfonic acid (AQDS) or 2,6-dihydroxy anthraquinone (DHAQ) lose capacity more rapidly while the molecules
are held in reduced form than in oxidized form. In addition, we utilize pauses in electrolyte flow to demonstrate the interaction of theelectrolyte with cell materials in the case of ferricyanide at highpH. Furthermore, we contrast two viologen-based examples to illustrate and quantify the benefits of molecular design for protection
against decomposition mechanisms. Finally, the pros and cons ofsymmetric cell methods are discussed, and standard protocols are suggested for careful characterization of loss rates of novel flow battery
chemistries.
Methodology
Cell assembly.—All unbalanced compositionally-symmetric flow
cell cycling tests were performed with a 5 cm2 cell (Fuel Cell Tech,
Albuquerque, NM) equipped with POCO sealed graphite flow plateswith serpentine flow fields. Each side of the cell contained a stackof four sheets of Sigracet SGL 39AA porous carbon paper that had
been pretreated by baking in air at 400◦ C for 24 h. In most cases thecell was assembled with a Nafion N117 membrane pretreated by immersion in a blank electrolyte (i.e. missing the redox-active species)
for at least 72 hours. Cells for alkaline electrolytes such as DHAQ or
Fe(CN)6 contained Nafion membranes soaked in 1 M KOH, whereas
those prepared for acid electrolytes contained membranes soaked in
1 M H2 SO4 . In contrast, cells prepared for neutral pH viologen-based
electrolytes contained Selemion DSV (Asahi Glass Co., Ltd.) anion
exchange membranes which had been soaked in 1 M NaCl. In allcases, Viton (PVDF) gaskets were used to seal the electrodes between the membranes and graphite flow fields. Flow of electrolytes at
100 mL/min. was forced with a Cole-Parmer Masterflex L/S peristaltic
pump, which required a small length of Viton peristaltic tubing. Allother tubing and electrolyte reservoirs were made from chemically resistant fluorinated ethylene propylene (FEP). An example of this setup
is shown in Fig. 1a. Static cells were also developed for unbalanced
compositionally-symmetric cell tests and the details for these can befound in the SI.
Electrolyte preparation.—All electrolytes were prepared inroughly 0.1 M concentrations to make comparisons as relevant to battery operation as possible while staying within the limits of solubilityof all species. In most cases, 50% SOC electrolytes were preparedby charging a portion of initially-uncharged electrolyte in a full cell(against ferrocyanide in the case of DHAQ, against ferrocene in thecase of viologen, and against hydrogen in the case of AQDS) until the theoretical capacity was attained or the current decayed to asteady state value typically under 5 mA, and then mixing this charged
electrolyte in equal proportions with uncharged material. In the caseof Fe2+ /Fe3+ , a 0.1 M 50% SOC solution was prepared directly byadding FeCl2 and FeCl3 to 1 M H2 SO4 . Similarly, K3 [Fe(CN)6 ] andK4 [Fe(CN)6 ] were mixed directly into 1 M KOH to make a 0.1 M
50% SOC [Fe(CN)6 ]4− /[Fe(CN)6 ]3− electrolyte.
Cell cycling protocol.—Potentiostatic cycling of all cells was
performed in a nitrogen-filled glove box with < 2 PPM of oxy-

Figure 1. a) Schematic of symmetric flow cell. b) Example of potentiostatic
symmetric cell cycling protocol.

gen with a Biologic VSP 300 potentiostat. Unless otherwise specified, the potential limits were set to ± 200 mV to access over 99%
of the CLS capacity. An example of this cell cycling protocol isshown in Fig. 1b. In all two-electrode symmetric cells, one reservoir was filled with an excess of 50% SOC electrolyte to make theother side of the cell capacity limiting. For high potential posolyte
species such as Fe2+ /Fe3+ and [Fe(CN)6 ]4− /[Fe(CN)6 ]3− , this excess
was only marginal (10-20% more volume) to mitigate any possible
effect of electrolyte asymmetry. For low-potential negolyte species
such as AQDS and DHAQ, this excess was increased (roughly double the volume) to ensure that oxidation of the NCLS electrolyte
by low levels of oxygen in the glove box would not lead to bothsides becoming capacity limiting and fading at the rate of oxidation.
In addition, small aliquots of 50% SOC electrolyte were occasionally added to the NCLS to confirm that the NCLS had not become
limiting.
Ex situ testing.—Cyclic voltammetry of cycled material was performed with a Gamry 3000 potentiostat using a silver/silver chloride
reference electrode, a platinum wire counter electrode and a glassy
carbon working electrode. For the data presented in this study, a scanrate of 100 mVsec−1 was used in all cases and the data were correctedby subtracting the background current of the supporting electrolyte.
For the Fe(CN)6 UV-vis spectra, 4 μL of each electrolyte was diluted
into 2 ml of 1 M KOH and the spectra recorded with 50 ms of integration time using an Ocean Optics Flame-S Spectrometer Assembly.

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A1468

Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)
Results and Discussion

Symmetric cell method description and validation.—Before using
the unbalanced compositionally-symmetric cell method to characterize capacity fade due to reactant instability, it is necessary to validate
and understand the limits of the method by using it on a redox couple
that is known to be stable. To this end, the Fe2+ /Fe3+ redox couple
was selected as a control test case. If all other forms of capacity fadeare controlled for, then a symmetric cell with this redox couple should
be indefinitely stable. Eliminating all other sources of capacity loss istherefore essential to this method. A first step in this direction is to utilize potentiostatic modes of cycling with current cutoff values, rather
than the typically-utilized galvanostatic cycling conditions with potential limits. The capacity realized during galvanostatic cycling withpotential limits is directly related to the polarization resistance of thecell. This includes effects such as changes in membrane and electrolyte resistance, changes in mass transport conditions within thecell, and changes in kinetics due to, e.g., changes in electrode surface
conditions. None of these effects relate to the amount of charge capacity available in the CLS of the cell. By using potentiostatic cycling
with steady state current cutoffs approaching the background current,
it becomes possible to access nearly 100% of the theoretical capacitywith negligible sensitivity to changes in cell polarization resistance.
As discussed later, in some cases it may be more appropriate to restrict the applied overpotentials to purposefully access significantly
less than 100% of the capacity.
Other sources of capacity loss include those due to side reactionsin the NCLS of the cell. With a pH-independent equilibrium potentialof +0.77 V vs. SHE, this redox couple sits near the middle of thewater stability window at pH 0. By constraining the applied overpotential within ± 0.500 V of the equilibrium potential of a 50% SOC
solution, it is therefore expected that no water splitting side reactions
will occur. To be even more conservative and still access above 95%of the theoretical capacity for this one-electron redox reaction, only
± 0.150 V is applied to the CLS of the flow cell in this case. Althoughoxidation of ferrous iron by atmospheric oxygen is not expected inacidic conditions, flow cell cycling was conducted within the glove
box to standardize conditions across all chemistries. In the absence ofsuch heterogeneous and homogeneous side reactions, we anticipate
that the only remaining causes of capacity fade are crossover, precipitation, and electrolyte leakage/permeation. As mentioned previously,
by using the same initial 50% SOC Fe2+ /Fe3+ electrolyte on bothsides of the cell with only a small excess volume on the NCLS, thenet flux of iron species from one side to the other should be canceledand crossover induced capacity fade should be eliminated. It is, however, possible that large disparities in both half-cycle durations andindividual ion permeabilities could eventually lead to differences inconcentration between both sides, especially if the starting SOC deviates significantly from 50%. Potentiostatic cycling to SOC extremes
produces larger concentration gradients across the membrane, which
could lead to changes in overall concentration if one species crosses
faster than the other. This effect would be reversed and nullified during
the following half-cycle if the duration of the reverse SOC extreme
is equal. Hence, it is theoretically possible for vestiges of crossover
effects to remain in symmetric cell testing, but these effects should be
negligible if the starting SOC is nearly 50% and starting volumes arenearly equal, or if a sufficiently thick membrane with sufficiently low
permeability is chosen. If such crossover can be neglected, only decomposition and precipitation would remain as mechanisms that are
intrinsic to the chemistry; neither of these is expected for Fe2+ /Fe3+
at this concentration.
Regarding electrolyte leakage and permeation into cell materials,
this can be overcome with judicious material selection and carefulcell assembly. Although some variation between cell builds is to be
expected, if the anticipated fade mechanisms are managed then a symmetric cell with this redox couple should not experience any capacity
fade. A representative example of such unbalanced compositionallysymmetric cell data is shown in Fig. 2a. In this test roughly 5 ml of
0.1 M Fe2+ /Fe3+ are placed in each reservoir for initial cycling. With

Figure 2. a) Unbalanced compositionally symmetric cell potentiostatic cycling of 0.1 M Fe2+ /Fe3+ in 1 M H2 SO4 , b) with zoomed in vertical scale.

equal volumes on both sides, it is unknown which side is capacity
limiting. Hence, after 15 hours of cycling, 1 ml is added to one sideand 3 ml added to the other to impose a capacity limiting conditionon the smaller reservoir. With 6 ml of electrolyte in the CLS and 8 ml
in the NCLS, the theoretical capacity of the cell is 57.9 C. The average reduction capacity measured over the remaining 15 days of theexperiment is 56.1 ± 0.2 C, which is 96.9% of the theoretical capacity.
It is informative to zoom into the cycling capacity data as in Fig.
2b. This reveals important phenomena that alter the data from theideal horizontal line. Firstly, the excessive noise in the data relates tothe experimental setup and the inherent complexity of a flow battery.
While capacity is contained within the electrolyte, it is ultimately measured via the diffusive mass transport of reactants to the electrodes.
Hence, any change in accessibility between the reactants and the electrodes affects the measured capacity. Formation of stagnant zones in
the reservoirs or permeation of electrolyte to small recesses in tubing
junctions will contribute to losses in capacity. More importantly, any
liquid that becomes stranded due to liquid droplets forming on the

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Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)
walls of the CLS electrolyte reservoir also lead to a spontaneous decrease in measured capacity. When the reservoir is manually shaken
to allow the droplets to fall back into the accessible electrolyte, thecapacity returns to the previous level, as shown in Fig. 2b. Counterintuitively, this capacity recovery happens over nearly a day andover several cycles, rather than in the expected incremental fashion
expected for a discrete amount of liquid returning to the reservoir ona timescale shorter than half a cycle. This demonstrates how capacity
measurements can be affected by other transient phenomena such asconcentration equilibration due to droplets being sequestered in a particular state of charge. Although it is difficult to avoid the formationof droplets due to peristaltic pump pulsation and splashing, there aresome sections of the data that do not appear to have this problem.
Between day 10 and 11, the capacity fade seems regular with a rateof 0.038 ± 0.005%/day; however, this slope is different from those ofother seemingly unwavering sections like between days 2.3 and 2.8,
and between days 3.0 and 3.9. If instead, we merely take the endpoints
to be the ‘recovered’ capacity after each shaking of the reservoir (days
2.2 and 15.6), then the overall fade rate between these points is closer
to 0.031%/day. These numbers highlight the inherent ambiguity of
selecting limits for determination of a true capacity fade rate in thepresence of significant experimental noise.
Once that fade rate is determined, how can it be assigned to a
specific mechanism for capacity loss? To the best of our understanding,there is no mechanism intrinsic to the chemistry that would lead to lossof concentration or redox activity of the dissolved species. Therefore,
only asymmetric crossover or experimental variation due to cell build,such as leakage or permeation, would be suspected in this case. With a
CLS reservoir volume of only 6 ml, however, 0.03%/day is equivalent
to a loss of only 1.8 μL per day. We present these symmetric celldata to emphasize the experimental difficulty of characterizing stableelectrolytes in flow cell geometries and to motivate the developmentof simpler and potentially higher accuracy systems such as the static
cells described in the supplementary information.
The reader accustomed to working with full cells will immediately
notice the difference between the reduction and oxidation capacity
seen in Fig. 2, which initially suggests some form of side reaction or
net crossover of active species. This apparent coulombic inefficiency,
present to a certain degree in all the data in this study, relates to thesteady state background current in each direction. In this case, this difference amounts to a relatively constant 0.24 C between the reduction
and oxidation half cycles. The possibility of side reactions was eliminated by testing for and observing no dependence of this backgroundcurrent on applied overpotential. The possibility of crossover was also
excluded due to the amount of hypothetical crossover required being
substantially more than the total capacity of the system (0.24 C × 792
cycles = 190 C). More generally, the constant nature of the apparent
current inefficiency is simply uncorrelated to any measurable fade rateand hence must be related to instrumental precision. Current offsets,
which fall within the specifications of the potentiostat, were observed
during CV measurements through a simple resistor and would accountfor the measured background currents responsible for the differences
in capacity.d Fortunately, this systematic current bias has no effect onlong term cycling trends such as capacity fade. As explained above,in this case the slow capacity fade is uncorrelated to the coulombic
efficiency and the precision with which the fade rate can be measured
depends on experimental noise due to the pumped electrolyte and cellvariation. In cases where capacity fade is significant and becomescorrelated to the coulombic efficiency, higher precision coulometry
would reduce such systematic biases and could enable more quantitative determination of which operation – oxidation or reduction – isresponsible for capacity fade.17–19 In the case of flow batteries with
dissimilar electrolytes, these measurements are likely to be dominateddAnother way of corroborating this interpretation is by looking at the dependence of theapparent current efficiency on the total capacity or cycling time. The reader will noticesuch differences in later comparisons between large capacity flow cells and small capacity
static cells or between slower galvanostatic cycling and faster potentiostatic cycling in
Fig. 11.

A1469

Figure 3. Semi-log plot of symmetric cell cycling of 50% SOC 0.1 M DHAQ
in 1 M KOH with cycling pauses in different states of charge. Exponential
decay time constants, τ, derived from slopes of each cycling segment denotedin black text. Capacity loss and SOC of CLS during 24 hour pauses denoted inblue text. Due to significant capacity loss, cycling rate increases from roughly
70 cycles/day during first segment, to 133 cycles/day during last segment.
Change in slope after significant capacity loss indicates a higher order processthat is not unimolecular.

by crossover of reactants. Alternatively, cycling pauses such as thosepresented in the next section can be used to gain similar information
for flow battery electrolytes.
Pauses in cycling for investigating different states of charge of
DHAQ.—In contrast to the stability of transition metal ions, organic
molecules such as DHAQ are expected to decompose under certain
conditions. When paired with potassium ferrocyanide at pH 14, Linet al. showed that DHAQ produces a 1.2 V battery with a previously reported fade rate of 0.1%/cycle when limited by the DHAQ negolyte.20
With no sign of decomposition of the oxidized form of DHAQ, andno sign of crossover into the posolyte, the authors attributed this faderate to hydraulic leakage. When converted to a temporal metric, thisyields a capacity fade of about 7.6%/day. As demonstrated above, thiskind of leakage can be reduced to a very low level and would varyacross cell builds. We observed similar fade rates during repeatedtests with DHAQ symmetric cells, however (Fig. S2), and set out toreassess the possibility of chemical or electrochemical decomposition by introducing cycling pauses to hold the electrolyte in different
SOCs.
With few exceptions, capacity fade is very nearly exponential intime (Fig. S2), which is characteristic of unimolecular decomposition
kinetics. Consequently we report the instantaneous temporal fade rate
as the slope of fits such as that in Fig. S2-d), i.e. –lnC/t, whereC(t) is the time-dependent capacity. With unimolecular kinetics, theinstantaneous fade rate is independent of time and its reciprocal is thecharacteristic decay time constant τ; the half-life is τ ln 2. Thus the faderate of the capacity of the 50% SOC 0.1 M DHAQ solution shown in
Fig. S2 is 8.8%/day, or 0.23%/cycle; the corresponding time constants
are 11 days and 430 cycles, respectively. For some experiments thatare long enough for a substantial fraction of the initial capacity to
be lost while the cycling regimen is punctuated by pauses at various
SOC, the instantaneous fade rate is observed to change with time.
This suggests kinetics that are more complicated than unimolecular;
nevertheless we report the instantaneous fade rate that we measure
and we leave it to future work to establish the kinetics and the time
constants.
The results of a long DHAQ cycling experiment in which holds
at various SOC were introduced are shown in Fig. 3. The data show

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Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)

capacity fade behavior for which electrolyte leakage is an untenable
explanation. Each data point corresponds to one potentiostatic cycle
and the horizontal axis is plotted as time to demonstrate the directconnection between capacity fade and time. The cell contains roughly
5 ml of 50% SOC electrolyte on the CLS, and 10 ml of the same 50%
SOC electrolyte on the NCLS with an initial cycle period of roughly
20 minutes. With no induction time, the capacity starts to fade immediately at an instantaneous fade rate of 0.076%/cycle or 5.8%/dayover the first segment of cycling data. The cycling is then stopped with
the CLS electrolyte in an oxidized state, still being pumped throughthe cell. The cell is held in this state for 24 hours and then cycling
is continued. As indicated in Fig. 3, no appreciable capacity was lost
during this pause in an oxidized state. When cycling resumes, the instantaneous fade rate of the 2nd segment is slightly lower at 5.3%/day.
Again, cycling is paused for 24 hours and pumping is continued, with
the CLS electrolyte in a 50% SOC this time. A small amount of capacity (0.7 C) is lost during this pause. Cycling is then continued forone more day, exhibiting a fade rate of 5.6%/day, at which point cycling is paused again and pumping continued with the CLS electrolyte
in a fully reduced SOC. The unbalanced compositionally-symmetriccell lost 11.2 C or 18%/day during this pause in a reduced state –considerably more so than during cycling. This entire process was
replicated once more to verify the phenomena observed. The faderates for the last four cycling segments decrease monotonically andequate to: 2.9%/day, 2.7%/day, 2.4%/day and 2.0%/day, respectively.
The capacity loss for the 2nd set of pauses also reflects the same behavior as the first set, in which no measurable capacity fade was detectedduring a pause with the CLS in the oxidized state and only a small
amount of fade during a pause at 50% SOC. Indeed, in both cases
the CLS electrolyte loses more capacity when paused in a reduced
state than it does during cycling and during pauses in other states.
This can be diagnosed most quickly as an offset in the capacity data
between periods of cycling. This indicates that something happens to
the reduced form of DHAQ that causes it to become redox inactive,
leading to capacity fade. This also explains why no sign of decomposition was observed in the previous study, which had focused purely
on the oxidized form of DHAQ. To verify that this loss of capacity
is neither an artifact of cycling conditions, nor merely a recoverable
capacity loss due to SOC asymmetry caused by side reactions, nor aloss of CLS electrolyte volume due to leakage, the concentration ofredox active species was measured via cyclic voltammetry. The CLS
electrolyte after cycling has a smaller peak current than before cycling
as shown in Fig. S3 of the supplementary information. After subtracting the KOH background, this change in peak current corresponds toroughly a 40% difference in concentration, which is slightly less than
the total loss in capacity seen in the symmetric cell cycling. Due to thecomplexity of the flow battery system, this kind of post-mortem testis necessary to confirm whether there is a real loss of redox active material. The mechanism behind this loss of redox activity and the slight
discrepancy between capacity measurements and CV measurements
are beyond the scope of this work but are the subject of an ongoingstudy.
That the instantaneous fade rate during cycling exhibits a gradual decrease as the capacity is lost in Fig. 3 suggests a mechanism
more complicated than simple unimolecular decomposition, e.g. thedecomposition rate may depend on the concentrations of oxidized andreduced species, and possibly on those of any decomposition products.
Indeed, the fade rates during pauses indicate a superlinear relationshipbetween fade rate and the concentration of reduced species. That thefade rate during pauses at 100% SOC is higher than during cycling
suggest a purely chemical form of decomposition that is unrelated tothe cycling rate but is time-denominated and related only to the SOC.
Due to the significant loss in capacity, the cycle period has decreased
from the initial 20 minutes to almost 10 minutes during the last cycling
segment, which means that each molecule is being cycled twice as frequently at the end; yet the temporal fade rate has decreased comparedto the beginning of the experiment. Clearly a cycle-denominated capacity fade rate is not at all representative of the expected lifetime of aflow battery based upon this material. If linear kinetic approximations

Figure 4. Semi-log plot of unbalanced compositionally symmetric cell cycling of 50% SOC 0.1 M DHAQ in 1 M KOH with changing reduction overpotential. Oxidation overpotential limits set at +300 mV throughout experiment,with changes in reduction overpotential limits indicated in graph. Demonstrates
decrease in temporal fade rate due to restricted SOC range.

of fade rates are to be used at all, then a time-denominated capacity
fade rate would be more appropriate than a cycle-denominated faderate.
In addition to using cycling pauses in different states of charge,
restricting the accessed SOC range during cycling can also provide
similar information. Although the SOC range can be constrained byincreasing current cutoffs or restricting the time for which potentials
are applied, this can also be done by changing the applied potentials.
Typically, potential limits for electrochemical stability are establishedby cyclic voltammetry during reactant screening. The example shownin Fig. S3 indicates no side reactions within 300 mV of the equilibrium potential. Using the Nernst equation, equivalent overpotentials
can be estimated for symmetric cell cycling, to achieve the desired
SOC range. Similar to cyclic voltammetry in which stability is determined by expanding the overpotential range to look for side reactions,
it is first recommended that stability limits be established by increasing overpotentials to look for variations in fade rate. In the first sectionof data in Fig. 4, large overpotentials (±300 mV) are applied during
cycling, which lead to an instantaneous fade rate during cycling of
6.3%/day before the reduction overpotential is changed to −200 mV.
This reduction overpotential is still sufficient to access over 99%
of the capacity for this two-electron reaction, leading to a relatively
unbroken trend in the capacity and fade rate (6.1%/day). The negligible change in fade rate further suggests that the reduced form of
DHAQ is chemically unstable, rather than electrochemically unstable, because electrochemical instability should be very sensitive toapplied overpotential. By the end of the second section the capacity
has dropped to 87.8 C, at which point the reduction overpotential ischanged to −100 mV and the accessed capacity drops to 80.9 C or
92% of the remaining amount. This slight constraint on the accessed
SOC (i.e., the concentration of reduced DHAQ) decreases the faderate during cycling to 4.4%/day. When the reduction overpotential isfurther changed to only −50 mV, the accessed SOC drops to 75%of the previous level and the fade rate during cycling drops to only
1.6%/day. This further demonstrates that the capacity fade seems to bea superlinear function of the concentration of the reduced state of themolecule. When the reduction potential is returned to −200 mV thecapacity climbs to 85.4 C and the fade rate during cycling returns to
6.2%/day, which is quite close to the value of 6.1%/day recorded forthe previous segment with an applied overpotential of −200 mV. The
total amount of capacity lost during the restricted SOC range cycling

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Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)

Figure 5. a) Semi-log plot of unbalanced compositionally symmetric cell
cycling of 0.1 M H2 AQDS/AQDS in 1 M H2 SO4 with cycling pauses invarious states of charge. b) Linear plot of close-up of data showing how fadeis quantified during cycling pauses. Average cycling rate during test is 80
cycles/day.

amounts to (80.9-79.4) + (64.3-63.6) = 2.2 C, which is in reasonableagreement with the difference between the end of the first −200 mVsegment and the beginning of the second −200 mV segment (87.885.4 = 2.4 C). This suggests that the capacity lost during the restricted
SOC cycling is real, and that, when taken with ex situ evidence suchas CV, full SOC range potentiostatic cycling is an effective way of
directly quantifying real loss in capacity.
Transient phenomena and strategies for measuring slow capacity fade.—The transient behavior observed in the case of
Fe2+ /Fe3+ appears to be primarily related to mechanical pulsationand hydraulically-induced noise that can be reduced with proper system design. It is possible, however, for such transient phenomena to
be intrinsic to the chemistry under study or the methodology being
used, as in the example provided in Fig. 5. Although this test uses thesame cycling-pause-cycling methodology as in the previous example,
the acidic AQDS electrolyte used here displays very different capacity fade behavior. For 5 ml of a 0.1 M limiting electrolyte the cycling
capacity starts off at 105.5 C, which is slightly above the expectedvalue for a two-electron reaction. The capacity fades immediately andregularly throughout the first cycling section. In subsequent sections,A1471

the cycling capacity contains reproducible transient behavior related
to the SOC during the prior pause. After pauses in the oxidized state,subsequent cycling capacity has a positive curvature in which the faderate starts high and then approaches a constant, steady state slope. After pauses in the reduced state, subsequent cycling capacity starts low
but increases; the resulting positive slope monotonically decreases until reaching the steady state fade rate. Both of these phenomena seem
to offset each other when the CLS AQDS electrolyte is in a 50% SOC,
as observed from the steady slope upon initial cycling and after thepause in that state. In all cases, these transient occurrences comprise< 1% of the capacity and are therefore difficult to measure with ex situ
methods. One possible explanation for such transient behavior could
be adsorption or deposition of AQDS on the carbon paper electrodes.
If such a process were not entirely reversible, it could potentially affectthe total capacity over time scales longer than one cycle. Our working
hypothesis for the transient capacity fade is molecular crossover toand from the CLS electrolyte. As explained previously, due to thetemporarily enforced concentration gradients during pauses in 0% or
100% states of charge, there exists a strong diffusional driving force
between the CLS and NCLS electrolyte. If the reduced form permeates more quickly through the membrane than the oxidized form, thenboth temporary phenomena can be rationalized. During a pause in the
reduced state, if more H2 AQDS crosses out of the CLS than is replaced by AQDS crossing in, then a loss of capacity will be observed.
Conversely, when paused with the CLS in the oxidized state, less
AQDS will cross out of the CLS than H2 AQDS crossing in from theNCLS, leading to an increase in capacity. Although crossover of either
form is very slow, the high concentration difference and long duration(3–5 days) of the pauses could lead to such an effect. This suggeststhat symmetric cell capacity measurements could also become a useful tool for characterizing the in-situ crossover of the reduced andoxidized forms of any redox active species. Crossover as the source oftransient capacity is also consistent with the observation that no transient phenomenon is observed upon initial cycling and after a pause in
50% SOC, when the electrolyte compositions on the two sides of thecell are balanced. As demonstrated later, potentiostatic cycling capacity is unaffected by cell resistance and therefore thicker membranes ormore membrane layers can be used to mitigate the effect of different
crossover rates.
This example of AQDS capacity fade is provided to illustrate a few
critical challenges for characterizing the stability of new redox active
molecules. Ignoring the pauses and transient behavior, the overall
fade rate calculated from the slopes of all linear sections averages to
0.08 ± 0.02%/day. Accurately determining such a slow rate of capacity
fade therefore requires very long cycling time, especially when transient phenomena such as those seen here are present. Cycling pauses
were also imposed to determine the source of the capacity fade, as wasdone for DHAQ. Due to the transient behavior and much slower faderate of AQDS, these pauses were extended up to 5 days to accentuate
any differences among the fade rates. To quantitatively compare theamount of fade during these pauses it was necessary to fit each linear
section of data independently and compare the vertical offset between
the fitted lines. This vertical distance would then be divided by the duration of the pause and then combined with the average fade rate during
cycling to give a fade rate during the pause. Although pauses in a reduced state of charge consistently led to more capacity fade (roughly
0.2%/day), the case is not so clear for other states. More pauses ofeven longer duration would be needed to obtain sufficient statistics,
but this experiment was stopped after 36 days due to hardware malfunctions. Maintaining controlled experimental conditions for suchextensive periods is challenging; furthermore, introducing pauses incycling further complicates the experiment because no data are collected during that time and any unexpected sources of noise would not
be revealed. Fortunately, in this case the hydraulically-induced noise
appears to have been minimal compared to the 0.08%/day fade rate.
Although significantly lower than that of DHAQ, the fade rate of
AQDS would still amount to a substantial loss over the first year offlow battery operation. In both cases, capacity is lost over time during
cycling pauses in the reduced form, indicating a chemical instability

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A1472

Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)mechanism. In the case of DHAQ the loss was sufficiently large tobe verified via cyclic voltammetry. In the AQDS case however, only
about 3% of the capacity was lost over the course of the 36-day
experiment, indicating a small concentration change that would be
difficult to verify via other experimental methods. Moreover, as morestable molecules are developed, it will become necessary to quantify
even slower rates of capacity fade and decomposition. One way to
do so would be to adopt accelerated stress testing procedures. If apurely chemical form of decomposition is expected, then elevatedtemperatures could be used. If electrochemical instability is suspected,
then more rapidly cycling cells could be utilized. Both methods were
implemented in the case of AQDS.
Instead of using symmetric flow cells with large electrolyte volumes due to unavoidable tubing and reservoirs, small static cells weredeveloped akin to coin cells for rapid cycling of materials. This cell,
pictured in Fig. S1, was used for testing the fade rate of AQDS using
the same potentiostatic cycling conditions and same capacity limiting
symmetric cell methodology as with the flow cell case. The low experimental noise in the capacity data in Fig. S4 attests to the hydraulicstability of the static cell. With only a tiny volume of electrolyte imbibed directly into the electrode the cell achieved roughly 35 cycles
per hour. In contrast, the flow cell test reported in Fig. 5 averaged
only about 3 cycles per hour. The fact that both cycling tests measure
similar temporal fade rates of around 0.1%/day (See SI) gives further
support to our conclusion that the capacity fade is time-denominated.
This order of magnitude example also underscores the need for standardizing parameters such as volume and concentration when using acycle-based metric.
Another way of confirming chemical decomposition mechanisms
is by looking for signs of decomposition after storage of electrolytes.
For slowly decomposing species such as AQDS, this process canalso be accelerated and analyzed for Arrhenius behavior by storageat elevated temperature. The example in Fig. S5 provides additional
evidence for the chemical instability of the reduced form of AQDS
after storage for two weeks at 45◦ C. The appearance of new peaks
and the decrease in integration of the original AQDS aromatic peaksin the NMR spectrum of the sample stored in the reduced form bothindicate decomposition.
Material interactions and pH effects: stability of Fe(CN)6 .—One
advantage of flow cells over static cells is the possibility of pumping
most of the electrolyte into the reservoirs during cycling pauses andthe ability to visually inspect electrolytes and take aliquots for ex situ
testing. This permits additional diagnostic methods for understanding
sources of capacity fade such as the possibility of material interactions
as in the case of Fe(CN)6 described here. In the first incarnation of thisexperiment, the CLS of the symmetric flow cell is filled with 4 ml of
50% SOC 0.1 M K4 [Fe(CN)6 ] /K3 [Fe(CN)6 ] in 1 M KOH, while the
NCLS contains 6 ml of the same electrolyte. It can be seen in Fig. 6athat the capacity starts off as expected for this one-electron reaction,
but after roughly 4 hours, the capacity begins to decay exponentiallyat an instantaneous rate of 67%/day, regardless of whether cycling ispaused with the CLS electrolyte in an oxidized or a reduced state.
As with the other redox-active compounds presented thus far, thisbehavior suggests some form of chemical capacity fade mechanism
independent of electrochemical cycling that, in this case, may seemto affect both forms of Fe(CN)6 . This is not the case however. The
initial few hours of stable cycling suggests that the capacity of the
CLS reservoir can always be accessed if there is sufficient capacity
in the NCLS. As the cell lost most of its capacity, the electrolytes
lost their initial clear yellow coloration. A post-mortem UV-vis absorption spectrum of both CLS and NCLS electrolytes showed mostly
ferrocyanide with very little ferricyanide present. This indicates aheterogeneous or homogeneous side reaction that reduces the ferricyanide to ferrocyanide. This reaction would produce an asymmetry
in the SOC and, at some point, make both reservoirs capacity limiting.
This is illustrated in Fig. 6b, where the accessed capacity of each electrolyte is the distance along each axis traversed by the projection of
the diagonal arrow. At the start of the experiment, t0 , the entire SOC

Figure 6. a) Semi-log plot of unbalanced compositionally symmetric cell
cycling of 0.1 M K4 [Fe(CN)6 ] /K3 [Fe(CN)6 ] in 1 M KOH showing capacity
fade rate that is unchanged during cycling pauses in alternating reduced oroxidized states, starting with a pause in the oxidized state, b) Representationof temporal evolution in accessed SOC in both CLS and NCLS for the same
cycling data, where each time point tn denotes a time point from cycling datain part a) and length of arrows signifies real cycle capacity, i.e.: Capacity andarrow length at t2 are half of initial at t0 .

range of the CLS and only a fraction of the NCLS is cycled. As time
progresses, excess ferricyanide in the NCLS is spontaneously reduced
to ferrocyanide, restricting the SOC range of the NCLS and moving
the arrow to the left in the SOC diagram.e Once there is no longer
any excess of ferricyanide in the NCLS at time t1 , both sides become
capacity limiting and the left point of the arrow reaches the left top
corner of the SOC diagram. At this point, spontaneous reduction of
ferricyanide shifts the arrow further to the left, e.g. t2 , restricting the
SOC range of both sides and causing precipitous capacity fade. It is
important to distinguish two separate capacity fade regimes in thisexperiment. Although the initial portion of the data is relatively stablecompared to the exponential decay observed after both sides become
limiting, it does exhibit some capacity fade. This ‘apparent capacity
loss’ may be due to a change in the open circuit potential of the cell, as
described by Chen et al.3 As the ferricyanide in the NCLS is reduced,
the average SOC and therefore equilibrium potential of the electrolytechanges. This affects the overpotential applied during cycling and,
consequently, the capacity accessed. This was termed an ‘apparent’capacity fade, as opposed to real capacity fade, because it can becompensated for by changing the applied potential. Once both sides
eFerricyanide in the CLS is also reduced by the same mechanism, shifting the arrow
downward as well. The overall trend is a superposition of these two changes.

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Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)

A1473

Figure 7. Semi-log plot of unbalanced compositionally symmetric cell cycling of 0.1 M K4 [Fe(CN)6 ] /K3 [Fe(CN)6 ] in 1 M KOH showing capacity faderate is decreased during pauses when pumping is stopped, cell is emptied, andelectrolyte is kept in reservoirs. Blue dashed line is linear fit to clarify offset
before and after third sections of cycling, and lack of offset between last two
cycling segments.

become limiting, however, the capacity fades more quickly, following
an exponential decay curve, and no change in cycling conditions canrestore this lost capacity. Fortunately, as described later, in this casethe capacity fade is reversible by other means.
In the 2nd incarnation of this experiment, the same composition of
initial electrolytes was placed on both sides of the cell, with 5.5 mlon the CLS and 7 ml on the NCLS to verify the same initial stablecapacity phase as before, as shown in Fig. 7. After about 5 hours,
which we interpret as the point at which both sides become capacity
limiting, the capacity fade rate becomes high during cycling periods
with an instantaneous rate of roughly 45%/day. In this experiment,
however, the electrolytes were occasionally pumped out of the celland into their respective FEP reservoirs during cycling pauses. The
first instances of these empty cell pauses show a slight positive effecton capacity retention, as indicated from the dashed blue linear fit line
to the third cycling segment. To accentuate this effect, the empty cellpause was repeated for a much longer duration. During this nearly
16 hour pause, the cell lost only 7% of its capacity, in contrast to thepreceding instantaneous cycling fade rate of 43%/day. This decreasein fade rate when the cell is emptied implies an interaction between theelectrolyte and the cell materials. Moreover, the 7% of lost capacity
corresponds roughly to the amount of electrolyte (∼400 μL) thattypically remains imbibed in the electrodes and is not pumped outof the cell – this suggests that the electrodes are the cell materials
reducing the ferricyanide.
The reduction of ferricyanide to ferrocyanide aided by carbon paper electrodes was established more systematically in a subsequent
experiment. In this case, a symmetric cell was loaded with 6 ml of
50% SOC electrolyte on each side and cycling begun immediately.
Meanwhile, 6 ml of the same electrolyte was also stored in two falcon tubes. In one tube, the same amount of electrode material as isused in the cell was immersed into the electrolyte at the same time
as electrolytes were pumped through the cell. Some manual shakingof the stored electrolytes was performed to promote a homogeneousconcentration. Both visual inspection and UV-vis spectrophotometry were used to compare the electrolytes before and after cycling
and storage. Cycling capacity had degraded to nearly nothing afterroughly a day (Fig. S7). Both the electrolytes in the cell looked nearlycolorless and visually identical after cycling (not shown), as did theelectrolyte stored with the carbon paper (Fig. S8). All electrolytes

Figure 8. UV-Vis spectra comparing 0% SOC 0.1 M [Fe(CN)6 ]4− , 100% SOC
0.1 M [Fe(CN)6 ]3− and initial 50% SOC 0.1 M K4 [Fe(CN)6 ] /K3 [Fe(CN)6 ]electrolyte with a) CLS and NCLS electrolytes after symmetric cell cycling,
and b) after immersion of carbon paper. c) CV of ferrocyanide, electrolyte
stored with carbon paper and cycled electrolytes.

were analyzed spectrophotometrically after roughly 33 hours, andcompared with freshly prepared 0% SOC (0.1 M ferrocyanide) and
100% SOC (0.1 M ferricyanide). As displayed in Fig. 8a, spectrafrom both electrolytes overlay nearly perfectly with that from thepure ferrocyanide and show no signs of ferricyanide remaining. Similarly, the initial 50% SOC electrolyte stored with carbon paper also
exhibits the same behavior, as seen in Fig. 8b. As a final verificationthat the ferricyanide is not destroyed but is simply reduced to ferrocyanide, cyclic voltammograms of final electrolytes were compared to

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A1474

Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)that of a 0.1 M ferrocyanide solution. The voltammograms, overlaid inFig. 8c, confirm the nearly equivalent redox activity of all electrolytes.
The exact mechanism by which ferricyanide is rapidly and heterogeneously reduced to ferrocyanide when in contact with carbon paper
is currently being investigated. It does, however, relate to the potential
vs. pH behavior of the [Fe(CN)6 ]4− /[Fe(CN)6 ]3− redox potential vs.
the water-splitting window, as suggested by the slower capacity fadeat pH 12 depicted in Fig. S9.
This example was chosen because of its importance to several
existing flow battery chemistries,21–23 because the spontaneous reduction of ferricyanide can lead to confusing results if this redox couple
is paired with other reactants in a full cell. The capacity of such acell would be expected to fade at similarly drastic rates as seen here
once the ferricyanide ceases to be in excess. In addition, regardless of
ferricyanide excess, the spontaneous reduction would always add tothe charging capacity of such a cell without enhancing the discharging capacity and would therefore lead to lower current efficiency.
This effect needs to be considered when attempting to deconvolute
other sources of current efficiency loss such as crossover. Fortunately,
this spontaneous reduction is not an irreversible decomposition of
the redox-active compound and would not necessarily affect overall
battery lifetime, as it can be reversed by rebalancing methods suchas electrolysis. It should be noted, however, that Fe(CN)6 may also
suffer from irreversible decomposition due to loss of CN− groups andeventual formation of iron hydroxide at high pH.24–27 These reactions
are expected to occur at much longer timescales than those studied
here.
Redox active molecule protection: A viologen-based example.—
The examples provided thus far may connote doom and gloom fororganic-based redox flow batteries; however, none of the molecules
described previously were specifically designed for long lifetime. If
the mechanism for molecular decomposition is known, then strategies may be developed to protect against the mechanism. The addition of protective functional groups to a redox active molecular core is one such approach. For example, Beh et al. usedbis(trimethylammoniopropyl) (BTMAP) functional groups to simultaneously enhance the aqueous solubility and stability of viologen.13
Beh et al. hypothesized that the coulombic repulsion imparted by
the charged ammonium end groups would suppress the bimolecularinteraction that is known to lead to viologen dimerization and decomposition. BTMAP functionalization leads to extremely high capacity
retention as illustrated by the comparison between the one-electron reduction of methyl viologen and BTMAP-viologen in Fig. 9. Whereas
the capacity of the methyl viologen symmetric cell faded at a rateof roughly 1.5%/day, no measurable capacity fade was observed for
the BTMAP-viologen. Due to the remarkable stability of BTMAPviologen and the inherent capacity fade measurement limits discussed
previously, accelerated stress testing may be necessary for an estimateof the lifetime at room temperature.
Another advantage of using symmetric cells is the ability to know
the overpotential being applied to the electrolyte under study, as opposed to the unknown overpotential split when using dissimilar electrolytes with different kinetics and mass transport properties. Thispermits investigation of stability limits much like in cyclic voltammetry with a three-electrode cell. Although BTMAP functionalization
can protect the singly reduced form of viologen, the doubly reduced
compound is still expected to be highly unstable. Using the same CLS
electrolyte that was cycled stably in Fig. 9b, a new test was performed
to probe this second reduction step. With a small aliquot removed
between tests the cycling begins with roughly the same stable capacity. The potential limits are then increased in steps from ± 150 mV,
to ± 200 mV, ± 250 mV and ± 350 mV. These steps only lead tomarginal gains in capacity as displayed in Fig. 10. Once the potential
is increased to ± 450 mV, however, the capacity increases from near
40 C to roughly 55 C, and then to 65 C once the potential is raised to± 500 mV. At this point the capacity begins to fade rapidly until thecycling potential is reset to the original ± 150 mV. While insufficient
to convert all the BTMAP-viologen molecules into their doubly re-

Figure 9. Comparison between unbalanced compositionally symmetric cell
cycling of a) 8 ml of 0.1 M Methyl Viologen in 1 M NaCl and b) 5 ml of 0.1 M
BTMAP-Viologen in 1 M NaCl.

duced forms, the potentials applied demonstrate the instability of the
doubly reduced form. Moreover, the loss of redox active material ispermanent. The capacity lost during large overpotential cycling wasroughly 5% of the initial 65 C. When stable cycling is recovered, thecapacity has decreased from the initial 39 C to 37 C, a similar lossof roughly 5%. Hence, destroyed molecules remain destroyed, but
surviving molecules remain redox active.
Recommendations and implications for the field.—Both galvanostatic cycling conditions and %/cycle metrics are inherited from thesolid-state battery field. In these systems, the rate at which ions areintercalated or deposited to and from electrode structures may impact the lifetime of the battery and hence galvanostatic cycling isused to control these rates. In the case of intercalation batteries, large
current spikes could lead to irreversible damage to the electrodes,
whereas for deposition-based batteries higher rates may lead to undesirable dendrite structures causing early battery failure. For redox
flow batteries with permanently dissolved active species, these kinds

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Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)

Figure 10. Unbalanced compositionally symmetric cell cycling of 5 ml of
0.1 M BTMAP-Viologen in 1 M NaCl at increasing overpotentials to access
2nd reduction step. Annotations indicate points at which cycling overpotentials
were changed.

of problems are not expected. Capacity fade measured in %/cycle
is relevant for solid-state batteries for the same reasons. For flow
batteries capacity-limited by crossover of reactants this metric may
also be indirectly relevant because crossover mainly occurs during
cycling (and therefore pumping) of reactants. Even for these systems,
however, cycle-based fade rates cannot be a meaningful basis of comparison unless other parameters that relate to total capacity, such asvolume and concentration, are standardized. Claims of high capacity retention such as the 0.0001%/cycle observed for the fast-cycling
symmetric AQDS static cell may seem impressive, but this becomes
irrelevant for long term applications if the species decomposes irreversibly at a rate of 0.1%/day. Although this rate is significantly slower
than the rate of bromine crossover in a full AQDS/Br2 cell,3 it is an
important underlying limitation that cannot be offset by flux balancing
or electrolyte rebalancing methods.28,29 Ideally, each novel electrolyte
being developed should be individually characterized to identify suchfundamental limitations to RFB lifetime. To the authors’ knowledge,
thus far, no data in the RFB literature and none of the results in thisstudy indicate any causal relationship between the number of times
a dissolved redox active species is cycled and the capacity retentionof a single electrolyte. For this reason, we stress the importance ofusing temporal or calendar fade rates for characterizing the stabilityof novel electrolytes. When paired together and used in a full cell,the stability of the shortest-lived active species may be the primarycontributor to capacity fade. Unless each electrolyte is known to be independently stable, temporal fade rates should be included in full cellstudies.
As in the choice of cycling metrics, cycling conditions should also
be selected based on the parameters to be characterized. On one hand,
purely galvanostatic cycling or constant-power cycling may be appropriate for holistic simulation of long-term battery performance andapparent capacity retention, which includes cell resistance and otherstack related effects. On the other hand, if the purpose of cycling is
to characterize the true charge capacity retention of a full cell or thestability of a single electrolyte, then the cycling conditions should bechosen so as not to be affected by parameters unrelated to this capacity. Galvanostatic cycling is not ideal for deconvoluting different
sources of real capacity fade due to the confounding effects of cell
polarization resistance that change the SOC at which the voltage limits are reached, leading to artifacts in the measured capacity. Apart

A1475

Figure 11. Semi-log plot of unbalanced compositionally symmetric cell cycling of DHAQ, showing equivalence of potentiostatic and galvanostatic cycling with same potential holds. In both cases cycling was performed by imposing voltage holds at ± 200 mV and switched when current density dropped to
1 mA/cm2 . Galvanostatic conditions were started at 10 mA/cm2 and increased
to 20 mA/cm2 after roughly 2 days. Vertical arrows indicate times at which
a 0.1  resistor was added and removed in series with the cell. The jumps inmeasured capacity demonstrate the dependence of measured cycling capacity
on cell resistance when strictly galvanostatic conditions are used.

from using ex situ methods, the most direct way of characterizing thetrue capacity of the CLS of a cell is with a potentiostatic charge anddischarge until a steady state background current is reached. If large
initial current spikes like those shown in Fig. 1b need to be avoided
due to hardware current constraints, then galvanostatic cycling withpotential holds at the end of every half-cycle can also be used. Suchcycling can also be useful for tracking the evolution of oxidation andreduction voltage profiles. Beside this, the only difference between
this and potentiostatic cycling is the time distribution of the applied
potential. Hence, these methods can be used interchangeably as longas the capacity fade is not sensitive to the applied overpotential. This is
the case for DHAQ, for which the capacity fade was previously shown
to be a function of the state of charge and independent of the applied
overpotential. Using the same potentials and background current cutoffs, the two methods yield the same capacities and fade rates during
symmetric cell cycling of DHAQ as shown in Fig. 11. To prove thatthese methods are unaffected by changes in cell resistance, a 0.13 
resistor was added in series (or subsequently removed) with the cellat approximately the midpoint of each cycling segment. Whether the
resistor is added or removed, both potentiostatic and galvanostatic cycling with potential holds show no change in the capacity fade trend.
In contrast, purely galvanostatic cycling segments show a step change
in the accessed capacity when the resistor is added or removed, whichis exacerbated at higher current densities. The value of the resistor
was chosen to represent a 25% increase in cell resistance (see Fig.
S10 in SI). Such sudden changes are unlikely during typical flow battery operation; however, a gradual drift of the polarization resistanceof this magnitude is entirely plausible due to changes in membrane
conductivity, changes in mass transport of electrolytes or changes inelectrode surface properties. Whether positive or negative, such a drift
in resistance would lead to a drift in capacity during strictly galvanostatic cycling, which would be superimposed onto the real capacity
fade curve. To deconvolute this effect, it is possible to craft a compromise between simulating battery usage in deployment and accuratecharge capacity by interposing potential holds at the end of one out of
every n cycles, as introduced by Beh et al. using n = 10.13 It is also
possible to use low current density (e.g., C/20) cycling to access the

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A1476

Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)

entire capacity of the cell with minimal effect from polarization.30–33
This method works for solid electrode single electrolyte systems suchas lithium ion cells but is not as useful for flow cell characterization.
Due to the presence of hydraulically-induced and other sources ofnoise in flow cells, performing a low current density capacity test over
20 hours is not the most time-efficient way for accurately determining
capacity fade trends. Moreover, if crossover of reactants is suspected,
as in the case of AQDS previously mentioned, then such a slow measurement would inherently amplify such crossover effects. Because
there is no experimental evidence nor any plausible mechanism thus
far indicating that dissolved redox reactants are intrinsically sensitive
to the cycling rate (or current density), then we see no reason to limit
cycling conditions to lower currents. As seen in Fig. 11, potentiostatic
cycling conditions lead to a higher sampling rate of the capacity,
thereby providing more detailed information for identifying transient
effects, quantifying experimental noise and measuring capacity fade
more quickly. For the purpose of establishing the stability of novel
electrolytes with the symmetric cell method described in this study, we
therefore recommend such potentiostatic cycling conditions wheneverpossible.
While it may be more industrially relevant for cycling of full cells
to be conducted galvanostatically, the potentiostatic methodology withholds at different states of charge is better for elucidating the limitsof electrolyte stability and the mechanisms behind these limits asquickly as possible. Complete characterization of new electrolytes is
considerably simplified with proper control of variables, as is enabledby the unbalanced compositionally-symmetric cell method.

e.

2. General methodological results:
a.

b.

c.
d.

Conclusions
This study contains specific results pertaining to the stability andcapacity retention of several flow battery electrolytes in addition tomore general results concerning the methodology to measure andcharacterize such properties. The primary conclusions of the study
can be summarized as follows:
1. Electrolyte-specific results:
a.

b.

c.

d.

Pauses in cycling of a DHAQ unbalanced compositionallysymmetric cell with the CLS at different SOC demonstrate thatthe capacity fade rate is independent of the number of chargedischarge cycles and depends only on SOC and time. The faderate is of order 10%/day when the electrolyte is in the reduced
state, of order 1%/day when the electrolyte is at 50% SOC, andis imperceptibly low when the electrolyte is in the oxidized state.
Consequently the projected lifetime during use in the field depends on the anticipated temporal SOC distribution.
Cycling of AQDS unbalanced compositionally-symmetric cells
leads to a capacity fade rate of 0.08 ± 0.02%/day, independentof cycling. The same rate is found in rapid cycling of a static,
non-flow cell. Transient behavior in the flow cell, spanning several hours and tens of cycles after the resumption of cycling,hindered the measurement of the SOC-dependence of the faderate, but the highest capacity fade rate was in the reduced form, atroughly 0.2%/day. Chemical decomposition was observed of the
reduced form but not of the oxidized form of AQDS with NMR
spectrometry of samples held for two weeks at 45◦ C.
Cycling of K4 [Fe(CN)6 ] /K3 [Fe(CN)6 ]in 1 M KOH is characterized initially by a low capacity fade rate that becomes high whena SOC imbalance between the two sides grows to the extent thatboth sides become capacity-limiting, as illustrated in Fig. 6b. The
imbalance development is caused by reduction of the [Fe(CN)6 ]3−
to [Fe(CN)6 ]4− by contact with the carbon paper electrodes. The
rate is slower by a factor of approximately 5 when the pH isdropped from 14 to 12.
It is possible to protect molecules from decomposition. Unbalanced compositionally-symmetric cell cycling of methyl viologen exhibited a capacity fade rate of 1.5%/day, whereas cyclingof BTMAP-viologen over a period of two weeks led to an imperceptibly small amount of capacity loss.

The overpotential imposed during oxidation and reduction canaffect molecular lifetimes if it is sufficient to create unstablespecies. When the reduction overpotential of BTMAP-viologen
was increased enough to initiate reduction by a second electron
per molecule, irreversible capacity fade resulted.

e.

The unbalanced compositionally-symmetric cell methodology
presented here is useful for revealing and quantifying different
mechanisms for capacity fade in redox flow batteries based onmolecular energy storage. The precision limitations of the method
are determined by flow cell and reservoir hardware, as shown by
the Fe2+ /Fe3+ test case.
Evaluating small capacity fade rates is better performed with potentiostatic cycling than with galvanostatic cycling because thelatter is vulnerable to artifacts caused by drifting cell resistance.
Galvanostatic cycling with potential holds at the voltage limits yields indistinguishable results from those of potentiostatic
cycling between the same voltage limits when the overpotentials at the voltage limits are low enough not to affect molecular
lifetimes.
Average temporal capacity fade rates during cycling can varywith the SOC range being accessed during cycling, if there is anonlinear dependence of temporal fade rate on SOC.
Cases in which capacity decay occurs during pauses in cycling
indicate a purely chemical mechanism. Such capacity fade canbe recoverable if it is due to a simple change in oxidation
state, as in the case of ferricyanide converting to ferrocyanide,or permanent if the molecule decomposes irreversibly as in thecase of the reduced form of AQDS. Electrochemical decomposition can also be assessed by increasing the applied overpotentials beyond the level necessary for essentially complete conversion of redox active species during cycling. No change incapacity fade rate beyond this level, as in the case of DHAQ,
similarly indicates a purely chemical mechanism for capacity
fade.
In all cases reviewed to date of organic or organometallic redoxactive species, capacity fade is independent of the numberof charge-discharge cycles imposed, suggesting that dissolved
molecular reactants are unaffected by changes in oxidation state
and that a cycle-based metric is unsuitable for predicting flow battery lifetime. Instead, capacity fade is reasonably consistent witha time-denominated exponential decay phenomenon in which thetime constant can be SOC-dependent. Hence, in all reports ofcycling performance of flow cells utilizing redox-active organics
or organometallics, the capacity fade rate per unit time should bereported.
Acknowledgments

This research was supported in part by U. S. DOE ARPA-E award
DE-AR-0000767, by Innovation Fund Denmark via the Grand Solutions project “ORBATS” file nr. 7046-00018B, by the Massachusetts
Clean Energy Technology Center, and by the Harvard School of Engineering and Applied Sciences. We thank Daniel Pollack for assistance with development of chemical storage methods, and Prof.
Jeff Dahn at Dalhousie University and Jaroslaw Syzdek at BioLogic USA LLC for technical advice concerning coulometric precision. We thank Mike Perry and James Saraidaridis at UTRC fora critical reading of the manuscript. We also thank David Kwabi,
Michael Gerhardt, Eugene Beh and Roy G. Gordon for helpful
discussions.
ORCID
Marc-Antoni Goulet https://orcid.org/0000-0002-9146-6759
Michael J. Aziz https://orcid.org/0000-0001-9657-9456

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Journal of The Electrochemical Society, 165 (7) A1466-A1477 (2018)
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Flow Battery Molecular Reactant Stability Determined by
Symmetric Cell Cycling Methods

Supporting Information

Marc-Antoni Goulet† and Michael J. Aziz†*
Harvard John A. Paulson School of Engineering and Applied Sciences, 29 Oxford
Street, Cambridge, Massachusetts 02138, USA*Corresponding Author Information:
Prof. Michael J. Aziz,
Address: Harvard John A. Paulson School of Engineering and Applied Sciences, 29 Oxford
Street, Cambridge, Massachusetts 02138, USA
Email:

maziz@harvard.edu

Tel: (617) 495-9884,
Fax: (617) 495-9837†

Active member of the Electrochemical Society (ECS)1

Methodology
Unless otherwise stated, the experiments in this supplementary information use the same
methodology as described in the main text. A few additional techniques and details pertaining
only to the supplementary data are added here.
Static cell
To avoid added noise due to pumping and the possibility of electrolyte becoming trapped in partsof the tubing junctions and reservoirs, and to understand the effects of high cycling rates; static
cells like the one shown in Fig. S1 were developed. For negolytes such as AQDS the current
collector was machined from solid niobium with a half inch diameter circular pocket deep enough(0.008”) to house one carbon paper electrode (0.01” thick each) for the limiting side and two
carbon paper electrodes (0.015” deep pocket) on the non-limiting side. The depth was chosen toprevent leakage of electrolytes while producing enough compression of the carbon paper for good
conductivity. The electrodes were placed within their respective current collectors and filled via
pipette with small, known quantities of electrolyte. A Nafion 117 membrane was sandwiched
between both electrodes and the cell was compressed together by vise clamp and sealed with PTFE
grease. Contact to each current collector was made via conductive copper tape.

Figure S1. Photograph of open niobium static cell with example carbon paper electrodes andmembrane
Chemical decomposition storage
Samples of both reduced and oxidized AQDS were prepared in a glovebox with less than 2 PPM
of oxygen content. The reduced form was made electrochemically by reducing a small amount of2

the initial 0.1 M AQDS in 1 M H2SO4 electrolyte within the glovebox with a full cell and thentransferring1 ml to an FEP vial with a tightly sealed cap. This FEP vial was then placed within acanning jar filled with tissues soaked in an aqueous sodium sulfite oxygen scavenger. The oxygen
scavenger was a 0.5 M Na2SO3 solution with 0.001 M KOH added to make it alkaline and avoid
hydrogen sulfide from forming. Roughly 20 ml of this oxygen scavenger was imbibed into thetissues which surrounded the FEP vial. The stable reduced form of BTMAP-viologen was used
as an oxygen indicator. For this purpose, a 20 µl aliquot of reduced 0.1 M BTMAP-viologen was
diluted into 1 ml of water and then placed within a two-dram glass vial with a loosely sealed cap.
With only 2 µM of reduced species this indicator vial contained only 2% of the oxidizable
molecules as compared to the sample under investigation and created a colorful blue solution thatwould turn colorless very rapidly (< 5 min) upon exposure to oxygen. This loosely capped
indicator vial was placed on the outside of the imbibed tissues within the same canning jar. The
jar was then closed tightly before being removed from the glovebox. The same procedure wasfollowed for a 1 ml sample of the initial oxidized AQDS electrolyte. The jars were then placed ina convection oven at 45 °C. After two weeks of storage, the jars were removed, and it was
confirmed that the BTMAP-viologen indicator vial was still blue in both cases before opening thejars. Once the jars were opened, both samples were removed and aerated to return the reduced
sample into oxidized AQDS. Aliquots were then taken for measurement and compared to thosetaken before storage.

NMR
NMR samples were prepared by diluting 100 µl of the AQDS electrolyte (initially 0.1 M AQDS
in 1 M H2SO4) in 700 µl of D2O containing roughly 10 mM of maleic acid as internal standard.
All samples were aerated after storage to ensure the electrolyte was converted into its oxidized
form. Testing was performed with a Varian INOVA 500 MHz spectrometer using 32 scans and a
10 second relaxation delay.

Results & discussion
Here we add results which provide further evidence for the claims stated in the associated paper.
Each section corresponds to the same section in the paper.
3Cycling pauses for investigating different states of charge of DHAQ
A representative example of DHAQ cycling in a symmetric cell is shown in Fig. S2. The data
show curvature whether plotted versus time or versus cycle number. On a semi-log plot instead,as in Fig. S2-c and Fig.S2-d, the data become more linear with a fitted R2 value > 0.999, which
implies that the capacity experiences exponential decay. When evaluating the instantaneous faderate this amounts to roughly 0.23 %/cycle and 8.8 %/day respectively, which is similar to what
was observed in other symmetric cell and full cell tests.

Figure S2. Symmetric flow cell cycling of 5 ml of 50% SOC 0.1 M DHAQ in 1 M KOH as
plotted a) capacity against cycle number, b) capacity against time, c) log(capacity) against cycle
number, and d) log(capacity) against time. Excellent linear fits in logarithmic data indicate
exponential decay.

4

The cycled DHAQ from Fig. 3 in the main text was analyzed via CV after cycling and comparedto the initial electrolyte. In this case, 100 µl of the initial 0.1 M electrolyte was added to 4.9 ml of
1 M KOH, after this supporting electrolyte was scanned at 100 mVsec-1. The initial electrolyte
was scanned within the potential limits seen in Fig. S3, after which point the electrode was thenrinsed, polished and rinsed again before repeating the procedure for the cycled electrolyte.

Figure S3. Background-corrected cyclic voltammograms of initial and cycled 0.1 M DHAQ in
1 M KOH vs. Ag/AgCl reference.

The background-corrected reduction peak currents of the initial material stand at around 40 µA,
whereas the cycled material produces only 24 µA. This indicates only roughly 60% of redox active
material remains in the electrolyte after cycling, which compares well with the 50% loss of
capacity observed during cycling.

Transient phenomena and strategies for measuring slow fade mechanisms
Using the niobium static cell depicted in Fig. S1, 30 µl of 50% SOC 0.7M H2AQDS/AQDS was
imbibed directly into the limiting electrode, while twice this volume was loaded into the nonlimiting electrode before sealing the cell in the glovebox. Similarly to the flow cell in the5associated paper, this static cell was cycled potentiostatically at ±200mV with a current cutoff of
5 mA. Due to the minute volumes involved, it was found that even small oxidative contributions
from the materials and ambient atmosphere could influence the starting SOC of the electrolytes.
Hence, a concentration higher than 0.1 M was used to preserve the initial SOC. The results fromthis test are displayed in Fig. S4 below.

Figure S4. Symmetric static cell cycling of 30ul of 0.7M AQDS.
The cell performed 1000 cycles in only 1.2 days, in which time the capacity of the cell decreased
by roughly 0.2%. Although slightly higher than in the flow cell test mentioned above, this could
be due to the higher concentration used in the static cell. The tiny volume used could also havebeen a factor due to a small amount of AQDS being lost by permeation into the membrane. A
small but noticeable initial transient is observed and if only the second half of the data is
considered, the rate of capacity fade falls to 0.11 ± 0.02%/day, which is within error of the flow
cell results. Considering that the rate of cycling in the flow cell is 10 times slower than in thisstatic cell, the similarity in fade rates therefore implies that AQDS is electrochemically stable, i.e.,
the number of times the molecule is reduced and oxidized has no effect on its lifetime. Although
the static cell shown here has the same membrane and electrolyte with an identical redox potential6

as in the corresponding flow cell, the data here show negligible average difference between the
reduction and oxidation capacity. This further demonstrates that this current efficiency is in partdue to the unavoidable background current of the potentiostat. This background current has a moresignificant effect on current efficiency in the flow cell due to the longer cycle time.
As shown in the main text, when the limiting electrolyte is held in a reduced H2AQDS
form, the temporal capacity fade rate is higher than the fade rate during cycling. This iscorroborated by NMR evidence of decomposition of the reduced H2AQDS form. After storing
both the oxidized and reduced forms of AQDS for 2 weeks at 45 °C, the reduced sample was reoxidized by aeration before being analyzed by NMR. In contrast to samples stored in the oxidized
form, Fig. S5 shows the appearance of new peaks in the NMR spectra of samples stored in reduced
form.

Figure S5. Proton NMR spectra of AQDS in D2O. a) baseline spectrum taken before storage, b)
spectrum of AQDS after storage in oxidized form for 2 weeks at 45°C and c) spectrum of AQDS
after storage in reduced form for 2 weeks at 45°C.
The appearance of side peaks in the sample stored in the reduced form indicates decompositionof the AQDS molecule. By integrating the AQDS peaks at 8.35 PPM and those between 8.00
and 8.10 PPM with respect to the maleic acid internal standard (not shown), it was also observedthat this sample suffered a loss of AQDS concentration on the order of 40 % with respect to theinitial and stored oxidized samples.
7

Material interactions and pH effects: Stability of Fe(CN)6
Capacity retention of a Fe(CN)6-bearing electrolyte was first studied using a static cell methodmuch like the niobium cell described above. It was found, however, that the alkaline electrolyte
would react with most metals either corroding them or depositing a blue layer onto them. Hencea similar design was used but with a flat pyro-sealed graphite current collector. Instead of amachined circular pocket, PVDF gaskets (0.02” and 0.03”) were used to contain two electrodes on
the CLS and three electrodes on the NCLS in this case. An example of results from this experimentis shown in Fig. S6 below.

Figure S6. Semi-log plot of unbalanced compositionally-symmetric static cell cycling of 60 µlof 50% SOC 0.1 M K4[Fe(CN)6] /K3[Fe(CN)6] in 1 M KOH
The extremely fast capacity fade of Fe(CN)6 seen in the static cell was unexpected and prompted
the use of flow cells to understand more about the process. The UV-vis spectrophotometry data
presented in the main text come from the cycling run portrayed in Fig. S7 in which 6 ml of 50%
SOC 0.1 M K4[Fe(CN)6] /K3[Fe(CN)6]in 1 M KOH was cycled from the start, with both sides
having equal capacity; hence the capacity fade starts immediately upon pumping of theelectrolytes.

8Figure S7. Semi-log plot of unbalanced compositionally-symmetric flow cell cycling of 6ml of
50% SOC 0.1 M K4[Fe(CN)6] /K3[Fe(CN)6]- in 1 M KOH.
The same starting electrolyte was also stored with and without carbon paper in clear centrifuge
tubes during the same time interval. Their colors were compared (Fig. S8) and UV-vis absorption
spectra were compared as well.

Figure S8. Photograph of initial 50% SOC 0.1 M K4[Fe(CN)6] /K3[Fe(CN)6]in 1 M KOH
electrolyte before and after storage of solutions with and without carbon paper.

9

The spontaneous reduction of ferricyanide to ferrocyanide, which is catalyzed by carbon paper, isdue to the redox potential of the molecule being above the stability limit of water at pH 14.
Therefore, this reduction reaction should be decreased at lower pH, as demonstrated in Fig. S9below for cycling at pH 12.

Figure S9. Semi-log plot of unbalanced compositionally-symmetric cell cycling of 0.1 M
K4[Fe(CN)6] /K3[Fe(CN)6]in 0.01 M KOH (pH 12).
The instantaneous fade rate for this cell is roughly 7 %/day, which is considerably lower than thesame symmetric cell at pH 14 (on the order of 50%/day). As explained in the main text, thiscapacity fade is due to a recoverable spontaneous reduction of the ferricyanide, and hence thissuggests that, when operating a flow battery with a Fe(CN)6 posolyte, less current efficiency will
be sacrificed at lower pH.

Recommendations and implications for the field
To demonstrate the effect of increased resistance on measured capacity, a 0.13 Ω resistor wasplaced in series with the cell during cycling. The impedance of the cell with and without theresistor was checked several times throughout the experiment, as shown in Fig. S10.
10

0.2

-Im(Z) (Ohm)With resistor
Without resistor

0.1

0.0
0.40

0.45

0.50

0.55

0.60

0.65

0.70

0.75

0.80

Re(Z) (Ohm)Figure S10. Potentiostatic EIS measurements of cell between symmetric cell cycling segmentsof 0.1 M DHAQ in 1 M KOH showing effect of resistor added in series. These impedance
spectra were measured after each potentiostatic cycling segment from Fig. 11 in the main text.
Overlapping data sets, two separate EIS spectra with resistor (black) and two spectra without
resistor (red), indicate little variability in cell resistance throughout experiment.

11