Article

Alkaline Quinone Flow Battery with Long
Lifetime at pH 12
David G. Kwabi, Kaixiang Lin,
Yunlong Ji, ..., Alán
Aspuru-Guzik, Roy G. Gordon,
Michael J. Aziz
maziz@harvard.edu

HIGHLIGHTS
Near-neutral chemistry forrechargeable, aqueous-soluble,organic redox-flow batteries
Capacity fade rate of 0.01%/
day, the longest lifetime for aquinone-based electrolyte
Meets technical criteria forcommercialization of redox-flowbatteries

This work demonstrates a new, organic redox-flow battery (RFB) that outlives its
predecessors, offering the longest-lived high-performance organic flow battery todate. It appears to be the first aqueous-soluble organic RFB chemistry to meet all
the technical criteria for commercialization. The potential low reactant andmembrane costs of this chemistry offer the potential for RFBs of this type to be
used cost effectively at the gigawatt scale in order to enable massive penetrationof intermittent renewable electricity.

Kwabi et al., Joule 2, 1894–1906
September 19, 2018 ª 2018 The Authors.
Published by Elsevier Inc.
https://doi.org/10.1016/j.joule.2018.07.005

Article

Alkaline Quinone Flow Battery
with Long Lifetime at pH 12
David G. Kwabi,1,4 Kaixiang Lin,2,4 Yunlong Ji,2 Emily F. Kerr,2 Marc-Antoni Goulet,1
Diana De Porcellinis,1 Daniel P. Tabor,2 Daniel A. Pollack,3 Alán Aspuru-Guzik,2 Roy G. Gordon,1,2
and Michael J. Aziz1,5,*

SUMMARY

Context & Scale

We demonstrate a long-lifetime, aqueous redox-flow battery that can operate
at a pH as low as 12 while maintaining an open-circuit voltage of over 1 V. We
functionalized 2,6-dihydroxyanthraquinone (2,6-DHAQ) with highly alkali-soluble carboxylate terminal groups. The resulting negative electrolyte material
4,40 -((9,10-anthraquinone-2,6-diyl)dioxy)dibutyrate (2,6-DBEAQ) was six times
more soluble than 2,6-DHAQ at pH 12. Symmetric cell cycling with 2,6-DBEAQ
on both sides of the cell demonstrates a capacity fade rate of <0.01%/day and<0.001%/cycle. By pairing 2,6-DBEAQ with a potassium ferri-/ferrocyanide
positive electrolyte and utilizing a non-fluorinated membrane, this near-neutral
flow battery shows a capacity fade rate that is the lowest of any quinone andrivals the lowest ever reported for any flow battery in the absence of rebalancing processes. This result adds the important attribute of long calendar life toquinone-based redox-flow batteries, which may enable massive penetrationof intermittent renewable electricity.

Electricity generation fromrenewable sources such as solar
and wind can eliminate fossil-fuelbased systems if their intermittent
output can be regulated using
safe, cost-effective energy
storage. Aqueous-soluble organic
redox-flow batteries (RFBs) are apotentially safer, less expensivealternative to lithium ion batteries
and vanadium flow batteries forlong-discharge duration storage.
We demonstrate a long-lifetime,organic-molecule-based RFB thatcan operate in weak alkalineconditions while maintaining an
open-circuit voltage of over 1 V.
The negative electrolytecomprises an anthraquinone
functionalized with solubilizing
carboxylate groups and thepositive electrolyte comprises thefood additive potassium
ferrocyanide. The RFB can
operate with an inexpensive
membrane and shows a capacity
fade rate that rivals the lowest
ever reported for any RFB in theabsence of rebalancing
processes. This result adds theimportant attribute of long
calendar life to quinone-based
RFBs.

INTRODUCTION
The cost of wind and photovoltaic electricity has dropped so much that one of thegreatest technical barriers to their widespread substitution for fossil electricity istheir intermittency. Cost-effective, safe, and scalable stationary electricity storage
could solve this problem. Traditional enclosed batteries such as lithium ion batteries
are common, but they cannot cost effectively store enough energy for the longdischarge durations at rated power that appear to be necessary for regulating
the intermittency of renewables. Redox-flow batteries (RFBs) comprise one classof energy storage systems particularly suited for long-duration discharge because
energy-storing species are held in liquid form in external tanks that are separated
from the power generation stack. Consequently, energy and power ratings can bescaled independently, and discharge durations (energy:power ratios) of several
hours to days can be achieved at the rated power.
The RFBs farthest along the commercial pathway are based on charge storageof vanadium ions in acidic electrolytes, but they have limited potential for
widespread adoption due to the high cost and low earth abundance of vanadium.1
As a consequence, RFBs based on redox-active organic and organometallic
reactants2–6 are receiving much research interest, as they are potentially muchless expensive than their vanadium-based counterparts; in addition, organic
reactants can be chemically modified through selective functionalization in orderto improve aspects of performance such as voltage, rate capability, and energy
density.

1894 Joule 2, 1894–1906, September 19, 2018 ª 2018 The Authors. Published by Elsevier Inc.
This is an open access article under the CC BY-NC-ND license (http://creativecommons.org/licenses/by-nc-nd/4.0/).

One drawback of many reported organic RFBs, however, is their low chemical stability because organic molecules are susceptible to degradation reactions such as
nucleophilic substitution, gem-diol formation, and self-polymerization. Several
organic and organometallic RFBs have been reported in the past 5 years, operatingin highly acidic, highly alkaline, or neutral conditions, based on quinone,2,4,7viologen,8–11 ferrocene,8 alloxazine,5 and nitroxide radical10,12–14 motifs. Most ofthem, however, experience high temporal capacity fade rates of 0.1%–3.5%/day,which limits their long-term use over many years. Such relatively short lifetimes
render most of these chemistries unsuitable for commercialization, while motivating
further research into the development of organic molecules with high stability inboth oxidized and reduced redox states.
Here, we report a new negative electrolyte (negolyte) molecule synthesized from ourpreviously reported DHAQ chemistry that exhibits record high chemical stability, leading to the lowest temporal capacity fade among quinone-based flow batteries, on theorder of 5%/year. This is among the lowest reported for any RFB in the absence of rebalancing procedures. The approach involves the introduction of ether-linked alkyl
chains with solubilizing carboxylate functional groups onto an anthraquinone core
and affords high chemical stability and high solubility (0.6 M at pH 12; 1.1 M at pH14) in alkaline electrolyte (Table S1). We hypothesized that solubility would be enhancedby frustrating crystallization if we attached solubilizing groups via short hydrophobic
chains that were not so long that their hydrophobicity dominates. The solubility
enhancement reduces the concentration of supporting electrolyte by two ordersof magnitude without compromising the ionic conductivity of the system. Pairing
2,6-DBEAQ with a ferro-/ferricyanide-based positive electrolyte (posolyte) results in abattery with an open-circuit voltage of 1.05 V and theoretical volumetric energy
density of 12 Wh/L at pH 14 and 17 Wh/L at pH 12 (Table S2). We further showthat this O-alkylated anthraquinone can be used in a flow battery at pH 12 with an inexpensive hydrocarbon-based membrane that features exceptionally high permselectivity
and low permeability of both the alkylated anthraquinone and ferricyanide species,
affording century-scale timescales for reactant crossover. These results highlight theimportance of synthetic approaches to increase the chemical stability of active species
and constitute a leap forward in realizing high-performance aqueous organic RFBs thatexhibit long cycle and calendar life at low capital cost.

RESULTS AND DISCUSSION
Synthesis and Cyclic Voltammetry Study
Figure 1 illustrates the synthetic route, chemical structures, and cyclic voltammograms(CVs) of various isomers of DBEAQ. All DBEAQ isomers can be synthesized fairly simply
by an O-alkylation reaction followed by hydrolysis of the ester to afford the highly alkalisoluble carboxylic acid terminal groups (Figure 1A). Since the second hydrolysis reactionis almost 100% in yield and the by-product is a volatile organic, i.e., methanol, we expectthat for the industrial process, the negolyte can be prepared directly after the hydrolysis
step and the removal of methanol without the additional cost incurred by two unnecessary acid-base treatments, namely, collection of the hydrolyzed DBEAQ product via
acidification for characterization and deprotonation of DBEAQ with a designatedamount of base for electrolyte preparation.

1Harvard School of Engineering and Applied

Sciences, Cambridge, MA 02138, USA
2Department of Chemistry and Chemical Biology,

Harvard University, Cambridge, MA 02138, USA
3Department of Physics, Harvard University,

Cambridge, MA 02138, USA

After the functionalization, the resulting DBEAQ isomers (Figures 1B and S1–S3) showeda significant improvement in solubility versus their hydroxyl counterparts at both pH 14
and pH 12. For instance, the room-temperature solubility of 2,6-DBEAQ exceeded 1 M
in pH 14 and 0.5 M in pH 12 KOH solution (Table S1), as opposed to 0.6 M4 and 0.1 M,

4These authors contributed equally
5Lead Contact*Correspondence: maziz@harvard.edu
https://doi.org/10.1016/j.joule.2018.07.005

Joule 2, 1894–1906, September 19, 2018 1895

Figure 1. Synthesis and Initial Characterization of DBEAQ(A) Synthetic approach to functionalize hydroxyanthraquinone with tethered highly alkali-soluble
carboxylic groups.
(B) Chemical structures of 1,2-, 2,6-, and 1,8-DBEAQ.
(C) Cyclic voltammograms of 1,2- (blue), 2,6-(gray), and 1,8- (green) DBEAQ. The redox potential vsSHE of each isomer is indicated.
See also Figures S1–S3.

for 2,6-DHAQ under similar conditions, respectively. Several avenues exist for further
increasing reactant solubility and the corresponding energy density, such as the useof mixed cations in the electrolytes 15 or mixtures of different DBEAQ isomers or
different molecules with nearly the same reduction potential.
All the DBEAQ isomers showed very similar redox potentials between 510 and540 mV versus standard hydrogen electrode (SHE) (Figure 1C), which would yieldbattery voltages above 1 V versus ferri-/ferrocyanide posolyte. Despite the increase
in reduction potential compared with 2,6-DHAQ (680 mV versus SHE), their CVs
exhibited high reversibility with redox peaks 40 mV apart, much smaller than the90 mV value from 2,6-DHAQ.4 A Levich analysis was used to obtain the diffusion
coefficient of the oxidized form of 2,6-DBEAQ (1.58 3 106 cm2/s; Figure S4), which
was then used in a CV simulation of its redox kinetics (Figure S5). The results show, inagreement with an analogous study of the 2,6-DHAQ CV, that the peak shape andseparation are consistent with two one-electron reduction steps at different potentials, E1 and E2, whose values are modulated by the energetics of semiquinone
reduction (see Figure S6 and associated discussion in the Supplemental Information). In comparison with 2,6-DHAQ, which exhibited an E1-E2 separation of
60 mV,4 the 2,6-DBEAQ CV exhibits a much smaller E1-E2 separation of 6 mV,
implying that 2,6-DHAQ is more thermodynamically susceptible to the formationof semiquinone radicals than 2,6-DBEAQ.
Chemical and Electrochemical Stability Study
We first demonstrate that 2,6-DBEAQ has superior chemical stability over other quinones reported to date, notably including 2,6-DHAQ. As a direct probe of the effectof O-alkylation of anthraquinone on its chemical stability, 2,6-DBEAQ was cycled
potentiostatically in a volumetrically unbalanced compositionally symmetric cell1896 Joule 2, 1894–1906, September 19, 2018

Figure 2. Symmetric Cell Cycling of 2,6-DBEAQ at pH 14(A) Cell schematic for unbalanced compositionally symmetric cell cycling and full cell cycling.
(B and C) (B) Unbalanced compositionally symmetric cell cycling of 0.10 M 2,6-DBEAQ and (C)0.65 M 2,6-DBEAQ, showing capacity as a function of time and temporal capacity fade rates fromlinear fits of the last 5 days of cycling. The capacity-limiting side was 5 mL 2,6-DBEAQ, while thenon-capacity-limiting side was 10 mL 2,6-DBEAQ, both at pH 14. Capacities were obtained by fullpotentiostatic reduction and oxidation at G0.2 V of the capacity-limiting side; the potential was
switched when the magnitude of the current density decayed to 2 mA/cm 2 . Note that both y axis
scales represent about 2% of the capacity of the capacity-limiting side.
See also Figure S7.

configuration. The symmetric cell configuration, in which the two sides have thesame electrolyte composition, is a simple and direct probe of chemical and electrochemical stability of RFB reactants and is described in full in a separate report.16 Anyobserved capacity fade is directly related to deactivation of the reactant on the capacity-limiting side in either its oxidized or reduced state because (1) there is negligible reactant crossover with symmetric compositions, (2) changes to membrane
resistance do not result in temporal capacity variations in potentiostatic cycling,and (3) the unbalanced volumes permit the capacity-limiting side to be taken to its
limiting states of charge (SOCs) despite potential side reactions.
Figure 2 shows a schematic of the cell setup (Figure 2A) and the results of unbalanced compositionally symmetric cell cycling for 2,6-DBEAQ at 0.10 and 0.65 M
for durations of 13 and 26 days, respectively (Figures 2B and 2C). In both cases,
2,6-DBEAQ exhibited temporal capacity fade rates of <0.01%/day or <3.0%/year,
which suggests a loss mechanism that is first order in 2,6-DBEAQ concentration.
In contrast, 2,6-DHAQ cycled in a similar cell configuration demonstrated a muchhigher temporal fade rate of 5%/day (Figure S7). Temporal capacity fade rates
for neutral organic/organometallic RFBs have been summarized elsewhere.8 Most
systems show fade rates in the range of 0.1%–3.5%/day. The temporal fade rateof 2,6-DBEAQ is the lowest ever reported for a quinone-based electrolyte, withthe lowest previously reported being for 9,10-anthraquinone-2,6-disulfonate(AQDS) at 0.1%–0.2%/day.17,18 It is on par with that of a recently reported

Joule 2, 1894–1906, September 19, 2018 1897

viologen-based flow battery,8 which exhibited the highest capacity retention rate forany flow battery in the absence of rebalancing processes, with a temporal capacity
fade rate of <0.01%/day in symmetric cell testing.16
The origin of the higher chemical stability of 2,6-DBEAQ compared with 2,6-DHAQ
is not completely understood; however, we highlight here a few key distinctions that
may be responsible. Notably, it is the reduced form of 2,6-DHAQ that has beenshown in a previous study to be involved in the loss of redox activity, whereas the
reduced form of 2,6-DBEAQ is shown here to be quite stable even at temperatures
up to 95 C (see Supplemental Information). Our first observation is that, because
the redox potential of 2,6-DBEAQ is 200 mV higher than that of 2,6-DHAQ,
2,6-DBEAQ should be more stable thermodynamically in its reduced form. Secondly, at high pH, we expect both the reduced hydroquinone core and solubilizing
groups to be deprotonated and therefore negatively charged in both molecules.
The closer proximity of the negatively charged deprotonated hydroxyl groups in
2,6-DHAQ should lead to larger intramolecular Coulombic repulsion forces, which
may contribute to this destabilization of the reduced form of the molecule. Finally,
the greater susceptibility to the formation of semiquinone radicals in 2,6-DHAQ,
as discussed above, may also be involved in its decreased stability.
The alkyl chain functionalization, while drastically improving lifetime, does not avoid
decomposition altogether. By performing elevated temperature chemical stability
studies, we have identified that the cleavage of g-hydroxybutyrate is involved in thedecomposition of the oxidized form of 2,6-DBEAQ (Figures S8 and S9), and we havecharacterized the time course of g-hydroxybutyrate cleavage at pH 12 and pH 14 andat 0.1 M and 0.5 M concentration (Figures S10 and S11). The results suggest that thehalf-life of 2,6-DBEAQ at room temperature, pH 14, and 0.1 M concentration in theoxidized form is on the order of 5 years, with substantially slower decomposition in
the reduced form relative to the oxidized form and at pH 12 relative to pH 14; these observations are consistent with long RFB lifetime at typical operating cell conditions.
Membrane and Full Cell Studies
Ion-conductive membranes play a critical role in RFB systems in governing the transport of counter ions across the membranes (ionic conductivity) and the simultaneoustransport of redox-active species (membrane crossover). In a compositionally asymmetric cell, e.g., 2,6-DBEAQ-ferrocyanide system, capacity fade rate can be greatly
exacerbated by the irrecoverable crossover of the reactant to the other side of theelectrolyte.19 To limit capacity fade by this mechanism, while retaining high permselectivity, we surveyed a range of cation-exchange membranes, including the industry
standard Nafion-based perfluorosulfonic acid (PFSA) membranes, by characterizing
both their K+ conductivity and permeability of 2,6-DBEAQ and ferricyanide (morepermeable than ferrocyanide) (see Experimental Procedures for the customized
lab setup). E-600 series membranes, which comprise a non-fluorinated, sulfonatedpolyaryletherketone-copolymer backbone, delivered the best performance. The
membrane displayed a low area specific resistance (ASR) of 1 U∙cm2 in 1 M K+
solution, which is comparable with Nafion 212. It also showed an extremely
low 2,6-DBEAQ and ferricyanide permeability of 5.26 3 1013 cm2/s and 4.4 3
1012 cm2/s, respectively (Figure S12), which are at least an order of magnitude
lower than Nafion 212 systems. To further prove the low permeability of redox-active
species, we constructed a low concentration 2,6-DBEAQ-ferri-/ferrocyanide full cellat pH 14, using a Fumasep E-620 (K) membrane. Over a period of galvanostatic
cycling testing for 4 days (880 cycles), the cell showed immeasurably low capacity
fade and a current efficiency around 99.94% (Figure S13). This test is prima facie

1898 Joule 2, 1894–1906, September 19, 2018

evidence that the low 2,6-DBEAQ permeabilities measured ex situ using thenutating table setup translate to negligible crossover of 2,6-DBEAQ in an operatingcell and that additional contributions to crossover that are not present in thenutating table tests, such as electromigration and pressure from active pumpingof electrolytes,20 are unimportant. These low permeabilities, together with the factthat Fumasep E-600 series membranes do not contain fluorine, imply that theiruse in a DBEAQ-ferri-/ferrocyanide full cell affords a potentially robust configuration
for a long-lasting RFB with low power cost (in $/kW) for fully installed systems (see A
Note on DBEAQ/DHAQ Permeability in Supplemental Information).
To reduce the corrosivity of the system and the ferricyanide decomposition rate,
which is exacerbated at high pH,21,22 we performed the full cell tests of 2,6-DBEAQ
at a more moderate electrolyte condition of pH 12. Symmetric cell cycling of 1,2(Figure S14A) and 1,8-DBEAQ (Figure S14B) at pH 12 showed higher temporalfade rates—greater than 0.1% and 10%/day, respectively—than 2,6-DBEAQ at thesame pH (<0.01%/day, Figure S15) and were therefore not tested in full cell experiments. Theoretical calculations suggest that these isomers are more thermodynamically susceptible to hydroxide or water-induced g-hydroxybutyrate cleavage than
2,6-DBEAQ (see Table S4). Given the 0.6 M solubility measured for 2,6-DBEAQ at
pH 12 (Table S2), subsequent cell tests were performed at 0.5 M in order to examine
the performance of a full cell with reasonable energy density. From the Pourbaix diagram of 2,6-DBEAQ (Figure S16A), its reduction potential becomes pH independent above pH 11.5 and is not expected to change during cell cycling. Polarizationand capacity utilization measurements with a negolyte containing 0.5 M 2,6-DBEAQ
are shown in Figure 3. Polarization studies (Figure 3A) at room temperature showeda near-linear relationship between current density and voltage at currents close to
the open-circuit voltage (OCV), which increased from 0.97 V at 10% SOC to 1.12 V
at 100% SOC (Figure 3B). Between 80% and 90% of the polarization ASR is accounted for by the high-frequency resistance measured using electrochemical
impedance spectroscopy (EIS), which largely reflects membrane resistance.
A peak galvanic power density of 0.24 W/cm2 was realized at 100% SOC (Figure S17).
This power density is about half of that previously reported in a 2,6-DHAQ-ferrocyanide cell,4 owing to the higher OCV of the latter (1.20 V as opposed to 1.05 V at 50%
SOC) and smaller ASR (0.858 U cm2 as opposed to 1.2 U cm2 at 50% SOC). Whenvoltage and current have a linear, ohmic relationship, the peak galvanic power den2
sity is given by pmax = VOC=r, where VOC is the open-circuit potential and r is the ASR.
The low permeability of the Fumasep membrane to the fastest-crossing species(i.e., 200 years required for 50% loss through crossover of ferricyanide) permits, inprinciple, a 4-fold reduction in the membrane thickness, which would raise the power
density significantly.
In order to avoid temporal variations in accessible capacity during full cell cycling
caused by changes in membrane resistance,16 each galvanostatic half cycle wasfinished with a potential hold at the potential limit (1.4 V after charge, 0.6 V after
discharge) until the magnitude of the current density fell below 2 mA/cm2. Over a
5-day test period, a capacity fade rate of 0.05%/day or 0.001%/cycle was observed(Figure 4A). A parallel cycling test was performed in which a potentiostatic chargedischarge cycle was executed after every 20 galvanostatic cycles; a capacity faderate of 0.04%/day was observed in that case (Figure 4B). It has been shown that capacity retention rates using both cycling protocols yield virtually identical results.16
These full cell measurements, however, showed roughly 4-fold increase in capacity
fade compared with the <0.01%/day observed during symmetric cycling tests

Joule 2, 1894–1906, September 19, 2018 1899

Figure 3. Polarization Measurements of 2,6-DBEAQ Full Cell at pH 12(A) Cell voltage versus discharge current density at room temperature at 10%, 30%, 50%, 70%, 90%,
and 100% SOC. Electrolytes comprised 5 mL of 0.5 M 2,6-DBEAQ (negolyte) at pH 12 (10 mM KOH)
and 38 mL of 0.3 M potassium ferrocyanide and 0.1 M potassium ferricyanide (posolyte) at pH 12.
(B) OCV, high-frequency, and polarization ASR versus SOC.
(C) Galvanostatic charge and discharge curves from 25 to 250 mA/cm 2 . The vertical dashed lines
indicate the maximum capacity realized with potentiostatic charge and discharge at the voltage
cutoffs (1.4 and 0.6 V, respectively), as well as the theoretical capacity.
(D) Coulombic efficiency, round-trip energy efficiency, and capacity utilization as a percentage ofpotentiostatic capacity versus current density.

(Figures 2B and 2C), suggesting additional capacity fade mechanisms not observed
during symmetric cell studies. In-depth chemical and electrochemical analysis (Figure S18) was performed to probe the chemical decomposition and crossover of
DBEAQ from the capacity-limiting side. Electrolytes in a cell compositionally identical to that in Figure 4A but cycled for 11 days were subjected to nuclear magnetic
resonance (NMR) (Figure S18A) and CV (Figure S18B) analysis. From these results, noevidence of DBEAQ decomposition and crossover was found after examining thenegolyte and posolyte before and after cycling. Based on the detection limit ofNMR (0.1 mM DBEAQ) and CV techniques under the experimental conditions chosen, the upper limit of the capacity fade rate caused by 2,6-DBEAQ decomposition
and/or membrane crossover was 0.01%/day, similar to our symmetric cell study.
We therefore hypothesize that other capacity fade mechanisms, such as precipitation of 2,6-DBEAQ in the posolyte after crossing over, might be operative but
untraceable by NMR and CV techniques due to the slow capacity fade rate, whichcorresponds to a total loss of <0.5% of 2,6-DBEAQ over a 6-day testing period.
One other such mechanism might be leakage of the negolyte due to poorer adhesion between the thinner Fumasep membrane and the Viton gasket in the full cellthan between Nafion N117 and the gasket in the symmetric cell. Indeed, the total
capacity fade in Figure 4B corresponds to a total loss over the entire 6-day cycling
period of 10 mL of negolyte volume, which is roughly one-fifth of a droplet.
When translated to an equivalent current density (0.5 mA/cm2), this negolyte loss
rate is well within expectation for seepage of the electrolyte into spaces between

1900 Joule 2, 1894–1906, September 19, 2018

Figure 4. Extended Full Cell Cycling of 2,6-DBEAQ at 0.5 M and pH 12(A) Current efficiency (squares) and charge (upward-pointing triangles) and discharge (downwardpointing triangles) capacity versus time and cycle number for a negolyte-limited 2,6-DBEAQFe(CN) 6 cell. The cell was cycled galvanostatically at 100 mA/cm 2 between 1.4 and 0.6 V, andeach half cycle ended with a potentiostatic hold until the magnitude of the current density fell
below 2 mA/cm 2 . The negolyte comprised 5 mL of 0.5 M 2,6-DBEAQ at pH 12 while the posolytecomprised 30 mL of 0.3 M potassium ferrocyanide and 0.1 M potassium ferricyanide at pH 12.
(B) Evolution of charge (upward-pointing triangles) and discharge (downward-pointing triangles)
capacity for extended cell cycling at 100 mA/cm 2 . After every 20 th galvanostatic cycle, apotentiostatic charge-discharge cycle was performed, the discharge capacity of which is shown(open circles). Because the concentration of oxidized 2,6-DBEAQ is at its maximum after eachpotentiostatic discharge, the capacity of the galvanostatic charge step immediately following
potentiostatic charge-discharge is higher than in subsequent cycles. This particular cell showed
only 65% of its theoretical capacity because a large fraction of salt was left over from the synthesisof 2,6-DBEAQ.

the gaskets and/or interface between the membrane and gaskets, compared with ananalogous leak rate estimated from a previous study of capacity fade in an anthraquinone-based flow battery with this cell architecture (0.09–0.12 mA/cm2).19
Quinones for aqueous flow batteries have been the subject of intensive research
since their recent debut in this application. This work demonstrates that thesimple functionalization with bulky charged groups can greatly extend the calendar
life and cycle life; raise the solubility, leading to a theoretical energy densityof 17 Wh/L with possible routes to further increases; operate with fluoropolymerfree membranes at pH as low as 12 while maintaining the voltage against
Fe(CN)63/4 above 1 V; and perform with peak power density of 0.24 W/cm2 atroom temperature with possible routes to further increases. These properties appear
to meet all of the technical requirements commonly understood to be necessary for
aqueous RFB commercialization.
At the large production volumes necessary for gigawatt-scale grid storage, it ispossible that DBEAQ may approach the cost of other functionalized anthraquinones, such as DHAQ and AQDS, the latter of which was recently estimated to be
between $0.92/kg and $3.92/kg.23 Using $2.40/kg as a mid-range cost, this results
in a capital cost of $17.60/kWh for the 2-electron-accepting DBEAQ in the negolyte
in a 1.05 V battery. Similarly, we estimate the industrial-scale cost of potassium ferrocyanide to be around $2.15/kg,24 which amounts to a capital cost of about $32/kWhfor the posolyte (see Estimated Cost of Electrolytes in the Supplemental Information). At $50/kWh, the total capital cost of these electrolytes stands at less than
one-third of the cost of an all-vanadium electrolyte, for which the vanadium alone
is currently priced at $160/kWh. Using a fluorine-free membrane such as Fumasep
E620 (K) should result in a product of membrane cost (<$25/m2 at large production

Joule 2, 1894–1906, September 19, 2018 1901

volumes) and ASR (<1.5 U cm2) that is below $5/mU, which adds negligible cost to
the system.25 Thus, if the DBEAQ production cost at scale turns out to be nearly as
low as for AQDS and DHAQ,23,26 then this chemistry may have serious commercial
potential, and DBEAQ-based flow batteries may be instrumental in accelerating
the penetration of wind and photovoltaic electricity.

EXPERIMENTAL PROCEDURES
Synthesis and Chemical Characterization
Synthesis of 4,40 -((9,10-Anthraquinone-2,6-diyl)dioxy)dibutyric Acid
2,6-DHAQ was purchased from AK Scientific. Methyl 4-bromobutyrate was purchased from VWR. All other chemicals were purchased from Sigma-Aldrich. All
chemicals were used as received unless specified otherwise.
Dimethyl 4,40 -((9,10-anthraquinone-2,6-diyl)dioxy)dibutyrate (1). 2,6-DHAQ was
first converted to its dipotassium salt (2,6-DHAQK2) by adding 2,6-DHAQ (5 g,
20.8 mmol) to a 250 mL oven-dried flask of dimethylformamide (250 mL). Under
vigorous stirring, potassium ethoxide (6.1 g, 72.9 mmol) was added. The mixture
solution was stirred at room temperature for 15 min. For the O-alkylation reaction,
2,6-DHAQK2 (6.5 g, 20.8 mmol) was mixed with anhydrous K2CO3 (14.3 g,
104 mmol) and methyl 4-bromobutyrate (12.4 mL, 104 mmol). The reaction mixture
was then heated to 95 C overnight. After cooling to 0 C, deionized (DI) water(150 mL) was added to the mixture to dissolve inorganic salt and to precipitate
the ester precursor of DBEAQ. The precipitate was vacuum filtered and washedthoroughly with DI water (50 mL). The product was analyzed by 1H NMR and used
for the next step reaction without further purification. Final yield, 87%.
4,40 -((9,10-Anthraquinone-2,6-diyl)dioxy)dibutyric Acid (2). The ester precursor of
DBEAQ (1 g, 2.23 mmol) was added along with KOH (0.52 g, 9 mmol) to a flask filledwith a water-isopropanol mixture (2:1 v/v, 60 mL). The solution was stirred vigorously
and heated to 60 C for 12 hr. During the reaction, all solids dissolved, and the solution became a dark red color. After the reaction, the solution was transferred to alarger 500 mL flask and diluted with DI water (200 mL). Glacial acetic acid was added
until the pH of the solution dropped to 4. The mixture was stirred vigorously for
1 hr, followed by vacuum filtration and thorough washing with DI water (100 mL).
The product was vacuum dried, analyzed by 1H NMR, and used for electrochemicalmeasurement without further purification. Final yield, 99%.
The 1H NMR spectrum of 4,40 -((9,10-anthraquinone-2,6-diyl)dioxy)dibutyricacid is shown in Figure S1. 1,8- and 1,2-isomers were synthesized similarly, using
1,2-DHAQ and 1,8-DHAQ as precursors. Potassium tert-butoxide (8.16 g,
72.9 mmol) replaced potassium ethoxide to deprotonate the 1,2-DHAQ and 1,8DHAQ isomers because of the higher pKa of the hydroxyl groups of these isomers.
1
H NMR spectra of the final products are shown in Figures S2 and S3, respectively.
Electrochemical Characterization
Cyclic Voltammetry and Rotating Disk Electrode Measurements
Glassy carbon was used as the working electrode for all three-electrode CV tests.
Rotating disk electrode experiments were conducted using a Pine Instruments
Modulated Speed Rotator AFMSRCE equipped with a 5 mm diameter glassy carbon
working electrode, a Ag/AgCl reference electrode (BASi, pre-soaked in 3 M NaCl solution), and a graphite counter electrode. The electrode was rotated at a specificspeed while the voltage was swept linearly from 0.4 to 0.11 V versus Ag/AgCl(Figure S4A). The diffusion coefficient of the oxidized form of 2,6-DBEAQ was

1902 Joule 2, 1894–1906, September 19, 2018

calculated using the Levich equation, which relates the mass-transport-limited current to the number of electrons transferred (n), the area of the electrode (A), andthe concentration of redox-active species in the electrolyte (C), by plotting the
mass-transport-limited current against the square root of the rotation rate (Figure S4B) with the following parameters: n = 2, F = 96,485 C/mol, A = 0.196 cm2,
C = 5 mM, kinematic viscosity of 1 M KOH n = 1.08 3 106 m2/s. The resulting valueof the diffusion coefficient for the oxidized form of 2,6-DBEAQ is 1.58 3 106 cm2/s.
Symmetric and Full Cell Measurements
Flow battery experiments were constructed with cell hardware from Fuel Cell Tech
(Albuquerque, NM) assembled into a zero-gap flow cell configuration, similar to aprevious report.4 Pyrosealed POCO graphite flow plates with serpentine flow patterns were used for both electrodes. Each electrode comprised a 5 cm2 geometricsurface area covered by a stack of three or four sheets of Sigracet SGL 39AA porous
carbon paper pre-baked in air overnight at 400 C.
For symmetric cell tests, a sheet of Nafion 117 membrane soaked overnight in 1 M
KOH served as the ion-selective membrane between the carbon electrodes,
whereas for full cell tests, a Fumasep E620 (K) membrane was used. The outerportion of the space between the electrodes was gasketed by Viton sheets withthe area over the electrodes cut out. Torque applied during cell assembly was 60
lb-in (6.78 Nm) on each of 8 bolts. The electrolytes were fed into the cell throughfluorinated ethylene propylene tubing at a rate of 60 mL/min, controlled by ColeParmer Masterflex L/S peristaltic pumps. All cells were run inside a nitrogen-filled
glove box with an O2 partial pressure of about 1–2 ppm. Cell polarization measurements, impedance spectroscopy, and charge-discharge cycling were performed using a Biologic VSP 300 potentiostat.
Potentiostatic symmetric cell cycling was performed according to a reported procedure16 in order to assess the temporal chemical stability of DBEAQ, independently of any variations in apparent capacity as a result of changes to membrane
resistance, and in a configuration where there is negligible water or reactant crossover due to the same DBEAQ concentration being used in both electrolytes. 7.5 mLof 2,6-DBEAQ was fully reduced (i.e., charged to 100% SOC) against a posolyte ofexcess potassium ferrocyanide and then mixed with 7.5 mL of oxidized 2,6-DBEAQ(0% SOC) to afford a 2,6-DBEAQ electrolyte at 50% SOC. 5 mL of the resulting
electrolyte was used as the capacity-limiting side of a volumetrically unbalanced
compositionally symmetric cell, while the remaining 10 mL was used as the
non-capacity-limiting side, i.e., with twice the nominal capacity of the capacitylimiting side. The capacity-limiting side was then cycled potentiostatically between
0% and 100% SOC with potential limits of G0.2 V versus the non-capacity-limiting
side, switched when the absolute value of the current density decayed to 2 mA/cm2,
which is 33 higher than the background current.
Full cell cycling (i.e., with a ferro-/ferricyanide-based posolyte) was performed withthe same flow cell hardware but with a Fumasep E620 (K) membrane due to its highconductivity and low permeability of DBEAQ and ferricyanide compared with othermembranes (see Supplemental Information). For studies at pH 14, the posolyte volume was 30 mL, and its composition, when assembled, was 0.20 M potassium ferrocyanide, 0.08 M potassium ferricyanide, and 1 M KOH. The negolyte was preparedby dissolving 0.1 M 2,6-DBEAQ in its oxidized form in 1.2 M KOH solution, resulting
in 0.1 M 2,6-DBEAQ, 1.2 M K+, and 1 M OH electrolyte solution. For studies at
pH 12, both electrolytes comprised 10 mM OH; the pH was checked both with aJoule 2, 1894–1906, September 19, 2018 1903

pH meter and pH paper and adjusted where necessary. For all full cell studies, thenegolyte was assembled in the fully discharged state.
Galvanostatic cycling was performed at G0.1 A/cm2 at room temperature with
voltage limits of 0.6 and 1.4 V. To obtain the polarization curves, the cell was firstcharged to the desired state of charge and then polarized via linear sweep voltammetry at a rate of 100 mV/s. This method was found to yield polarization curves very
close to point-by-point galvanostatic holds, yet to impose minimal perturbation to
the SOC of the small-electrolyte-volume cell. EIS was performed at SOCs between
10% and 100% at open-circuit potential with a 10 mV perturbation and with frequency ranging from 1 to 300,000 Hz.
Permeability Measurements
The permeability of the oxidized form of 2,6-DBEAQ and potassium ferricyanide
across a Fumasep E-620 (K) membrane was evaluated with a lab-made two-compartment cell. In the first case, the donating side was filled with a solution of 2,6-DBEAQ(0.1 M) in 1.6 M KOH, while the receiving side was filled with 1.6 M KOH. For potassium ferricyanide permeability, the donating side was filled with potassium ferricyanide (0.3 M) in 1 M KOH, while the receiving side was filled with 1.9 M KOH. Bothcompartments had the same volume. The cell was continuously agitated on anutating table. At different time intervals, aliquots were taken from the receivingside, diluted, characterized by UV-visible spectrophotometry, and replaced by fresh
KOH solution. The electrolyte volumes on both sides were checked periodically toensure that there was negligible water flux across the membrane, which might
affect the apparent reactant permeability. The concentration was calculated froma calibration curve and the permeability of 2,6-DBEAQ was calculated based on
Fick’s law using the following equation:
 

2ct
Vo l
Dln 1 
2A
c0
;
P=
Dt
where P is permeability (cm2/s), A is the effective membrane area (cm2), t is elapsed
time (s), ct (mol/L) is the concentration of active species in the receiving side at time t,
Vo is the volume of the solution in either compartment (5 cm3), l is the thickness of themembrane (20 mm), co is the concentration of 2,6-DBEAQ in the donating side attime zero (0.1 mol/L), and D represents a finite difference.
Solubility Tests
Solubility limits of all DBEAQ isomers were measured in their oxidized forms by adding the potassium salt of DBEAQ (prepared by reacting DBEAQ with potassium hydroxide in water) until no further solid could be dissolved. After filtering the mixture
through a polytetrafluoroethylene 0.45 mm syringe filter, a saturated solution of
DBEAQ in KOH was obtained. The saturated solution was then diluted by a known
amount, and the concentration was evaluated by UV-visible spectrometry (Ocean
Optics Flame-S Spectrometer Assembly) at 280 and 364 nm. The concentration
was calculated according to a pre-calibrated absorbance-concentration curve ofknown concentrations of DBEAQ. Solubility of the oxygen-sensitive reduced formof DBEAQ was not measured but was assumed to be higher than the oxidized
form because (1) no precipitation after full electrochemical reduction of DBEAQ
was observed and (2) increasing the number of negative charges from two for
DBEAQ in the oxidized form to four for DBEAQ in the reduced form is expected
to make quinone-quinone interactions even more unfavorable and increase its
solubility.27

1904 Joule 2, 1894–1906, September 19, 2018

SUPPLEMENTAL INFORMATION
Supplemental Information includes Supplemental Experimental Procedures, 18 figures, 2 schemes, and 4 tables and can be found with this article online at https://doi.
org/10.1016/j.joule.2018.07.005.

ACKNOWLEDGMENTS
This research was supported initially by U.S. DOE ARPA-E award DE-AR-0000348 andsubsequently by U.S. DOE award DE-AC05-76RL01830 through PNNL subcontract
304500, by Innovation Fund Denmark via the Grand Solutions project ‘‘ORBATS’’ file
no. 7046-00018B, by the Massachusetts Clean Energy Technology Center, and by
the Harvard School of Engineering and Applied Sciences. D.A.P. acknowledges
funding support from the NSF Graduate Research Fellowship Program. We thankNada Aissami for help with permeability measurements.

AUTHOR CONTRIBUTIONS
K.L. designed the molecule. Y.J. designed the synthetic route and synthesized themolecule; subsequent syntheses were performed by K.L. and E.F.K. K.L., Y.J., and
E.F.K. performed molecular characterization. M.-A.G. and D.G.K. performed symmetric cell tests. D.G.K. performed full cell tests. D.D.P. performed permeability
tests and membrane selection and assisted with cell assembly. E.F.K. performed
chemical analysis of cycled electrolytes. D.A.P. performed chemical analysis and stability studies. D.P.T. performed theoretical calculations under the supervision of
A.A.-G. R.G.G. supervised the molecular synthesis and characterization. M.J.A. supervised the cell assembly, electrolyte composition selection, and performance
evaluation. D.G.K. and M.J.A. drafted the manuscript. All authors edited themanuscript.

DECLARATION OF INTERESTSHarvard University has filed a patent application on the materials described in thispaper.
Received: March 16, 2018
Revised: May 8, 2018
Accepted: July 3, 2018
Published: July 23, 2018

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JOUL, Volume 2

Supplemental Information

Alkaline Quinone Flow Battery
with Long Lifetime at pH 12
David G. Kwabi, Kaixiang Lin, Yunlong Ji, Emily F. Kerr, Marc-Antoni Goulet, Diana De
Porcellinis, Daniel P. Tabor, Daniel A. Pollack, Alán Aspuru-Guzik, Roy G.
Gordon, and Michael J. Aziz

Figure S1. 1H NMR spectrum of 2,6-DBEAQ. 1H NMR (500 MHz, DMSO-d6): δ 11.79 (s, 2H), 8.11 (d, J =8.6 Hz, 2H), 7.54 (d, J = 2.6 Hz, 2H), 7.38 (dd, J = 8.7, 2.6 Hz, 2H), 4.20 (t, J = 6.4 Hz, 4H), 2.43 (t, J =7.3 Hz, 4H), 2.00 (m, 4H). Solvent peaks are those that are not integrated. Final yield: 87%.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

18DBEAQ1-18
STANDARD
PARAMETERS
Name FLUORINE
Shift Range
H's Integral Classs

2 M (m) -0.09 -0.05 ..
-0.14

1

0.70

m

3 I (s)1.89 1.91 ..
1.87

0

0.39

s

4 H (m)1.98 2.03 ..
1.93

2

1.87

m

5 G (t)2.52 2.58 ..
2.48

2

1.90

t

7.27, 7.27

6 F (t)4.15 4.18 ..
4.12

2

1.92

t

6.27, 6.27

7 D (ddt) 7.52 7.55 ..
7.49

1

0.95

ddt

1.29, 8.29, 18.94,
0.00

8 C (m)7.69 7.76 ..
7.62

2

1.84

m

9 B (s)12.17 12.19 ..
12.15

1

1.06

s

14

13

12

11

10

9

8

7

6
f1 (ppm)

5

4

3

2.03

2.76

1.95

3

2.00

J's

-0.08 -0.06 ..
-0.10

0.98
0.97
1.00

1 N (s)submitted 3/16/18, revised 6/18/18

2

1

0

-1

-2

Figure S2. 1H NMR spectrum of 1,8-DBEAQ. 1H NMR (500 MHz, DMSO-d6): δ 12.17 (s, 2H), 7.71 (m,2H), 7.65 (m, 2H), 7.52 (dd, J = 8.3, 1.3 Hz, 2H), 4.15 (t, J = 6.3 Hz, 4H), 2.53 (t, J = 7.3 Hz, 4H), 1.89 (m,4H). Final yield: 64.4%.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

14

13

12

11

10

9

8

7

6
f1 (ppm)

5

4

3

4.08

1.77
2.09

1.96
1.93

1.95
0.96
2.07
0.95

1.90

30April2018old12dbeaqsampleC

2

1

0

-1

-2

Figure S3. 1H NMR spectrum of 1,2-DBEAQ. 1H NMR (500 MHz, DMSO-d6): δ 12.15 (s, 2H), 8.11 (m,2H), 7.97 (dd, J = 8.6, 1.4 Hz, 1H), 7.85 (m, 2H), 7.51 (d, J = 8.7 Hz, 1H), 4.14 (t, J = 6.4 Hz, 2H), 4.00 (t,J = 6.2 Hz, 2H), 2.51 (t, J = 7.4 Hz, 2H), 2.43 (t, J = 7.2 Hz, 2H), 2.01 (m, 4H). Final yield: 72.1%.

CompoundPotassium Ferrocyanide
Potassium Ferricyanide
2,6-DBEAQ

Conditions
KOH, pH 14
KOH, pH 12
KOH, pH 14
KOH, pH 12
KOH, pH 14
KOH, pH 12

Maximum Solubility (M)
0.5
1.2 – 1.3
1.05 –1.13
1.8 – 2.0
1.10
0.60

Table S1. Effect of pH and counter-ion on maximum solubilities of potassium ferrocyanide, potassium
ferricyanide, and 2,6-DBEAQ in KOH. Ranges represent uncertainties in solubility based on different peaks
used in UV-Vis calibration measurements.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

Negolyte
2,6-DHAQ
2,6-DBEAQ
1,2-DBEAQ
1,8-DBEAQ

Solubility (M)at pH 14 (pH 12)
0.6
1.1 (0.60)
0.9
0.75

submitted 3/16/18, revised 6/18/18

Cell Voltage (V)
1.20
1.05
1.04
1.01

Theoretical Energy Density(Wh/L) at pH 14 (pH 12)
11.4
11.5 (17.2)
10.9
10.2

Table S2. Effect of solubility and full cell voltage when paired against a potassium ferrocyanide-based
posolyte on theoretical energy densities of 2,6-DHAQ and different DBEAQ isomers, with the solubility ofpotassium ferrocyanide estimated to be 0.5 M at pH 14 and 1.25 M at pH 12. Solubilities for potassium
ferricyanide and 2,6-DBEAQ at pH 14 and 12 are reported in Table S1. For 2,6-DBEAQ and 2,6-DHAQ,
cell voltages were obtained from the OCV at 50% SOC from full cell tests, whereas differences in CVobtained redox potentials between DBEAQ and potassium ferri/ferrocyanide were used for 1,2- and 1,8DBEAQ.

Figure S4. (a) Rotating Disk Electrode study of the reduction of 5 mM 2,6-DBEAQ in 1M KOH on a glassy
carbon electrode at rotation rates between 400 and 2500 rpm. (b) Levich plot (limiting current vs squareroot of rotation rate in rad/s) of 5 mM 2,6-DBEAQ in 1 M KOH. Data are taken from the current at -0.8 V vsSHE. The slope yields a diffusion coefficient for the oxidized form of 2,6-DBEAQ of 1.58 × 10-6 cm2/s.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

Figure S5. Capacitance-corrected CV of 5 mM 2,6-DBEAQ in 1 M KOH (solid grey line). The dashed redline represents the simulated total current arising from two successive one-electron reductions with
reduction potentials of -0.517 and -0.511 V vs SHE, according to the procedure outlined in a previous
publication. [1] Each simulated reduction has a rate constant ko = 7 × 10-3 cm/s and α = 0.5. The diffusion
coefficient of all redox states of 2,6-DBEAQ (oxidized, semiquinoid, and reduced) was assumed to be 1.58
× 10-6 cm2/s (as obtained from the Levich plot in Figure S4).
Redox chemistry of 2,6-DBEAQ vs 2,6-DHAQ
Following our previous study on 2,6-DHAQ [1] and literature reports [2,3] on quinone electrochemistry in
the absence of proton donors, we simulated the CV of 2,6-DBEAQ as representing two reversible electrontransfer steps at two different potentials E1 and E2. The separation between simulated redox potentials forthe first (E1) and second (E2) electron reductions for 2,6-DBEAQ is considerably narrower than thecorresponding peak separation for 2,6-DHAQ [1] (6 mV vs 60 mV). An important implication of thisdifference is that there is a smaller driving force for the formation of semiquinone radicals between fully
oxidized and fully reduced 2,6-DBEAQ species than there is for 2,6-DHAQ. In order to demonstrate why
this is the case, we first explain the effect of E1-E2 separation on the reduction mechanism and then draw
a connection between the reduction potential difference and the semiquinone concentration.
The pH dependence of the proton-coupled electron transfer (PCET) redox reaction for a generalized
anthraquinone (AQ) is often visualized by the nine-membered square scheme (Figure S6). The horizontal
direction indicates electron transfer (ET), and the vertical direction indicates proton transfer (PT or
protonation). Under strongly alkaline conditions, where the pH is higher than the second pKa (pKa 2) of the
reduced AQ (AQ2-), no protonation of the semiquinone anion will take place, and reduction proceeds via
two ET steps, from AQ to a semiquinone (AQ-●) at E1, and from AQ-● to AQ2- at E2 which is typically < E1.
Depending on the value of E1–E2, the mechanism of AQ reduction in alkaline conditions can either be aconcerted one-step, two-electron process or two single-electron processes. When E2 becomes more

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

negative relative to E1, a stepwise process via a semiquinone radical is followed, as in AQ reduction in
aprotic solvents. [3,4] Under these conditions, a comproportionation reaction, where AQ and AQ 2- combine
to generate AQ-● semiquinone (Equation 3) is thermodynamically favorable. The equilibrium constant forthis comproportionation (Keq_comp) can be obtained by dividing the equilibrium constants of E1 (Equation 4)
by that of E2 (Equation 5), which now becomes Equation 6.[2] The ratio between semiquinone concentrationand the geometric mean of the concentrations of the oxidized and reduced forms of AQ can now be seen
as an increasing function of E1–E2.

𝐴𝑄 + 𝑒 − ⇌ 𝐴𝑄 −· (Eq 1)
𝐴𝑄 −· + 𝑒 − ⇌ 𝐴𝑄 2− (Eq 2)
𝐴𝑄 + 𝐴𝑄 2− ⇌ 2𝐴𝑄 −· (Eq 3)
𝐸1 =
𝐸2 =

𝑅𝑇

𝑙𝑛𝐾𝑒𝑞1 ; 𝑲𝒆𝒒𝟏 =

𝑧𝐹
𝑅𝑇
𝑧𝐹

𝑙𝑛𝐾𝑒𝑞2 ; 𝑲𝒆𝒒𝟐 =𝑲𝒆𝒒_𝒄𝒐𝒎𝒑 =

𝐾𝑒𝑞1
𝐾𝑒𝑞2

=

[𝐴𝑄 −· ][𝐴𝑄][𝐴𝑄2− ]
[𝐴𝑄−· ][𝐴𝑄 −· ]2[𝐴𝑄]∙[𝐴𝑄 2− ]= 𝑒𝑥𝑝 (𝑧𝐹
𝑅𝑇

= 𝑒𝑥𝑝 (

= 𝑒𝑥𝑝 [

𝑧𝐹
𝑅𝑇

𝑧𝐹
𝑅𝑇

𝐸1 ) (Eq 4)
𝐸2 ) (Eq 5)(𝐸1 − 𝐸2 )] (Eq 6)This equation directly relates the concentration of semiquinone species to the E 1-E2 separation andsuggests that Keq_comp for semiquinone formation is 2.83 times higher for 2,6-DHAQ than it is for 2,6-DBEAQ
for the same AQ:AQ2- ratio, based on their respective E1-E2 separations.
The differences in energy between the first and second reduction can also be estimated with electronic
structure calculations. Conformers of the oxidized, one-electron reduced, and two-electron reduced formsof 2,6-DHAQ and 2,6-DBEAQ were generated with the UFF force field using RDKit.[5] The redox active
sites of the one-electron and two-electron reduced forms were paired with one and two potassium cations,
respectively. The deprotonated carboxyl and hydroxyl groups were also paired with potassium cations. Thispairing is done to approximate the ion pairing behavior of the molecules in solution in lieu of an explicit
solvent, as has been done previously.[6] The relative energies of the first and second reduction were
calculated with both semi-empirical methods (PM7 COSMO) and density functional theory (DFT) methods
in a polarizable continuum model (PCM) implicit solvent (B3LYP/6-311+G(d,p)). The DFT single point
energies were evaluated at the B3LYP/6-31G(d) (no PCM) minima. Typically, when calculating reduction
potentials, a calibration scheme is employed, [6] or the relative energetics of all molecules are computed
relative to a fixed molecule where the experimental result is known;[4] the latter strategy is used in this casewith 2-6 DHAQ used as the internal standard. We find that using DFT is necessary to qualitatively predictthese differences (though the shift is overestimated) in the first and second reduction. This suggests thatsuch a behavior can be screened for in future studies, at least for quantitative estimation of semiquinone
presence.
Negolyte

B3LYP/6-311+G(d,p) PCM
Estimated E1 – E2 (V)2,6-DHAQ*
0.060
2,6-DBEAQ-0.133
Table S3. Estimated difference in potential between the first and the second reduction for 2,6-DHAQ and
2,6-DBEAQ. *Note that the energy differences are calibrated so that the differences for 2,6-DHAQ are fixed
to experiment.

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submitted 3/16/18, revised 6/18/18

Figure S6. 9-membered Square Scheme of AQ Reduction. The horizontal direction indicates electrontransfer onto the molecule at a given potential, i.e. E1 and E2, and the vertical direction indicates protonationof the species depending on the pKa of the generated anion and the solution pH. AQ states that are unstable
under alkaline conditions are shown in grey. Blue arrows indicate PCET.
Theoretical calculations of DBEAQ isomer susceptibility to alkyl chain cleavage
The stability of disubstituted DBEAQ isomers against the loss of the alkyl groups was evaluatedcomputationally. The first approach taken was to evaluate the relative reaction energies of the loss of thealkyl chain of both the oxidized and reduced forms of different isomers of DBEAQ. These reactions areshown in Scheme S1.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

Scheme S1. Reaction energies evaluated to determine the relative thermodynamic stability of different
isomers of DBEAQ in different oxidation states.
Optimizations were conducted at the PM7 level of theory with the COSMO solvation model using the
MOPAC package.[7] Up to 20 conformers for each molecule were generated using RDKit.[5] Because themolecule is in alkaline conditions, the energies were evaluated for the deprotonated forms where explicit
potassium cations are placed near the negative charges. The relative energies are shown in Table S4. Two
important observations from these results are: (1) the reaction energies associated with chain loss are more
thermodynamically favorable for the oxidized form in all cases, except in 2,6-DBEAQ, where the reduced
form is slightly more susceptible. Second, the thermodynamic stability of both the oxidized and reduced
forms of 2,6-DBEAQ are among the highest of all the isomers examined.
Thermodynamic Loss of Chain (eV)

Reaction Energy to form
Tetrahedral Intermediate (eV)
From Oxidized Form a,c

DBEAQ Chain Positions

Oxidized Forma,bReduced Forma,b

2,6

0.00

-0.06

0.00

1,2

-0.38

-0.19

-0.21

1,3

-0.27

-0.09

-0.38

1,4

-0.33

0.03

-0.20

1,5

-0.32

-0.13

-0.35

1,6

-0.35

0.00

-0.29

1,7

-0.56

-0.21

-0.21

1,8

-0.34

0.02

-0.27

2,3

-0.30

0.71

-0.21

2,7

0.00
0.86
-0.04
Table S4. Relative Energies of DBEAQ. a Energies calculated at the PM7 level of theory with the COSMO
solvation model b Reaction energies are tabulated relative to the energy of 2,6-DBEAQ losing an alkyl
chain in its oxidized form c Reaction energies to form tetrahedral intermediate are relative to the reactionenergy for the oxidized form of 2,6-DBEAQ to form a tetrahedral intermediate.

Because the relative stability of the reduced form over the oxidized form cannot be completely explainedby looking at only the thermodynamics of the final products, we investigated the thermodynamic landscapeof possible intermediates. As the oxidized form is more electron deficient, we estimate the susceptibility tochain loss by evaluating the energy of a hydroxide-substituted intermediate of the DBEAQ molecule(Scheme S2). This intermediate should more easily form for the oxidized form. As done previously withbromination of anthraquinones,[8] the energy is estimated by evaluating the relative energy of theSI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

tetrahedral intermediate to that of the oxidized molecule (right-most column of Table S4). Here, we see that
2,6-DBEAQ is the most stable isomer, whereas 1,2- and 1,8-DBEAQ are less stable than 2,6-DBEAQ,
which is consistent with experimental observation. Based on these findings, it appears that DBEAQ stability
against alkyl chain loss is highly sensitive to the exact positioning of the two alkyl chains, particularly withrespect to the ketone groups on the anthraquinone core.

Scheme S2. Reaction of OH- with oxidized form to form tetrahedral intermediate. Used as a means ofestimating the vulnerability of the oxidized form to decomposition.

96
0.10 M 2,6-DBEAQ, pH 14< 0.01%/day

charge (C)

92

88

84

reductionoxidationoxidation
reduction

0.10 M 2,6-DHAQ, pH 14~5%/day80
0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0
time (days)

Figure S7. Unbalanced compositionally-symmetric cell cycling of 0.10 M 2,6-DBEAQ and 2,6-DHAQ,
showing capacity vs time, both in 1 M KOH (pH 14). The capacity-limiting side was 5 mL 0.1 M 2,6-DBEAQ
or DHAQ, while the non-capacity-limiting side was 10 mL of the same. Capacities were obtained by fullpotentiostatic reduction and oxidation of 5 mL of capacity-limiting side; potential was switched between ±0.2
V when magnitude of current dropped to 2 mA/cm2. Cell cycle period is ~ 20 min.

Identification of a chemical decomposition pathway of 2,6-DBEAQ
Chemical stability studies performed at 75C, 0.1 M concentration, and pH 14 indicate that the oxidized
form of 2,6-DBEAQ is susceptible to decomposition (Figure S8). Samples of 2,6-DBEAQ were stored in theoxidized form at 0.1 M concentration in pH 14 aqueous solution (1 M KOH) in FEP vials in an oven at 75C

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submitted 3/16/18, revised 6/18/18

for 8 days. Decomposition peaks were observed in the 1H NMR spectra (Figure S8c). Samples of 2,6DBEAQ that were electrochemically reduced and treated under identical conditions exhibit no evidence of
decomposition (Figure S8b).
We have identified that -hydroxybutyrate cleavage is involved in the decomposition of the oxidized form of
2,6-DBEAQ (Figure S9). The aliphatic region of the 1H NMR spectrum of 2,6-DBEAQ treated at 75C for 8
days at 0.1 M concentration in pH 14 aqueous solution (1 M KOH) in the oxidized form (Figure S9c) was
compared to that of -butyrolactone (Sigma Aldrich) prepared in a pH 14 aqueous solution (1 M KOD in
D2O with 10 mM NaCH3SO3 internal standard) (Figure S9b), as -butyrolactone hydrolyzes to its acyclic
form, -hydroxybutyrate, under basic conditions. The spectrum of 2,6-DBEAQ treated at 75C (Figure S9c)
includes protons D, E, and F, corresponding to the 2,6-DBEAQ aliphatic protons (Figure S9a); protons D´´,E´´, and F´´, corresponding to those of -hydroxybutyrate (Figure S9b); and protons D´, E´, and F´,corresponding to the aliphatic protons of the monoalkylated decomposition product (4-((6-oxido-9,10-dioxo9,10-dihydroanthracen-2-yl)oxy)butanoate) after the cleavage of -hydroxybutyrate. The spectrum of 2,6DBEAQ treated at 75C with 10 mM (0.5 µL) -hydroxybutyrate added to the NMR sample (Figure S9d)
exhibits an increase in the size of the peaks corresponding to protons D´´, E´´, and F´´, verifying the
involvement of -hydroxybutyrate cleavage in the decomposition of the oxidized form of 2,6-DBEAQ. The
relative peak integrals of the spectra of 2,6-DBEAQ treated at 75C are also consistent with the productsof -hydroxybutyrate cleavage. This decomposition pathway for the oxidized form of 2,6-DBEAQ, along withthe stability of the relatively more electron-rich reduced form of 2,6-DBEAQ, is consistent with a nucleophilichydroxide-mediated decomposition mechanism. As we observe that the monoalkylated decomposition
product further decomposes into 2,6-DHAQ via cleavage of a second -hydroxybutyrate moiety, the byproducts of -hydroxybutyrate cleavage may be susceptible to further decomposition leading tocorresponding capacity fade, as previously observed in 2,6-DHAQ (Figure S7).

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submitted 3/16/18, revised 6/18/18

Figure S8. NMR evidence of decomposition of 2,6-DBEAQ in the oxidized form. 1H NMR spectra (500 MHz,
1 M KOD in D2O with 10 mM NaCH3SO3 internal standard) of (a) the oxidized form of 2,6-DBEAQ in pH 14
aqueous solution (1 M KOH); (b) 2,6-DBEAQ treated at 75C for 8 days at 0.1 M concentration in pH 14
aqueous solution (1 M KOH) in the reduced form and then re-oxidized in order to compare to samples
tested in the oxidized form; (c) 2,6-DBEAQ treated at 75C for 8 days at 0.1 M concentration in pH 14
aqueous solution (1 M KOH) in the oxidized form.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

Figure S9. Identification of a decomposition product of 2,6-DBEAQ. 1H NMR spectra (500 MHz, 1 M KOD
in D2O with 10 mM NaCH3SO3 internal standard) of (a) aliphatic region of the oxidized form of 2,6-DBEAQ
in pH 14 aqueous solution (1 M KOH); (b) -hydroxybutyrate; (c) aliphatic region of 2,6-DBEAQ treated at
75C for 8 days at 0.1 M concentration in pH 14 aqueous solution (1 M KOH) in the oxidized form; (d)aliphatic region of 2,6-DBEAQ treated at 75C for 8 days at 0.1 M concentration in pH 14 aqueous solution(1 M KOH) in the oxidized form with 10 mM -hydroxybutyrate added to the NMR sample.

Characterization of the time course of -hydroxybutyrate cleavage
Further chemical stability studies were performed to characterize the time course of -hydroxybutyratecleavage of 2,6-DBEAQ. Samples of 2,6-DBEAQ in the oxidized form at 0.1 M and 0.5 M concentrations indeuterated aqueous solutions at pH 12 and pH 14 were stored in sealed 5 mm Precision Low
Pressure/Vacuum Valve (LPV) NMR sample tubes (Wilmad-LabGlass Product No. 535-LPV-7), lined with
5 mm FEP tube liners (Wilmad-LabGlass Product No. 6005-8) to prevent etching of the NMR tubes underalkaline conditions. The samples were heated in ovens at 65C and 95C for a period of several weeks,during which reaction progress was monitored by NMR. Two equivalent samples were prepared and tested
for each condition. Sealed LPV NMR tubes were used to prevent the escape of water vapor from the sample
tubes during heating. pH 14 solutions were prepared by dissolving 2,6-DBEAQ in D2O with theconcentration of KOD required to fully deprotonate the molecule plus an additional 1 M KOD (2 M KOD for
0.5 M 2,6-DBEAQ and 1.2 M KOD for 0.1 M 2,6-DBEAQ). pH 12 solutions were prepared by dissolving 2,6DBEAQ in D2O with the concentration of KOD required to fully deprotonate the molecule plus an additional

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

10 mM KOD (1.01 M KOD for 0.5 M 2,6-DBEAQ and 0.21 M KOD for 0.1 M 2,6-DBEAQ), with additional
KOD added as necessary until pH 12 was reached, as measured by a pH meter (Oakton WD-35419-12 pH
700 Benchtop Meter, Oakton WD-35805-06 pH electrode). The fraction of 2,6-DBEAQ remaining in eachsample was measured at intervals as the ratio of the total integral of the aromatic 1H peaks of 2,6-DBEAQ
to the total integral of all aromatic 1H peaks in the sample. To perform this analysis on samples at 0.5 M
concentration, in which 1H peaks overlap due to peak broadening, deconvolution was performed using theline fitting feature in MestReNova to obtain 1H peak integrals. To ensure that the ratio of aromatic 1H peaksof 2,6-DBEAQ to the total integral of all aromatic 1H peaks corresponds to the concentration of 2,6-DBEAQ
remaining in a sample relative to the initial 2,6-DBEAQ concentration, the integral of one 1H peak of 2,6DBEAQ was measured relative to that of an internal standard (0.1 M NaCH 3SO3 in D2O). The internal
standard solution was prepared in a sealed capillary tube that was inserted in each NMR tube before ameasurement was taken and removed from the sample during heating to prevent decomposition of andreactions with the internal standard. NMR tubes were cooled in a water bath before the valve was removed
to insert the capillary tube to prevent the escape of water vapor from the sample tubes. Good agreement
was found between both methods of tracking the time course of 2,6-DBEAQ decomposition, confirming thatthe ratio of aromatic 1H peaks of 2,6-DBEAQ to the total integral of all aromatic 1H peaks could be used asa measure of reaction progress.
Figure S10 illustrates the time course of 2,6-DBEAQ decomposition for samples treated at 65C and 95C,
at pH 12 and pH 14, and at 0.1 M and 0.5 M concentration in the oxidized form. Error bars correspond to
the standard deviation of measurements obtained from equivalent samples. Samples of 2,6-DBEAQ thatwere electrochemically reduced and stored for 13 days in sealed FEP-lined LPV NMR tubes at the same
temperature, pH, and concentration conditions as those reported here for the oxidized form exhibit lessthan 10% decomposition in all conditions and less than 5% decomposition in all except the 95C, pH 12,
0.5 M condition.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18[DBEAQ] vs. Time
1.0
0.9

[DBEAQ] / [DBEAQ]0

0.8
0.7
0.6
0.5
65C, pH 12, 0.5 M
65C, pH 12, 0.1 M
65C, pH 14, 0.5 M
65C, pH 14, 0.1 M
95C, pH 12, 0.5 M
95C, pH 12, 0.1 M
95C, pH 14, 0.5 M
95C, pH 14, 0.1 M

0.4
0.3
0.2
0.1
0.0
0

2

4

6

8

10

12

14

16

Time (days)
Figure S10. Time course of 2,6-DBEAQ decomposition at various temperature, pH, and concentration
conditions. The remaining concentration of 2,6-DBEAQ relative to the initial concentration vs time (indays) for samples treated at 65C and 95C, at pH 12 and pH 14, and at 0.1 M and 0.5 M concentration in
the oxidized form. Error bars correspond to the standard deviation of measurements obtained fromequivalent samples.
For samples at pH 14 and at 0.1 M 2,6-DBEAQ concentration (in which the concentration of hydroxide ionsexceeds that of 2,6-DBEAQ by tenfold so a pseudo reaction order with respect to 2,6-DBEAQ may beinferred), the linear fit of the semi-logarithmic plot of 2,6-DBEAQ concentration vs time, shown in Figure
S11a for the 95C condition, is consistent with pseudo-first order reaction kinetics with respect to 2,6DBEAQ (R2 = 0.994 for the 95C condition, R2 = 0.987 for the 65C condition). Comparison of the semilogarithmic plot (Figure S11a) to the same data plotted on a linear scale (Figure S11b, R2 = 0.919) and as
the reciprocal of 2,6-DBEAQ concentration vs time (Figure S11c, R2 = 0.968) confirms the inference ofpseudo-first order kinetics rather than pseudo-zero order or pseudo-second order kinetics with respect to
2,6-DBEAQ. The corresponding effective (pseudo-first order) rate constants are 0.10 days -1 at 95C and0.012 days-1 at 65C. The corresponding activation energy estimated from the Arrhenius equation is 70
kJ/mol (20 kcal/mol). The effective rate constant extrapolated to 25C is 0.0004 days-1, corresponding to aroom temperature half-life on the order of 5 years for decomposition of the oxidized form of 2,6-DBEAQ at
pH 14 and at 0.1 M concentration. In the parameter space studied (Figure S10), we report slower
decomposition at pH 12 relative to pH 14 and slower decomposition at 0.5 M relative to 0.1 M concentration.
We also report substantially slower decomposition of the reduced form relative to the oxidized form of 2,6DBEAQ.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

Asubmitted 3/16/18, revised 6/18/18

ln([DBEAQ]) vs. Time0.0

R2 = 0.994

ln([DBEAQ] / [DBEAQ]0)

-0.2
-0.4
-0.6
-0.8
-1.0
-1.2
-1.4
-1.6
-1.8
-2.0
0

2

4

6

8

10

12

14

16

18

20

Time (days)

C

B
R2 = 0.919

R2 = 0.968

Figure S11. Evidence of pseudo-first order decomposition kinetics with respect to 2,6-DBEAQ. The
remaining concentration of 2,6-DBEAQ relative to the initial concentration vs time (in days) for samples
treated at 95C, at pH 14, and at 0.1 M concentration in the oxidized form, plotted (a) on a semilogarithmic scale; (b) on a linear scale; (c) as the reciprocal of the 2,6-DBEAQ concentration relative tothe initial concentration vs time. Error bars correspond to the standard deviation of measurements
obtained from equivalent samples.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

A note on DBEAQ/DHAQ permeability
The permeability of the oxidized form of 2,6-DBEAQ (5.26 × 10-13 cm2/s) in Fumasep E-600 series
membrane (with ion exchange capacity of 1.70 meq/g) is about an order of magnitude lower than is typicalfor permeabilities of quinone-based molecules under alkaline conditions across more widely usedcommercial membranes (e.g. Nafion). Permeability of the oxidized form of 2,6-DBEAQ (with two negative
charges) was measured because we expect it to have a higher permeability than the reduced form, whichshould cross over more slowly through the cation-exchange membrane owing to it having four negative
charges instead of two.

2.0x10-6

1.5x101.0x10-6

2,6-DHAQ in Nafion 212
2,6-DHAQ in Fumasep E-610 (K)6.0x10-6

Ferricyanide= 4.4 E-12 cm2/s
2,6-DHAQ= 7.2 E- 13 cm2/s
2,6 DBEAQ= 5.26 E-13 cm2/s

5.0x10-7

Ln[1-2Cr/Cd] [-Vd/2A]-6

Ln[1-2Cr/Cd] [-Vd/2A]

7.0x10-6

Ferricyanide in Fumasep E-610 (K)2,6-DHAQ in Fumasep E-610 (K)2,6-DBEAQ in Fumasep E-610 (K)5.0x10-6
4.0x10-6
3.0x10-6
2.0x10-6
1.0x10-6
0.0

0.0
0

2

4

6

time (days)

8

10

12

14

0

1

2

3

4

5

6

7

8

9

10

11

12

time (days)

Figure S12. (a) The permeability of the oxidized form of 2,6-DBEAQ, 2,6-DHAQ, and potassium ferricyanide
across a Fumasep E-610 (K) membrane. (b) 2,6-DHAQ in Nafion 212 and Fumasep E-610 (K). The slope
represents the permeability.
We found that 2,6-DHAQ had a similarly low permeability across Fumasep E-610 (K) (10 μm, 7.24 × 10-13
cm2/s), in contrast to Nafion 212 (50 μm, 1.39 × 10-11 cm2/s).
These permeabilities imply that, assuming the membrane remains chemically stable, it would take at least
3,000 years for 2,6-DBEAQ/DHAQ permeation through Fumasep E-610 (K) to lead to a 50% loss in cell
capacity; it would take 200 years for potassium ferricyanide permeation to lead to a similar loss, as it is thefaster-crossing species. Fluoropolymer membranes for vanadium redox batteries comprise over 40% of thecost of the power stack and a significant fraction of the cost of the entire system.[9] The anticipated cost for
Fumasep E-620 (K) (20 μm) at large production scale is < $25/m 2, which is more than an order of magnitudeless expensive than fluoropolymers (Bernd Bauer, private communication, March 17, 2018).

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

Figure S13. Capacity and current efficiency (capacity on discharge divided by the capacity of the preceding
charge step) as functions of time during galvanostatic cycling of a cell with a negolyte of 5.5 mL 0.1 M 2,6DBEAQ and posolyte of 15 mL 0.20 M potassium ferrocyanide and 0.08 M potassium ferricyanide, both at
pH 14 (880 cycles with each cycle 6.5 minutes long). Galvanostatic cycling was performed at 100 mA/cm 2with voltage cutoffs of 1.4 and 0.6 V on charge and discharge, respectively.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

Figure S14. Unbalanced compositionally-symmetric cell cycling of 0.10 M (a) 1,2- and (b) 1,8-DBEAQ,
showing capacity vs time, both in 0.01 M KOH (pH 12). The capacity-limiting side was 5 mL 0.1 M DBEAQ,
while the non-capacity-limiting side was 10 mL of the same. Capacities were obtained by full potentiostatic
reduction and oxidation of 5 mL of capacity-limiting side; potential was switched between ±0.2 V whenmagnitude of current dropped to 2 mA/cm2. Linear fits between 0.01 and 10 %/day are drawn to compare
against experimental data.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

Figure S15. Unbalanced compositionally-symmetric cell cycling of 0.10 M 2,6-DBEAQ, showing capacity
vs time in 0.01 M KOH (pH 12). The capacity-limiting side was 5 mL 0.1 M 2,6-DBEAQ, while the noncapacity-limiting side was 10 mL of the same. Capacities were obtained by full potentiostatic reduction andoxidation of 5 mL of capacity-limiting side; potential was switched between ±0.2 V when magnitude ofcurrent dropped to 2 mA/cm2. Dashed lines with slopes of 0.01 and 0.1 %/day are drawn to compare against
experimental data, which showed a temporal fade rate of approximately 0.007 %/day.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

Figure S16. (a) Pourbaix diagram of 2,6-DBEAQ with a slope fit to the data of -35 mV/pH below pH 12.
Above pH ~ 11.5, the potential is pH-independent, indicating that both oxidized and reduced forms of 2,6DBEAQ are deprotonated. The dashed line has zero slope. (b) Capacitance-corrected CVs of 5 mM 2,6DBEAQ in 1 M KOH (black curve) and pH 12 buffer (red curve) solution. We hypothesize the wider peakseparation in the buffered pH 12 solution to be a result of slower kinetics for the first and second electron
reduction steps[3] rather than a larger difference between their redox potentials.

Figure S17. Power density vs current density for 0.5 M 2,6-DBEAQ at 20C at 50% and ~100% SOC.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

Figure S18. (a) NMR analysis of electrolytes from a cell compositionally similar to that in Figure 4a andcycled for 11 days. The origin of the peak shifts in NMR is unknown but could stem from differences slightdifferences in pH. (b) CV measurement of cycled posolyte showing no 2,6-DBEAQ redox peaks within thepotential window (vs Ag/AgCl).

Estimated cost of electrolytes
The U.S. Department of Energy has a cost target of $150/kWh for the entire storage system. The current
price for vanadium pentoxide (98% pure) is $15.15/lb. (www.vanadiumprice.com accessed 6/15/2018),
which equates to $160/kWh for both electrolytes of a 1.45-volt all-vanadium flow battery. To this, one
should add the cost of purification, but we do not do so here.
The mass-production cost of 2,6-DBEAQ, a new composition of matter, is not known at present, but costprojections for similar anthraquinones DHAQ (Fig. S1 of ref. [10]) and AQDS (ref. [11]) were recently
reported to be between $0.92/kg and $3.92/kg when manufactured in the quantities necessary for GWh
storage. (We note that the world’s largest battery, the Rongke Power vanadium redox flow battery underconstruction near Dalian, China, is to have a capacity of 3.2 GWh.) The similarity of 2,6-DBEAQ to, and its
synthesis from, 2,6-DHAQ make it plausible that its cost will not be much greater than that of DHAQ. Before
we can perform a similar estimate of the cost of 2,6-DBEAQ vs scale, an industrially-compatible synthesis
method must be developed with attributes such as high yield, and then the process analyzed for cost. Until
then, we use $0.92/kg - $3.92/kg as an estimate of the cost range for anthraquinones at large scale,
resulting in between $7/kWh and $29/kWh for the 2-electron-accepting DBEAQ in the negolyte.

SI: Alkaline Quinone Flow Battery with Long Lifetime at pH 12

submitted 3/16/18, revised 6/18/18

A recent paper by Yuan et al. [12] quotes the cost of sodium ferrocyanide at $1.07/kg. Although we usedpotassium ferrocyanide in our system, we do not expect this to cost much more than twice the price of thesodium form based on similar comparisons from the same source (from Yuan et al., NaOH costs $0.4/kg,
whereas KOH costs $0.84/kg). Hence, we estimate a cost of roughly $2.15/kg for potassium ferrocyanide,amounting to $32/kWh for the posolyte in our 1.05 V battery. Therefore, at large scale, we expect the
combined cost of our electrolytes to be between $39/kWh and $61/kWh. To the extent that we are able tomix Na+ with K+ in the electrolyte, as discussed above to raise solubility, the cost of chemicals will be further
decreased.

Supplemental References
1K. Lin, Q. Chen, M.R. Gerhardt, L. Tong, S.B. Kim, L. Eisenach, A.W. Valle, D. Hardee, R.G. Gordon,

M.J. Aziz, and M.P. Marshak, "Alkaline Quinone Flow Battery", Science 349, 1529 (2015).
2C. Costentin, "Electrochemical Approach to the Mechanistic Study of Proton-Coupled Electron Transfer",

Chem. Rev. 108, 2145 (2008).
3M. Quan, D. Sanchez, M.F. Wasylkiw, and D.K. Smith, "Voltammetry of Quinones in Unbuffered

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