RE S EAR CH | R E P O R T S

32. K. Niefind, J. Müller, B. Riebel, W. Hummel, D. Schomburg,
J. Mol. Biol. 327, 317–328 (2003).
ACKN OWLED GMEN TS

The research leading to these results received funding from theEuropean Union’s Seventh Framework Programme FP7/2007-2013
under grant agreement 266025 (BIONEXGEN). The project leading
to this application has received funding from the European
Research Council under the European Union’s Horizon 2020
research and innovation program (grant agreement no. 638271,
BioSusAmin). T.K. and N.S.S. received funding from UK
Biotechnology and Biological Sciences Research Council (BBSRC;
BB/K0017802/1). N.J.T. is grateful to the Royal Society for a
Wolfson Research Merit Award. Some aspects of the results
reported here are the subject of a provisional patent application.
Author contributions: F.G.M. and N.J.T. conceived the project andBATTERIES

Alkaline quinone flow battery
Kaixiang Lin,1 Qing Chen,2 Michael R. Gerhardt,2 Liuchuan Tong,1 Sang Bok Kim,1Louise Eisenach,3 Alvaro W. Valle,3 David Hardee,1 Roy G. Gordon,1,2*
Michael J. Aziz,2* Michael P. Marshak1,2*
Storage of photovoltaic and wind electricity in batteries could solve the mismatch problem
between the intermittent supply of these renewable resources and variable demand. Flowbatteries permit more economical long-duration discharge than solid-electrode batteries byusing liquid electrolytes stored outside of the battery. We report an alkaline flow battery basedon redox-active organic molecules that are composed entirely of Earth-abundant elements
and are nontoxic, nonflammable, and safe for use in residential and commercial environments.
The battery operates efficiently with high power density near room temperature. These results
demonstrate the stability and performance of redox-active organic molecules in alkaline
flow batteries, potentially enabling cost-effective stationary storage of renewable energy.

wrote the manuscript; F.G.M. planned the experiments, expressedand purified the AmDHs, performed the biocatalytic reactions, andanalyzed the data; T.K. performed the gene cloning of all AmDHs
and purified the ADHs; N.S.S. and M.B. provided intellectual andtechnical support; and M.B. and BASF provided the ADHs. We
thank R. Heath for a preliminary kinetic assay of the Ph-AmDH.
SUPPLEMENTARY MATERIALS

www.sciencemag.org/content/349/6255/1525/suppl/DC1
Materials and Methods
Figs. S1 to S12
Tables S1 to S20
References (33–36)
30 June 2015; accepted 14 August 201510.1126/science.aac9283

energy components of the system by storing theelectro-active species outside the battery container
itself (3–5). In a flow battery, the power is generated in a device resembling a fuel cell, whichcontains electrodes separated by an ion-permeable
membrane. Liquid solutions of redox-active species are pumped into the cell, where they can be
charged and discharged, before being returned tostorage in an external storage tank. Scaling theamount of energy to be stored thus involves simply making larger tanks (Fig. 1A). Existing flow
batteries are based on metal ions in acidic solution, but challenges with corrosivity, hydrogen evolution, kinetics, materials cost and abundance, andefficiency thus far have prevented large-scale commercialization. The use of anthraquinones in anacidic aqueous flow battery can dramatically1T

he cost of photovoltaic (PV) and wind electricity has dropped so much that one of thelargest barriers to getting most of our electricity from these renewable sources is theirintermittency (1–3). Batteries provide a means

to store electrical energy; however, traditional, enclosed batteries maintain discharge at peak power
for far too short a duration to adequately regulate
wind or solar power output (1, 2). In contrast, flow
batteries can independently scale the power andFig. 1. Cyclic voltammetry of electrolyte and cell schematic. (A) Schematicof cell in charge mode. Cartoon on top of the cell represents sources ofelectrical energy from wind and solar. Curved arrows indicate direction ofelectron flow, and white arrows indicate electrolyte solution flow. Blue arrow indicates migration of cations across the membrane. Essential components of electrochemical cells are labeled with color-coded lines and

SCIENCE sciencemag.org

Department of Chemistry and Chemical Biology, Harvard
University, 12 Oxford Street, Cambridge, MA 02138, USA.
Harvard John A. Paulson School of Engineering and Applied
Sciences, 29 Oxford Street, Cambridge, MA 02138, USA.
3
Harvard College, Cambridge, MA 02138, USA.
2

*Corresponding author. E-mail: aziz@seas.harvard.edu (M.J.A.);
gordon@chemistry.harvard.edu (R.G.G.); michael.marshak@colorado.
edu (M.P.M.)

text. The molecular structures of oxidized and reduced species are shownon corresponding reservoirs. (B) Cyclic voltammogram of 2 mM 2,6-DHAQ(dark cyan curve) and ferrocyanide (gold curve) scanned at 100 mV/s on
glassy carbon electrode; arrows indicate scan direction. Dotted line represents CV of 1 M KOH background scanned at 100 mV/s on graphite foil
electrode.
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25. N. H. Schlieben et al., J. Mol. Biol. 349, 801–813(2005).
26. K. Tauber et al., Chemistry 19, 4030–4035 (2013).
27. See supplementary materials on Science Online.
28. A. C. Price, Y.-M. Zhang, C. O. Rock, S. W. White, Structure 12,417–428 (2004).
29. The Ph-AmDH tolerates different buffers well (i.e., ammonium
chloride, formate, borate, citrate, acetate, oxalate). However, thehomotetrameric LBv-ADHs contains two Mg2+ centers; therefore,
strong chelators such as oxalate, citrate, and acetate must be
avoided as buffer salts. Removal of these Mg2+ ions inactivates the
enzyme (32).
30. B. R. Bommarius, M. Schürmann, A. S. Bommarius, Chem.
Commun. 50, 14953–14955 (2014).
31. R. Cannio, M. Rossi, S. Bartolucci, Eur. J. Biochem. 222,
345–352 (1994).

R ES E A RC H | R E PO R TS

reduce battery costs (6, 7); however, the use ofbromine in the other half of the system precludes
deployment in residential communities owing to
toxicity concerns.
We demonstrate that quinone-based flow batteries can be adapted to alkaline solutions, wherehydroxylated anthraquinones are highly soluble
and bromine can be replaced with the nontoxic
ferricyanide ion (8, 9)—a food additive (10). Functionalization of 9,10-anthraquinone (AQ) withelectron-donating groups such as OH has beenshown to lower the reduction potential and expand the battery voltage (6). In alkaline solution,these OH groups are deprotonated to provide solubility and greater electron donation capability,
which results in an increase in the open-circuit
voltage (OCV) of 47% over the previously reported
system. Because functionalization away from the
ketone group provides molecules with the highest
solubility (11, 12), we initially targeted commerciallyavailable 2,6-dihydroxyanthraquinone (2,6-DHAQ),

which we find exhibits a room-temperature solubility of >0.6 M in 1 M KOH. This system canachieve power densities of >0.45 W cm−2 at room
temperature and 0.7 W cm−2 at 45°C.
The use of alkaline electrolyte exploits pH as aparameter to shift the thermodynamic potentials
of proton-dependent reactions to more negativevalues. In acid solutions, AQ undergoes a twoelectron two-proton reduction at a single potential, which shifts to more negative values as the
pH increases (6). When the pH exceeds 12, thereduction potential of 2,6-DHAQ becomes pHindependent because the reduced species is generated in its fully deprotonated form (fig. S1). In
contrast with the pH-dependent electrochemical
behavior of quinones (negative terminal), theferro/ferricyanide redox couple (positive terminal)
has a pH-independent redox potential. This contrasting pH dependence can be exploited throughthe development of low–reduction potential quinones at high pH. The cyclic voltammograms(CVs) of 2,6-DHAQ and ferro/ferricyanide predictan equilibrium cell potential of 1.2 V upon combination of these two half-reactions (Fig. 1B). A
quantitative analysis of the CV of 2,6-DHAQ at
pH 14 (fig. S2) revealed redox behavior consistent
with two one-electron reductions at potentials
separated by only 0.06 V, with a rapid kinetic rate
similar to that of quinones in acid (6). This behavior raises interesting questions about the relationship between quinone redox and hydrogenbonding (13).
Cell testing was performed at 20°C with solutions of 0.5 M 2,6-DHAQ dipotassium salt and
0.4 M K4Fe(CN)6, both dissolved in 1 M KOH.
These solutions were pumped through a flow cell
constructed from graphite flow plates and carbon
paper electrodes, which were separated by a Nafion
membrane. A charging current of 0.1 A cm−2 wasapplied to charge the cell, and polarization curves
were measured at 10, 50, and 100% states ofquinone charge (SOC). The OCV is 1.2 V at 50%

Fig. 2. Cell performance. (A) Cell open-circuit voltage versus state of charge. All potentials were taken when cell voltage stabilized to within ±1 mV. One
hundred percent SOC was reached by potentiostatic holding at 1.5 V until the current decreased to below 20 mA/cm2. (B and C) Cell voltage and power density
versus current density at 20° and 45°C, respectively, at 10, 50, and ~100% SOC. Electrolyte composition: At 20°C, 0.5 M 2,6-DHAQ and 0.4 M ferrocyanide
were used in negative electrolyte and positive electrolyte, respectively. At 45°C, both concentrations were doubled. In both cases, potassium hydroxide content
started at 1 M for both sides in the fully discharged state.

Fig. 3. Cell cycling performance. (A) Representative voltage versus time curves during 100 charge-discharge cycles at 0.1 A/cm2, recorded between the10th and 19th cycles. (B) Capacity retention, current efficiency, and energy efficiency values of 100 cycles. Normalized capacity is evaluated based on thecapacity of the first charge and discharge cycle.

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SOC; its dependence on SOC is shown in Fig. 2A.
The polarization curves (Fig. 2B) show no sign ofredox kinetic limitations and exhibit a peak galvanic power density exceeding 0.4 W cm−2.
The cell was cycled at a constant current density of ±0.1 A cm−2 for 100 cycles (Fig. 3A). The
current efficiency exceeded 99%, with a stableround-trip energy efficiency of 84%. A 0.1% loss
in capacity per cycle was observed during cycling,which appears to be a continuous loss of electrolyte over the 100 cycles. Three possible loss
mechanisms were explored: chemical decomposition, electrolyte crossover through the membrane, and leakage from the pumping system.
Chemical and electrochemical stability studies
showed that the negative electrolyte is stable.
Ten millimolar 2,6-DHAQ was heated in 5 M KOH
solution at 100°C for 30 days and was characterized by proton nuclear magnetic resonance(NMR). Cycled negative electrolyte was also collected and characterized by the same method;both studies showed no degradation product at
the sensitivity level of 1% (fig. S3). Membranecrossover contamination has been a common challenge in acid-based redox flow batteries, wheremost electro-active molecules are either neutralor positive and tend to migrate through protonconductive membranes (5). In this alkaline system,

however, all the electro-active molecules remain
negatively charged in all charge states, leading toa dramatic decrease in the degree of crossoverduring cell cycling. Cyclic voltammetry of the ferro/
ferricyanide electrolyte collected at the end of
cycling showed no evidence of the presence of
2,6-DHAQ. This observation places an upper
limit on crossover of 0.8% of the DHAQ, implying a crossover current density of <2.5 mA cm−2
(fig. S4). Finally, hydraulic leakage was investigated because an apparent but unquantifiable
small decrease in fluid levels was observed in thereservoirs. After cell cycling, the cell was washedwith deionized water until no coloration of eluent
could be observed. The cell was then disassembled;coloration was found on the gaskets, indicating
the likely site of electrolyte leakage (fig. S5). This
source of capacity loss—equivalent to roughlyeight drops in our system—is expected to become
negligible as system size is scaled up.
By increasing the temperature to 45°C, thepeak galvanic power density increases from 0.45
to ~0.7 W cm−2 (Fig. 2C), as the cell area-specificresistance (ASR) decreases from about 0.878 to
0.560 ohm cm2, estimated from the linear partsof the polarization curves in Fig. 2. Most of thisASR decrease comes from a change in the highfrequency ASR (rhf) measured by electrochemical

impedance spectroscopy (fig. S6). In both cases,
the rhf contributes more than 70% of the ASR
and is indeed the limiting factor to the cell current and power outputs. The rhf is dominated by
the resistance of the membrane, which is an order of magnitude higher than the resistance ofthe same membrane in a pH 0 acid solution (14).
The sluggish kinetics of the hydrogen evolution
reaction in alkaline solution on carbon electrodes
results in a larger practical stability window inbase rather than in acid (fig. S7). Consequently,
quinones with substantially more negative reduction potentials are feasible as negative electrolytematerials. Preliminary investigations into the synthesis of different hydroxy-substituted anthraquinones suggest that further increases in cellpotential are possible. Self-condensation reactionsof substituted benzene yield 2,3,6,7-tetrahydroxyAQ (THAQ) (15) and 1,5-dimethyl-2,6-DHAQ (15DMAQ) (figs. S8 and S9). The CVs of these species
in 1 M KOH suggest cell potentials versus ferri/ferrocyanide approaching 1.35 V (Fig. 4, A and B),
which exceeds that of many aqueous rechargeable
batteries (Fig. 4C).
The cyanide ions in both ferro- and ferricyanide are bound too tightly to be released under most conditions. Consequently, it is nontoxic
in both oxidized and reduced forms and is evenFig. 4. Molecular structure and cyclic voltammetry of 2,6-DHAQ derivatives. (A and B) Molecular structures and CV of 2,3,6,7-THAQ (pink) and1,5-DMAQ (olive), respectively, plotted along with ferrocyanide (orange curve)scanned at 100 mV/s on glassy carbon electrode. Both 2,6-DHAQ derivatives/
ferrocyanide couples showed higher equilibrium potential than 2,6-DHAQ/
ferrocyanide. (C) Selected aqueous secondary batteries showing voltage andflow status. Literature data from (3) and (16).

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R ES E A RC H | R E PO R TS

permitted for use as a food additive (10). The use offerrocyanide offers notable advantages over bromine because it is nonvolatile and noncorrosive,
allowing simpler and less expensive materials of
construction. In addition, these tri- and tetraanionic organometallic molecules exhibit low crossover rates through cation-exchange membranes.
The results reported herein highlight the abilityof hydroxy-substituted anthraquinone and ferrocyanide to function as stable flow battery electrolytes in alkaline solution. The use of organic andorganometallic coordination complexes in base,rather than aqueous metal ions in acid, resolves
serious cost, corrosion, and safety concerns ofprevious flow battery chemistries. Alkaline flow
batteries can compensate for higher membrane
resistance with higher voltage, leading to performance similar to that of their acidic counterparts. In addition, quinone-ferrocyanide alkalinechemistry avoids the membrane crossover, corrosivity, toxicity, and regulations associated withbromine. This reduced corrosivity can lead to a
substantially lower materials cost because many
components can be made of inexpensive polyolefin or poly(vinyl chloride) plastics.
RE FE RENCES AND N OT ES

1. B. Dunn, H. Kamath, J.-M. Tarascon, Science 334, 928–935(2011).
2. Z. Yang et al., Chem. Rev. 111, 3577–3613 (2011).
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4. D. Biello, Sci. Am. 311, 66–71 (November 2014).
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F. S. Mjalli, M. Saleem, J. Electrochem. Soc. 158, R55–R79(2011).
6. B. Huskinson et al., Nature 505, 195–198 (2014).
7. R. M. Darling, K. G. Gallagher, J. A. Kowalski, S. Ha,
F. R. Brushett, Energy Env. Sci. 7, 3459–3477(2014).
8. G. G. I. Joseph, A. J. Gotcher, G. Sikha, G. J. Wilson, Highperformance flow battery (2011); www.google.com/patents/
US20110244277.
9. J. R. Goldstein, Novel flow battery and usage thereof(2015); www.google.com/patents/US20150048777.
10. “Seventeenth Report of the Joint FAO/WHO Expert
Committee on Food Additives. Report No. 539,” Wld Hlth
Org. Techn. Rep. Ser. (World Health Organization,
Geneva, 1974).
11. H. Pal, T. Mukherjee, J. P. Mittal, J. Chem. Soc., Faraday Trans.
90, 711–716 (1994).
12. S. Er, C. Suh, M. P. Marshak, A. Aspuru-Guzik, Chem. Sci. 6,
885–893 (2015).
13. M. Quan, D. Sanchez, M. F. Wasylkiw, D. K. Smith, J. Am. Chem.
Soc. 129, 12847–12856 (2007).
14. Q. Chen, M. R. Gerhardt, L. Hartle, M. J. Aziz, J. Electrochem.
Soc. 163, A5010–A5013 (2015).
15. T. S. Balaban, A. Eichhöfer, M. J. Krische, J.-M. Lehn, Helv.
Chim. Acta 89, 333–351 (2006).
16. D. Linden, T. B. Reddy, Handbook of Batteries (McGraw-Hill,
New York, 2002).
ACKN OW LEDG MEN TS

This work was funded by the U.S. Department of Energy
Advanced Research Projects Agency–Energy award
no. DE-AR0000348 and the Harvard John A. Paulson
School of Engineering and Applied Sciences. Methods, alongwith any additional extended data display items andsource data, are available in the supplementary materials;
references unique to these sections appear only in the online
version of the paper. M.P.M., K.L., R.G.G., and M.J.A.
formulated the project. K.L., S.B.K., and D.H. synthesized,
analyzed, and purified the compounds. K.L., L.T., and S.B.K.
collected and analyzed the NMR and mass spectroscopy
data. K.L., A.W.V., and L.E. measured solubility. K.L., Q.C.,
L.E., and M.R.G. collected and analyzed the electrochemical

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data. K.L., Q.C., M.R.G., M.P.M., R.G.G., and M.J.A. wrote thepaper, and all authors contributed to revising the paper.

Supplementary Text
Fig. S1 to S9References (17–22)

SUPPLEMENTARY MATERIALS

www.sciencemag.org/content/349/6255/1529/suppl/DC1
Materials and Methods

7 April 2015; accepted 24 August 201510.1126/science.aab3033

H-BONDING CATALYSISO–H hydrogen bonding promotes
H-atom transfer from a C–H bonds
for C-alkylation of alcohols
Jenna L. Jeffrey,* Jack A. Terrett,* David W. C. MacMillan†
The efficiency and selectivity of hydrogen atom transfer from organic molecules are often
difficult to control in the presence of multiple potential hydrogen atom donors and acceptors.
Here, we describe the mechanistic evaluation of a mode of catalytic activation that
accomplishes the highly selective photoredox a-alkylation/lactonization of alcohols withmethyl acrylate via a hydrogen atom transfer mechanism. Our studies indicate a particularrole of tetra-n-butylammonium phosphate in enhancing the selectivity for a C–H bonds inalcohols in the presence of allylic, benzylic, a-C=O, and a-ether C–H bonds.

C

omplex molecules, such as medicinal agentsand natural products, often possess multiple types of C–H bonds, each with a different
inherent reactivity. This intrinsic reactivity depends on a multifaceted interplayof steric effects, inductive and conjugative influences, as well as innate strain (1, 2). The intermolecular catalytic functionalization of C(sp3)–H
bonds in a selective manner represents a longstanding challenge that has inspired decadesof effort within the synthetic community. Notable early studies by Bergman (3), as well as recent
advances in selective intermolecular transition
metal catalyzed C(sp3)–H activation—including,among others, Hartwig’s rhodium-catalyzed
borylation of terminal methyl groups (4) andWhite’s iron-catalyzed oxidation of both secondary (2°) and tertiary (3°) aliphatic C–H bonds(5)—highlight the importance of catalyst structure on site selectivity.
Catalyst structure has also proven critical to theselectivity of C(sp3)–H functionalization via hydrogen atom transfer (HAT) catalysis. HAT—the effective movement of a hydrogen atom between
two molecular sites—represents a ubiquitous elementary reaction step in organic chemistry (6–8).
The rate of hydrogen abstraction from a C–H bonddepends not only on the C–H bond dissociation
enthalpy (BDE) but also on polar effects in thetransition state. In 1987, Roberts noted that certain electrophilic radicals (such as t-butoxyl) preferentially abstract hydrogen from electron-rich
C–H bonds, whereas nucleophilic radicals (suchMerck Center for Catalysis, Princeton University, Princeton,
NJ 08544, USA.
*These authors contributed equally to this work. †Corresponding
author. E-mail: dmacmill@princeton.edu

as amine-boryl) selectively cleave electron-deficient
C–H bonds (9). The generality of this concept was
subsequently delineated through the broad application of polarity reversal catalysis (PRC), whichtakes advantage of favorable polar effects to control the regioselectivity of HAT from multiple C–H
groups of similar strength (10).
We questioned whether the basic principles of
PRC could be integrated into a catalytic system forthe selective activation of alcohol a-C–H bonds inthe presence of a wide range of other C–H bonds(such as a-C=O, a-ether, or allylic or benzylic C–H)
(11, 12). Specifically, we postulated that the selective C-alkylation of alcohols could be achieved
via a photoredox-catalyzed, H-bond–assisted bondactivation strategy (Fig. 1) (13–15), in which thehydroxyalkyl C–H bond is selectively polarized
and weakened via O–H hydrogen bonding.
It is well known that the strength of a C–H
bonds of alcohols decreases upon deprotonationof the alcohol O–H group. This so-called “oxyanionic substituent effect” (16, 17) leads to theacceleration of a wide range of organic reactions[such as oxyanionic [1,3] and [3,3] sigmatropic
rearrangements and HAT from alkoxides (18)].
More recently, it has been shown that intermolecular hydrogen bonding between alcohols andvarious acceptor molecules gives rise to a similar
polarization and weakening of the adjacent C–Hbond (19), the strength of which is reflected in the13
C nuclear magnetic resonance (NMR) chemical
shift and the one-bond 13C–1H coupling constant(1JCH) (20, 21). In particular, it was found that a
1 kJ/mol increase in the enthalpy of the H-bondresulted in a 0.2-Hz decrease in 1JCH for hexafluoroisopropanol complexed to various amines (20).
On the basis of these studies, we reasoned that theefficiency and selectivity of alcohol C–H activation
sciencemag.org SCIENCEAlkaline quinone flow battery
Kaixiang Lin et al.
Science 349, 1529 (2015);
DOI: 10.1126/science.aab3033

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Supplementary Materials for
Alkaline Quinone Flow Battery
Kaixiang Lin, Qing Chen, Michael R. Gerhardt, Liuchuan Tong, Sang Bok Kim, Louise
Eisenach, Alvaro Valle, David Hardee, Roy G. Gordon*, Michael J. Aziz*, & Michael P.
Marshak**

Correspondence to: aziz@seas.harvard.edu (M.J.A.); gordon@chemistry.harvard.edu(R.G.G.); michael.marshak@colorado.edu (MPM)

This PDF file includes:
Materials and Methods
Supplementary Text
Figs. S1 to S91

Materials and Methods
Materials
All chemicals were purchased from Sigma Aldrich and used as received unlessstated otherwise. 2,6-dihydroxyanthraquinone (2,6-DHAQ) (90% purity) was purchasedfrom AK Scientific Inc. and was recrystallized from DMF-water mixture. 2-methyl-3hydroxybenzoic acid (98% purity) was purchased from AK Scientific Inc. and was used
as received. Nafion 212 and Sigracet SGL 10AA porous carbon paper were purchasedfrom Ion Power and both pretreated (see method below). The cell hardware and Pocopyrosealed graphite flow plates were purchased from Fuel Cell Technologies.
Chemical characterization and preparation
Proton nuclear magnetic resonance (1H NMR) spectra were recorded using Varian
INOVA 500 (500 MHz) NMR spectrometers at 23 °C. Proton chemical shifts are
expressed in parts per million (ppm, δ scale) and are referenced to residual protium in the
NMR solvent (CDCl3, δ 7.26 ppm; (CD3)2CO, δ 2.05 ppm; (CD3)2SO, δ 2.50 ppm;
CD3OD, δ 3.31 ppm). Carbon nuclear magnetic resonance (13C NMR) spectra wererecorded using Varian INOVA 500 (125 MHz) NMR spectrometers at 23 °C. Carbon
chemical shifts are expressed in parts per million (ppm, δ scale) and are referenced to thecarbon resonances of the NMR solvent (CDCl3, δ 77.16 ppm; (CD3)2CO, 29.84 ppm;
(CD3)2SO, 39.50 ppm; CD3OD, 49.00 ppm).
Mass spectrum was collected using a maXis impact UHR time-of-flight mass
spectrometer system (Bruker Daltonics Inc, Billerica, MA) equipped with an electrosprayionization (ESI) source. The sample was dissolved in acetonitrile (5 uM) and analyzed innegative mode.
2,3,6,7-tetrahydroxyanthraquinone (2,3,6,7-THAQ) was prepared using a previously
reported synthetic route (15). The overall yield was 29% (1H and 13C NMR spectra shownin Fig. S8). 1,5-dimethyl-2,6-dihydroxyanthraquinone (1,5-DMAQ) was prepared by thefollowing method: AlCl3 (10 g) and NaCl (1 g) were mixed and heated to 200 oC to afford
a molten salt. 2-methyl-3-hydroxybenzoic acid (2 g) was slowly added in four portions
into the AlCl3-NaCl melt. The mixture was continuously heated at 200 oC for 1 hour.
After cool down, distilled water (100 mL) was slowly added to this solution, which wasthen stirred at room temperature overnight. The dark yellow precipitate was collected on
frit glass filter and extracted with ethyl acetate (100 mL). Distilled water (50 mL) was
then added to this solution. Ethyl acetate layer was separated from aqueous layer by
separatory funnel and dried on magnesium sulfate. The organic phase was concentrated(to approximately 25 mL) and cooled in an ice bath. The resulting precipitates were
filtered on celite and washed with cold ethyl acetate (50 mL). This filter cake was thendissolved in acetone (50 mL) where ethyl acetate (50 mL) was added. This mixture wasgently heated to remove acetone. Precipitates from this resulting solution were removed
on celite. After concentrating the ethyl acetate filtrate, hexane was added to precipitateout desired product which was then collected on frit glass filter and dried under vacuum
to afford pale yellow solid. The yield of this product was 12% (1H, 13C NMR spectra andmass spectroscopy results are shown in Fig. S9).
Solubility measurement

2Room temperature solubility in 1 M KOH solution was determined by measuring theabsorbance at 413 nm and comparing to an absorbance-vs.-concentration calibration
curve determined by preparing known concentrations of 2,6-DHAQ. UV-Vis
spectrophotometry measurements were performed using an Agilent Cary 60
spectrophotometer equipped with a Quantum Northwest T2 temperature regulator.
Appropriate aliquots of 2,6-DHAQ stock solution were added to 1 M KOH blank solution
and their UV-Vis absorbance spectra measured. A saturated solution of 2,6-DHAQ in 1
M KOH was prepared by adding 2,6-DHAQ potassium salt into 1 M KOH solution (10mL) until a thin layer of precipitate formed, the remaining solution was diluted by known
proportions, and the absorbance of the resulting solution was compared to the calibration
curve.
Electrochemical characterization
A glassy carbon electrode was used for three-electrode cyclic voltammetry tests,
except in the electrochemical window tests where graphite foil was used to mimic thecondition of porous carbon paper. Rotating-disk electrode (RDE) experiments of 2,6DHAQ (1 mM) in 1 M KOH solution were performed using a BASi RDE-2 rotating-disk
electrode system (Fig. S2A). All tests were carried out using a Gamry Reference 3000
potentiostat, with a Pt counter electrode and a Ag/AgCl reference electrode (equilibrated
with 3M NaCl, 213 mV vs. standard hydrogen electrode). A Levich plot was constructedfrom the RDE data by plotting the mass-transport limited current vs. the square root of
the rotation rate (Fig. S2B). The diffusion coefficient of 2,6-DHAQ was calculated foreach of three runs from the slope of the line fit to the Levich equation (17). The kinematic
viscosity was taken to be 1.08 × 10−6 m2 s−1 (18). The resulting value of the diffusion
coefficient is 4.8(2) × 10−6 cm2 s−1.
Membrane and electrode pretreatments
Pretreatment of Nafion ® 212 membrane was performed by first heating in 80 oC
de-ionized water for 20 minutes and then soaking in 5% hydrogen peroxide solution for
35 minutes. These pre-treated membranes were stored in 0.1 M KOH solution at room
temperature.
Sigracet ® SGL 10AA (~ 400 µm thickness) porous carbon paper was pretreated bybaking in air at 400 °C for 24 hours.
Full cell measurement
20 oC cycling study
The positive electrolyte was prepared by dissolving potassium ferrocyanide
trihydrate (2.56 g) in 1 M KOH solution (15 mL) to afford a 0.4 M ferrocyanide and 2.6
M potassium electrolyte solution. The negative electrolyte was prepared by dissolving
2,6-DHAQ (0.48 g) in 2 M KOH solution (4 mL) resulting a 0.5 M 2,6-DHAQ and 1 M
potassium electrolyte solution.
20 oC cell performance study
The positive electrolyte was prepared by dissolving potassium ferrocyanide
trihydrate (2.56 g) in 1 M KOH solution (15 mL) to afford a 0.4 M ferrocyanide and 2.6
M potassium electrolyte solution. The negative electrolyte was prepared by dissolving3

2,6-DHAQ (0.6 g) in 2 M KOH solution (5 mL) resulting a 0.5 M 2,6-DHAQ and 1 M
potassium electrolyte solution.
45 oC cell performance study
The positive electrolyte was prepared by heat dissolving potassium ferrocyanide
trihydrate (5.1 g) in 1 M KOH solution (15 mL) at 45 oC to afford a 0.8 M ferrocyanide
and 4.2 M potassium electrolyte solution. The negative electrolyte was prepared by heatdissolving 2,6-DHAQ (1.2 g) in 3 M KOH solution (5 mL) at 45 oC resulting a 1 M 2,6DHAQ and 3 M potassium electrolyte solution.
For all full cell studies, the electrolytes were assembled in the fully discharged state.
Cell hardware from Fuel Cell Tech. (Albuquerque, NM) was used to assemble azero-gap flow cell configuration, similar to previous reports (19). Serpertine flow patternpyrosealed POCO graphite flow plates were used for both sides. A 5 cm2 geometricsurface area electrode comprised a stack of three sheets of Sigracet SGL 10AA porous
carbon paper. A sheet of pretreated Nafion 212 membrane served as the ion-selective
membrane. The rest of the space between the plates was gasketed by Teflon sheets. The
electrolytes were fed into the cell through PFA tubing, at a rate of 60 mL/min controlled
by Cole-Parmer Micropump gear pumps.
Before and during the tests, the electrolytes were purged with ultra-high purity argon
to ensure deaeration. Galvanostatic cycling was performed at ±0.1 A/cm2, with voltage
limits of 0.6 and 1.7 V, controlled by a Gamry 30K Booster potentiostat. To obtain thepolarization curves, the cell was first charged to the desired SOC, and then polarized via
linear sweep voltammetry at a rate of 100 mV/s. This method was found to yieldpolarization 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 50% SOC, open-circuit potential, and 10 mA/cm2 AC current density, with frequency
ranging from 1 to 100,000 Hz.
Supplementary Text
Cyclic voltammogram modeling
Computation was performed using Mathematica 10.0.1.0 according to algorithms by
Oldham and Myland (20). The diffusivity of 2,6-DHAQ and both the one-electron andtwo-electron reduction products were assumed to be 4.8 × 10−6 cm2 s−1, based on our
RDE studies. The temperature was 293 K. Time was discretized into 40 ms increments,
and the current at each time was computed based on the sum of previous currents as
described by Oldham (20).
To model the two-electron reduction process (Fig. S2C), the Butler-Volmer equation
with the number of electrons n = 2 was used in place of Oldham equation (4:5), and theright-hand-side of equation (8:3) was multiplied by 2 to account for a two-electron
reduction (17, 20). The reduction potential was assumed to be equal to the voltage
equidistant from the two peaks on the measured cyclic voltammogram (−0.684 V vsSHE) and the charge transfer coefficient α was assumed to be 0.5. The electrochemicalrate constant k0 was then varied from 10−3 cm s−1 to 10−5 cm s−1.
To model the two successive one-electron reductions (Fig. S2D), Oldham equations(12:13) and (12:14) were used (20). The charge transfer coefficients α1 and α2 wereassumed to be equal to 0.5, and the electrochemical rate constants for the two reductions
were assumed equal. The values of the single rate constant and the first and second4reduction potentials were permitted to vary freely. A reasonable fit to the experimental
data was observed when the reduction potential of the first reduction E1 = −0.657 V vsSHE, the reduction potential of the second reduction E2 = −0.717 V vs SHE, and the twoelectrochemical rate constants k0,1 and k0,2 both equaled 7 × 10−3 cm s−1, a value very
close to that observed in other anthraquinone systems (6). The simulation was repeatedfor k0,1 and k0,2 equal to 7 × 10−4 cm s−1 to establish a lower bound on possible k0 values.
Energy Density
The negative electrolyte of 0.5 M DHAQ, with two electrons per molecule, has acharge capacity of 27 Ah/L. The positive electrolyte of 0.4 M ferrocyanide has a charge
capacity of 11 Ah/L. In the experimental cell used here, in which there is 50% morepositive electrolyte than needed to balance the negative electrolyte, the energy density is
6.8 Wh/L. However, eliminating the positive electrolyte volume excess and using
established methods of raising the ferrocyanide solubility to 1.4 M (21), the energy
density would become 18.8 Wh/L.

5

-600
Slo

-620

pe
m
61
=-

-640

V/p
H

Potential (mV vs. SHE)

-580

-660

Slope = 0 mV/pH

-680
-700
10

11

12

13

14

pH
Fig. S1. Pourbaix diagram of 2,6-DHAQ.
Above pH ~11.7, the equilibrium potential of 2,6-DHAQ is pH-independent, indicatingthat both oxidized and reduced form are fully deprotonated. The solid line is linear fit to
the data between pH 10 and 12; the dashed line is not a fit, but zero slope.

6

A

B

C

D

Fig. S2 Study and simulation of the electrochemistry of 2,6-DHAQ.
(A) Rotating disk electrode study of the reduction of a 1 mM solution of 2,6-DHAQ in 1
M KOH on a glassy carbon electrode at various rotation rates. (B) Levich plot of 1 mM
2,6-DHAQ in 1 M KOH. Data are an average of the current at −1.187 V vs SHE for eachof three runs; error bars indicate the standard deviation. (C) Cyclic voltammogram of 1
mM 2,6-dihydroxyanthraquinone in 1 M KOH (solid line). Dashed lines represent
simulated cyclic voltammograms of a two-electron reduction of varying electrochemicalrate constant k0 and a reduction potential E0 of −0.684 V vs SHE. The simulations
assumed α = 0.5. The scan rate in all cases was 25 mV s−1. (D) Cyclic voltammogram of
1 mM 2,6-dihydroxyanthraquinone in 1 M KOH (solid black line). Dashed red lines
represent the simulated components of two successive one-electron reductions with
reduction potentials of −0.657 V vs SHE and −0.717 V vs SHE. The solid red line
represents the simulated total current arising from such a reaction. Each simulated
reduction has a rate constant k0 = 7 × 10−3 cm s−1. The same simulation was performed
assuming each reduction has a rate constant k0 = 7 × 10−4 cm, and the total current plotted
as a solid green line. The slower rate constants do not adequately describe the observed
cyclic voltammogram, setting a lower bound on the value of the rate constant. The
simulations assumed α = 0.5.
7

A

B
B

C

Fig. S3 Chemical and electrochemical stability of 2,6-DHAQ.
1
H NMR (500 MHz, DMSO-d6) spectra. (A) 2,6-DHAQ δ: 8.04 (d, 2H, J = 8.3 Hz, 2 x
ArCH), 7.47 (d, 2H, J = 2.5 Hz, 2 x ArCH), 7.19 (dd, 2H, J = 8.3, 2.5 Hz, 2 x ArCH). (B)
2,6-DHAQ, after 30 days heating in 5 M KOH solution at 100 oC. (C) 2,6-DHAQ, after
100 charge-discharge cycles.

8

A
2

Current Density (mA/cm )

80
60
40
20
0
-20

2,6-DHAQ

-40
-60
-1.5

ferri/ferrocyanide

-1.0

-0.5

0.0

0.5

1.0

1.5

Potential (V vs. SHE)
B
4

2

Current Density (mA/cm )B

2

7.6 mM 2,6-DHAQ
added

0
-2
Cycled Posolyte

3.9 mM 2,6-DHAQ
added

-4
-6
-8
-1.0

-0.8

-0.6

-0.4

Potential (V vs. SHE)
Fig. S4 Test of electrolyte permeation through membrane.
(A) Full and (B) Zoomed in cyclic voltammograms of 0.4 M ferro-ferricyanide
electrolyte after 100 charge-discharge cycles (black trace). Same posolyte to which we
added 3.9 mM 2,6-DHAQ (orange trace), and 7.6 mM 2,6-DHAQ (pink trace).

9Fig. S5 Leakage of electrolyte into gaskets.
Image showing the Teflon gaskets and graphite flow plates (A) before and (B) after cell
cycling. Orange colored area indicates leakage of the negative electrolyte (top plate:
negative side; bottom plate: positive side) leading to capacity fade.

10

0.3
o

2

-Zimg(: cm )

0.2

0.1

o

2

2

20 C, rhf = 0.625 :cm

45 C, rhf = 0.419 :cm

1 Hz
1 kHz

1 Hz

1 kHz

0.0
10 Hz

10 Hz

-0.1

-0.2
0.4

0.5

0.6

2

Zreal(: cm )

0.7

0.8

Fig. S6 Electrochemical impedance spectroscopy (Nyquist plot).
Frequency decreases from left to right. The high frequency area-specific resistance (rhf)values discussed in the main text are obtained by fitting the high frequency parts to a
resistor (rhf) in series with an inductance component. The inductance came from thickcables and current collectors that connect the cell to the potentiostat. The rhf includes
contributions from the Nafion membrane, carbon electrodes, and electrical leads between
the cell and the potentiostat. The rhf values at two different temperatures are indicatedwithin the figure.

11

o

o

40

E = 0.40 V

E = -0.83 V

1 M KOH-2

Current Density (PA cm )

30
20
10
0
-10

1 M H2SO4

1 M KOH

-20
-30
o

o

E =0V

-40
-1.5

-1.0

-0.5

0.0

E = 1.23 V
0.5

1.0

1.5

Potential (V vs. SHE)

Fig. S7 Background cyclic voltammograms.
Cyclic voltammograms of 1 M KOH and 1 M H2SO4 background scanned at 100 mV/s
using graphite foil electrode. Dashed lines indicate commonly-reported equilibrium
potentials of water splitting reactions. This illustrates the practical stability window for
aqueous flow batteries.

12

A

B
B

C

Fig. S8 Synthetic scheme and NMR characterization of 2,3,6,7tetrahydroxyanthraquinone (2,3,6,7-THAQ).
(A) Synthetic scheme of 2,3,6,7-THAQ from inexpensive, commodity chemicals. i.
condensation with acetaldehyde to afford 2,3,6,7-tetramethoxydimethylanthracene(MeCHO, H2SO4), ii. oxidation to afford 2,3,6,7-tetramethoxyanthraquinone (Na2Cr2O7
in acetic acid), iii. hydrolysis of methoxy (HBr, reflux) We carried out all three steps ofreaction in this work according to literature reference (15). (B) 1HNMR (500 MHz,DMSO) spectrum of 2,3,6,7-THAQ δ 10.42 (br, 4H, 4 x Ar-OH), 7.43 (s, 4H, 4 x ArCH).
(C) 13CNMR (125 MHz) spectrum δ 181.50, 151.23, 127.46, 113.31, 113.28. Solvent
peaks are labeled with asterisks.

13

A
A

B

C

Fig. S9 Synthetic scheme and NMR characterization of 1,5-dimethyl-2,6dihydroxyanthraquinone (1,5-DMAQ).
(A) Synthetic scheme of 1,5-DMAQ from inexpensive, commodity chemicals. i.
sulfonation of naphthalene to afford 1,3,5-nanphthalenetrisulfonic acid (H2SO4/SO3), ii.
preparation of sodium salt to afford trisodium naphthalene-1,3,5-trisulfonate (NaOH), iii.
reaction of trisodium naphthalene-1,3,5-trisulfonate with alkali (NaOH in autoclave athigh temperature (280-310 °C)), iv. Dimerization to afford 1,5-DMAQ (AlCl3-NaClmelt). We carried out step iv in this work. Steps i-iii were reported in a patent application(22). (B) 1HNMR (500 MHz, DMSO-d6) spectra of 1,5-DMAQ δ: 10.78 (br, 2H, 2 x ArOH), 7.93 (d, 2H, J = 8.5 Hz, 2 x ArCH), 7.21 (d, 2H, J = 8.5 Hz, 2 x ArCH), 2.55 (s,6H, 2 x Ar-CH3). (C) 13CNMR (125 MHz, DMSO-d6) δ 185.63, 161.87, 133.10, 128.94,
128.04, 126.74, 119.83, 14.03. Solvent peaks are labeled by asterisk. Mass spectroscopy:
C6H12O4 M-H calculated 267.0657 found 267.0659

14

References, Supplementary Materials

17. A. J. Bard, L. R. Faulkner, Electrochemical Methods: Fundamentals and Applications, 2nded. (John Wiley & Sons, New York, 2001)
18. L. B. Hitchcock, J. S. McIlhenny, Viscosity and Density of Pure Alkaline Solutions andTheir Mixtures. Ind. Eng. Chem. 27, 461–466 (1935).
19. Q. H. Liu et al., High Performance Vanadium Redox Flow Batteries with Optimized
Electrode Configuration and Membrane Selection. J. Electrochem. Soc. 159, A1246–A1252(2012).
20. K. B. Oldham, J. C. Myland, Modelling cyclic voltammetry without digital simulation.
Electrochimica Acta. 56, 10612–10625 (2011).
21. A. J. Esswein, J. Goeltz, D. Amadeo, High solubility iron hexacyanides (2014);
www.google.com/patents/US20140051003
22. H. Behre, F. Mueller-Hauck, J. Scherer, G. Schroeder, Method for producing 3-hydroxy-2methylbenzoic acid (2004); www.google.com/patents/WO2003080542A3