LETTER

doi:10.1038/nature12909

A metal-free organic––inorganic aqueous flow battery
Brian Huskinson1*, Michael P. Marshak1,2*, Changwon Suh2, Süleyman Er2,3, Michael R. Gerhardt1, Cooper J. Galvin2,
Xudong Chen2, Alán Aspuru-Guzik2, Roy G. Gordon1,2 & Michael J. Aziz1

As the fraction of electricity generation from intermittent renewable sources—such as solar or wind—grows, the ability to store large
amounts of electrical energy is of increasing importance. Solid-electrodebatteries maintain discharge at peak power for far too short a time tofully regulate wind or solar power output1,2. In contrast, flow batteries
can independently scale the power (electrode area) and energy (arbitrarily large storage volume) components of the system by maintaining
all of the electro-active species in fluid form3–5. Wide-scale utilization of flow batteries is, however, limited by the abundance and costof these materials, particularly those using redox-active metals andprecious-metal electrocatalysts6,7. Here we describe a class of energy
storage materials that exploits the favourable chemical and electrochemical properties of a family of molecules known as quinones.
The example we demonstrate is a metal-free flow battery based on theredox chemistry of 9,10-anthraquinone-2,7-disulphonic acid (AQDS).
AQDS undergoes extremely rapid and reversible two-electron twoproton reduction on a glassy carbon electrode in sulphuric acid. An
aqueous flow battery with inexpensive carbon electrodes, combining the quinone/hydroquinone couple with the Br2/Br2 redox couple, yields a peak galvanic power density exceeding 0.6 W cm22 at
1.3 A cm22. Cycling of this quinone–bromide flow battery showed.99 per cent storage capacity retention per cycle. The organic anthraquinone species can be synthesized from inexpensive commodity
chemicals8. This organic approach permits tuning of important properties such as the reduction potential and solubility by adding functional groups: for example, we demonstrate that the addition of twohydroxy groups to AQDS increases the open circuit potential of thecell by 11% and we describe a pathway for further increases in cellvoltage. The use of p-aromatic redox-active organic molecules insteadof redox-active metals represents a new and promising direction forrealizing massive electrical energy storage at greatly reduced cost.
Solutions of AQDS in sulphuric acid (negative side) and Br2 in HBr(positive side) were pumped through a flow cell as shown schematicallyin Fig. 1a. The quinone–bromide flow battery (QBFB) was constructedusing a Nafion 212 membrane sandwiched between Toray carbon paper
electrodes (six stacked on each side) with no catalysts; it is similar to acell described elsewhere (see figure 2 in ref. 7). We report the potential–
current response (Fig. 1b) and the potential–power relationship (Fig. 1cand d) for various states of charge (SOCs; measured with respect tothe quinone side of the cell). As the SOC increased from 10% to 90%,
the open-circuit potential increased linearly from 0.69 V to 0.92 V. In
the galvanic direction, peak power densities were 0.246 W cm22 and
0.600 W cm22 at these same SOCs, respectively (Fig. 1c). To avoid
significant water splitting in the electrolytic direction, we used a cutoff voltage of 1.5 V, at which point the current densities observed
at 10% and 90% SOCs were 22.25 A cm22 and 20.95 A cm22, respectively, with corresponding power densities of 23.342 W cm22 and21.414 W cm22.
In Fig. 2 we report the results of initial cycling studies for this battery,
to test for consistent performance over longer timescales. Figure 2ashows cycling data at 60.2 A cm22 using 50% of the total capacity ofthe battery. The cycles are highly reproducible and indicate that current
efficiencies for the battery are around 95%. Figure 2b shows constantcurrent cycling data, collected at 60.5 A cm22, using voltage cut-offs of
0 V and 1.5 V. These tests were done using the identical solutions usedin the battery for Fig. 1b–d. The galvanic discharge capacity retention(that is, the number of coulombs extracted in one cycle divided by the
number of coulombs extracted in the previous cycle) is above 99%,
indicating the battery is capable of operating with minimal capacity
fade and suggesting that current efficiencies are actually closer to 99%.
Full characterization of the current efficiency will require slower cycling experiments and chemical characterization of the electrolyte solutions after extended cycling.
To gain a better understanding of the quinone half-reaction on
carbon, AQDS was subjected to half-cell electrochemical measurements. Cyclic voltammetry of a 1 mM solution of AQDS in 1 M sulphuricacid on a glassy carbon disk working electrode shows current peaks corresponding to reduction and oxidation of the anthraquinone species9–11(Fig. 3d, solid trace). The peak separation of 34 mV is close to the valueof 59 mV/n, where n is the number of electrons involved, expected fora two-electron process. Rotation of this disk at a variety of rates yields
mass-transport-limited currents (Fig. 3a) from which the AQDS diffusion coefficient (D 5 3.8(1)3 1026 cm2 s21) can be determined; throughout this paper, the numbers reported in parentheses indicate thestandard deviation in the last reported digit. Koutecký–Levich analysis
at low overpotentials (Fig. 3b) can be extrapolated to infinite rotation
rate and fitted to the Butler–Volmer equation (Extended Data Fig. 3a) togive the kinetic reduction rate constant k0 5 7.2(5) 3 1023 cm s21. Thisrate constant is greater than that found for other species used in flow
batteries such as V31/V21, Br2/Br2 and S422/S222 (see table 2 in ref. 3).
It implies that the voltage loss due to the rate of surface electrochemical
reactions is negligible. The high rate is apparently due to an outersphere two-electron reduction into the aromatic p system requiring
little reorganizational energy. The electrochemical reversibility of thetwo-electron redox reaction was confirmed by fitting the slope to the
Butler–Volmer equation (Extended Data Fig. 3a), giving the transfer
coefficient a 5 0.474(2), which is close to the value of 0.5 expected for
an ideally reversible reaction. The Pourbaix diagram (Extended Data
Fig. 4) confirms that a two-electron, two-proton reduction occurs inacidic solution, and yields approximate pKa values of 7 and 11 for the
reduced AQDS species11.
Functionalization of the anthraquinone backbone with electrondonating groups such as hydroxy can be expected lower the reduction
potential of AQDS (E0), thereby raising the cell voltage12. Hydroxysubstituted anthraquinones are synthesized through oxidation reactions that may be performed at minimal cost. They are also natural
products that have been extracted for millennia from common sources
such as rhubarb and could even provide a renewable source for futureanthraquinone-based electrolyte solutions.
Quantum chemical calculations of un-substituted and hydroxysubstituted AQDS were performed to predict how substitution patterns
would change both E0 of the quinone/hydroquinone couples (Fig. 3c)

1

Harvard School of Engineering and Applied Sciences, 29 Oxford Street, Cambridge, Massachusetts 02138, USA. 2Department of Chemistry and Chemical Biology, Harvard University, 12 Oxford Street,
Cambridge, Massachusetts 02138, USA. 3Molecular Materials and Nanosystems, Eindhoven University of Technology, PO Box 513, 5600 MB Eindhoven, The Netherlands.
*These authors contributed equally to this work.
9 J A N U A RY 2 0 1 4 | VO L 5 0 5 | N AT U R E | 1 9 5©2014 Macmillan Publishers Limited. All rights reservedLoad(source)

a
Pump

b

25%
50%

10% SOCAQDSH2

Br2

HBr/Br2
electrolyte
storage
tank

2 H+
2 HBr

AQDS

Cell potential (V)

1.2

AQDS/
AQDSH2electrolyte
storagetank

75% 90%

1.4

2 e–

2 e–

1.6

Pump

Equilibrium potential (V)RESEARCH LETTER
1.00
0.90
0.80
0.70
0.60
0

1.0

20 40 60 80 100
State of charge (%)

0.8
0.6
0.4

Porous carbon electrodes

0.2
Charge0.0–2.4 –2.0 –1.6 –1.2 –0.8 –0.4

Proton exchange membrane

Discharge
0.0

0.4

0.8

1.2

1.6

Current density (A cm–2)

d

0.7
90%
75%

Galvanic power density (W cm–2)

0.6

Electrolytic power density (W cm–2)

c

0.5
50%
0.4
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10% SOC
0.2
0.1
0.0
0.0

0.2

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1.0
Current density (A cm–2)

1.2

1.4

0.0
–0.5
–1.0
90%

–1.5
75%

–2.0
50%
–2.5
–3.0

25%
10% SOC

–3.5
–2.5

1.6

–2.0

–1.5

–1.0

-0.5

0.0

Current density (A cm–2)

Figure 1 | Cell schematic and cell performance in galvanic and electrolytic
modes. a, Cell schematic. Discharge mode is shown; the arrows are reversed
for electrolytic/charge mode. AQDSH2 refers to the reduced form of AQDS.
b, Cell potential versus current density at five different states of charge(SOCs; average of three runs); inset shows the cell open circuit potential

versus SOC with best-fit line superimposed (Eeq 5 (0.00268 3 SOC) 1 0.670;
R2 5 0.998). c, Galvanic power density versus current density for the same
SOCs. d, Electrolytic power density versus current density. All data here werecollected at 40 uC using a 3 M HBr 1 0.5 M Br2 solution on the positive side anda 1 M AQDS 1 1 M H2SO4 solution on the negative side.

and the solvation free energy (G0solv ) in aqueous solution (Extended
Data Table 1). The addition of 2OH groups is calculated to lower theE0 by an average of 250 mV per 2OH and provide a wide window fortuning E0 by almost 0.6 V. In addition, increasing numbers of hydroxy
substituents are expected to raise the aqueous solubility due to hydrogen bonding.

In confirmation of the theory, the experimental reduction potentialof 1,8-dihydroxy-9,10-anthraquinone-2,7-disulphonic acid (DHAQDS)
was found to be 118 mV (versus the standard hydrogen electrode),which is very close to the 101 mV calculated for this species (Fig. 3cand d). The experimental E0 of DHAQDS was 95 mV lower than
AQDS, and would result in an 11% increase in QBFB cell potential.

ab

Cycle number
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15

1.2

100

1.8

± 0.2 A cm–2

95

Cell potential (V)

90

0.9
0.8

85

0.7

80

0.6
75

0.5
0.4

Current efficiency (%)

1.0

0

2

4

6

8

10

12

14

16

18

20

70

Time (h)Figure 2 | Cell cycling behaviour. a, Constant-current cycling at 0.2 A cm22
at 40 uC using a 2 M HBr 1 0.5 M Br2 solution on the positive side and a 0.1 M
AQDS 1 2 M H2SO4 solution on the negative side; current efficiency is
indicated for each complete cycle. b, Constant-current cycling at 0.5 A cm22

Cell potential (V)

1.5

1

2

3

4

5

6

b

7

8

9

99.2% average

10
100
99

± 0.5 A cm–2

98
97

1.2

96
0.9

95
94

0.6

93
0.3

92

0.0

91

0

10

20

30

40

50

60

70

80

Discharge capacity retention (%)

1.1

Cycle number
090
90 100Time (h)at 40 uC using a 3 M HBr 1 0.5 M Br2 solution on the positive side and a
1 M AQDS 1 1 M H2SO4 solution on the negative side (same solution usedin Fig. 1); discharge capacity retention is indicated for each cycle.

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LETTER RESEARCH
b –60

0
–200

–50

200 r.p.m.

–400

–40
i –1 (mA–1)

Current density (μA cm–2)

a

–600
–800

–30

η (mV)
13
18
23
28
33
38
363

–20
–1,000–10–1,2003,600 r.p.m.
100

150
200
250
Potential (mV vs SHE)cd6
Number of −OH groups

0
0.00

300

5
4
3
DHAQDS

2
1

AQDS

Current density (μA cm–2)–1,400
50

0.04

0.08
0.12
0.16
0.20
(Rotation rate)–1/2 (s1/2 rad–1/2)

0.24

200
100
0
O
HO3S

–100

SO3H

AQDS

–200

O
OH O

HO3S

OH
SO3H

–300
DHAQDS

0O–300 –200 –100

0

100

200

300

Potential (mV vs SHE)

–100

0

100

200

300

400

500

Potential (mV vs SHE)

Figure 3 | Half-cell measurements and theory calculations. a, Rotating
disk electrode (RDE) measurements of AQDS using a glassy carbon electrode in
1 M H2SO4 at 11 rotation rates ranging from 200 r.p.m. (red) to 3,600 r.p.m.
(black). b, Koutecký–Levich plot (current21 versus rotation rate21/2) derived
from a at seven different AQDS reduction overpotentials, g. c, Calculated

reduction potentials of AQDS substituted with –OH groups (black), calculated
AQDS and DHAQDS values (blue), and experimental values for AQDS andDHAQDS (red squares). d, Cyclic voltammogram of AQDS and DHAQDS(1 mM) in 1 M H2SO4 on a glassy carbon electrode (scan rate 5 25 mV s21).

DHAQDS was also found to have faster reduction kinetics (k0 5 1.56(5)3 1022 cm s21), possibly due to intramolecular hydrogen bonding of
the 2OH to the ketone (Extended Data Fig. 3b).
The organic approach liberates battery redox chemistry from the constraints of the limited number of elemental redox couples of the periodic table. Although quinones have been used previously in batteries
using redox-active solids13–15, their incorporation into all-liquid flow
batteries offers the following advantages over current flow-battery technologies. First, scalability: AQDS contains only the Earth-abundant atoms
carbon, sulphur, hydrogen and oxygen, and can be inexpensively manufactured on large scales. Because some hydroxy-anthraquinones arenatural products, there is also the possibility that the electrolyte material
can be renewably sourced. Second, kinetics: quinones undergo extremely rapid two-electron redox on simple, inexpensive carbon electrodes and do not require a costly precious-metal catalyst. Furthermore,this electrode permits higher charging voltages by suppressing the parasitic water-splitting reactions. Third, stability: quinones should exhibit
minimal membrane crossover owing to their relatively large size andcharge in aqueous solution as a sulphonate anion. Furthermore, althoughbromine crossover is a known issue in zinc-bromine, vanadium-bromine and hydrogen-bromine cells, AQDS is stable to prolonged heating inconcentrated Br2/HBr mixtures (Extended Data Figs 5 and 6), and the
QBFB can be cycled in HBr electrolyte solutions (Extended Data Fig. 9).
Fourth, solubility: AQDS has an aqueous solubility greater than 1 M at
pH 0, and the quinone solution can thus be stored at relatively high energy
density—volumetric and gravimetric energy densities exceed 50 W h l21
and 50 W h kg21, respectively. Last, tunability: the reduction potential
and solubility of AQDS can be further optimized by introduction offunctional groups such as 2OH. Use of DHAQDS is expected to leadto an increase in cell potential, performance and energy density.

These features lower the capital cost of storage chemicals per kilowatt
hour, which sets a floor on the ultimate system cost per kilowatt hourat any scale. The precursor molecule anthracene is abundant in crude
petroleum and is already oxidized on large scale to anthraquinone. Sulphonated anthraquinones are used on an industrial scale in wood pulp
processing for paper16, and they can be readily synthesized from thecommodity chemicals anthraquinone and oleum8. In fact, a cyclic voltammogram of the crude sulphonation product of these two reagents isvirtually identical to that of pure AQDS (Extended Data Fig. 8). Basedon this simple electrolyte preparation that requires no further productseparation, we estimate chemical costs of $21 per kilowatt hour for AQDS
and $6 per kilowatt hour for bromine17 (see Methods for informationon cost calculations). The QBFB offers major cost improvements overvanadium flow batteries with redox-active materials that cost $81 perkilowatt hour (ref. 18). Optimization of engineering and operating parameters such as the flow field geometry, electrode design, membrane separator and temperature—which have not yet even begun—should leadto significant performance improvements in the future, as it has forvanadium flow batteries, which took many years to reach the power
densities we report here6. The use of redox processes in p-aromaticorganic molecules represents a new and promising direction for costeffective, large-scale energy storage.

METHODS SUMMARYThe QBFB comprised a mixture of commercially available and custom-madecomponents. Pretreated 2 cm2, stacked (36) Toray carbon paper electrodes (eachof which is about 7.5 mm uncompressed) were used on both sides of the cell. Nafion
212 (50 mm thick) was used as a proton-exchange membrane, and PTFE gasketing
was used to seal the cell assembly. On the positive side of the cell, 120 ml of 3 M
HBr and 0.5 M Br2 were used as the electrolyte solution in the fully discharged state;
9 J A N U A RY 2 0 1 4 | VO L 5 0 5 | N AT U R E | 1 9 7©2014 Macmillan Publishers Limited. All rights reserved

RESEARCH LETTER
on the negative side, 1 M 2,7-AQDS in 1 M H2SO4 was used. AQDS disodium salt
was flushed twice through a column containing Amberlyst 15H ion-exchange resin
to remove the sodium ions. Half-cell measurements were conducted using a Ag/
AgCl aqueous reference electrode (3 M KCl filling solution), a Pt wire counterelectrode and a 3-mm-diameter glassy carbon disk electrode. For theoretical calculations, the total free energies of molecules were obtained from first-principles
quantum chemical calculations within density functional theory at the level ofthe generalized gradient approximation (GGA) using the PBE functional. Threedimensional conformer structures for each quinone/hydroquinone molecule were generatedusing the ChemAxon suite with up to 25 generated conformers per moleculeusing the Dreiding force field. Generated conformers were used as input structures
for the density functional theory geometry optimization employed for determining
the formation energy, which in turn is used to evaluate the reduction potential. In
the QBFB cost calculation, a price of $4.74 per kilogram (eBioChem) was used foranthraquinone. To get the sulphonated form actually used here, anthraquinone
must be reacted with oleum (H2SO4/SO3), which adds a negligible cost at scale; thiscost is not included here. The price of bromine was $1.76 per kilogram, based on
estimates from the US Geological Survey17. The cell voltage used to calculate costs
here was 0.858 V.
Online Content Any additional Methods, Extended Data display items and Source
Data are available in the online version of the paper; references unique to these
sections appear only in the online paper.
Received 26 June; accepted 25 November 2013.
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3577–3613 (2011).
Weber, A. Z. et al. Redox flow batteries: a review. J. Appl. Electrochem. 41,
1137–1164 (2011).
Leung, P. et al. Progress in redox flow batteries, remaining challenges and theirapplications in energy storage. RSC Adv. 2, 10125–10156 (2012).
Nguyen, T. & Savinell, R. F. Flow batteries. Electrochem. Soc. Interface 19, 54–56(2010).
Skyllas-Kazacos, M., Chakrabarti, M. H., Hajimolana, S. A., Mjalli, F. S. & Saleem, M.
Progress in flow battery research and development. J. Electrochem. Soc. 158,
R55–R79 (2011).
Huskinson, B., Rugolo, J., Mondal, S. K. & Aziz, M. J. A high power density, highefficiency hydrogen–chlorine regenerative fuel cell with a low precious metal
content catalyst. Energy Environ. Sci. 5, 8690–8698 (2012).
Crossley, M. L. The separation of mono-b, 2,6- and 2,7-sulfonic acids ofanthraquinone. J. Am. Chem. Soc. 37, 2178–2181 (1915).
Conant, J. B., Kahn, H. M., Fieser, L. F. & Kurtz, S. S. An electrochemical study of thereversible reduction of organic compounds. J. Am. Chem. Soc. 44, 1382–1396 (1922).

10. Kelsall, G. H. & Thompson, I. Redox chemistry of H2S oxidation by the British Gas
Stretford process. Part III: Electrochemical behaviour of anthraquinone 2,7disulphonate in alkaline electrolytes. J. Appl. Electrochem. 23, 296–307 (1993).
11. Forster, R. J. & O’Kelly, J. P. Protonation reactions of anthraquinone-2,7disulphonic acid in solution and within monolayers. J. Electroanal. Chem. 498,
127–135 (2001).
12. Song, Y. & Buettner, G. R. Thermodynamic and kinetic considerations for thereaction of semiquinone radicals to form superoxide and hydrogen peroxide. Free
Radic. Biol. Med. 49, 919–962 (2010).
13. Xu, Y. et al. in World Non-Grid-Connected Wind Power and Energy Conference 1–4(IEEE, 2009).
14. Wang, W. et al. Anthraquinone with tailored structure for a nonaqueous metalorganic redox flow battery. Chem. Commun. 48, 6669–6671 (2012).
15. Yao, M. et al. High capacity organic positive-electrode material based on abenzoquinone derivative for use in rechargeable lithium batteries. J. Power Sources
195, 8336–8340 (2010).
16. Gordon, O. W., Plattner, E. & Doppenberg, F. Production of pulp by the sodaanthraquinone process (SAP) with recovery of the cooking chemicals. US patent
5,595,628 (1997).
17. US Geological Survey 2010 Minerals Yearbook: Bromine (2010); available at http://
minerals.usgs.gov/minerals/pubs/commodity/bromine/myb1-2010-bromi.pdf.
18. US Geological Survey Vanadium Mineral Commodities Summary (2013); available
at http://minerals.usgs.gov/minerals/pubs/commodity/vanadium/mcs-2013vanad.pdf.
Acknowledgements This work was partially funded through US Department ofEnergy ARPA-E Award DE-AR0000348 and partially funded through the Harvard
School of Engineering and Applied Sciences. Theoretical work was funded in partthrough the Extreme Science and Engineering Discovery Environment (XSEDE),
which is supported by National Science Foundation grant number OCI-1053575.
B.H. was supported by an NSF Graduate Research Fellowship. S.E. performed workas part of the Fellowships for Young Energy Scientists programme of the Foundation for
Fundamental Research on Matter (FOM), which is part of the Netherlands Organizationfor Scientific Research (NWO). We thank T. Betley, L. Hartle, R. Burton and R. Duncan for
discussions.
Author Contributions B.H. and M.P.M. contributed equally to this work. B.H. and M.P.M.
designed and tested the battery, with direction from M.J.A. Both M.P.M. and M.R.G.
conducted electrochemistry experiments, with direction from M.J.A. M.P.M. and C.J.G.
synthesized chemicals with direction from R.G.G. Theoretical calculations were done by
C.S. and S.E., with input from M.P.M. and R.G.G. and direction from A.A.-G. X.C.
contributed NMR results. B.H., M.P.M., C.S., M.R.G., S.E., A.A.G., R.G.G. and M.J.A. allcontributed to writing the manuscript.
Author Information Reprints and permissions information is available atwww.nature.com/reprints. The authors declare competing financial interests: details
are available in the online version of the paper. Readers are welcome to comment on
the online version of the paper. Correspondence and requests for materials should be
addressed to M.J.A. (maziz@harvard.edu).

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LETTER RESEARCHMETHODS
Full cell measurements. The QBFB comprised a mixture of commercially available and custom-made components. Circular endplates were machined out of solid
aluminium. Current collectors were 3 inch 3 3 inch pyrolytic graphite blocks withinterdigitated flow channels (channel width 5 0.0625 inch, channel depth 5 0.08 inch,
landing between channels 5 0.031 inch, Fuel Cell Technologies). Pretreated 2 cm2,
stacked (36) Toray carbon paper electrodes (each of which is about 7.5 mm uncompressed) were used on both sides of the cell. Pretreatment consisted of a 10 min
sonication in isopropyl alcohol followed by a five hour soak in a hot (50 uC) mixtureof undiluted sulphuric and nitric acids in a 3:1 volumetric ratio. Nafion 212 (50 mm
thick) was used as a proton-exchange membrane (PEM, Alfa Aesar), and PTFE
gasketing was used to seal the cell assembly. Membrane pretreatment was done
according to previously published protocols7. Six bolts (3/8’’, 16 threads per inch)
torqued to 10.2 N m completed the cell assembly, and PTFE tubing was used totransport reactants and products into and out of the cell. The cell was kept on a hot
plate and wrapped in a proportional integral derivative (PID)-controlled heating
element for temperature control. On the positive side of the cell, 120 ml of 3 M HBr
and 0.5 M Br2 were used as the electrolyte solution in the fully discharged state; on
the negative side, 1 M AQDS in 1 M H2SO4 was used. HBr was used on the negativeside instead of H2SO4 for stability testing results displayed in Extended Data Fig. 9.
State-of-charge calculations are based on the composition of the quinone side of
the cell. 2,7-Anthraquinone disulphonate disodium salt 98% (TCI) was flushed
twice through a column containing Amberlyst 15H ion-exchange resin to remove
the sodium ions. Measurements shown here were done at 40 uC. March centrifugal
pumps were used to circulate the fluids at a rate of approximately 200 ml min21.
For characterization, several instruments were used: a CH Instruments 1100C potentiostat (which can be used up to 62 A), a DC electronic load (Circuit Specialists)
for galvanic discharge, a DC regulated power supply (Circuit Specialists) for electrolytic characterization, and a standard multimeter for independent voltage measurements. The cell was charged at 1.5 V until a fixed amount of charge ran throughthe cell. During this process, the electrolyte colours changed from orange to darkgreen (AQDS to AQDSH2) and from colourless to red (Br2 to Br2). Periodically,
the open circuit potential was measured, providing the data inset in Fig. 1b. Also, atvarious SOCs, potential–current behaviour was characterized using the equipment
described above: a fixed current was drawn from the cell, and the voltage, once
stabilized, was recorded (Fig. 1b). For the cycling data in Fig. 2b, the potentiostat
was used for constant-current (60.5 A cm22) measurements with cut-off voltagesof 0 V and 1.5 V. For the cycling data in Fig. 2a, a more dilute quinone solution(0.1 M as opposed to 1 M) was used. Here, the half-cycle lengths were programmed
to run at constant current for a fixed amount of time, provided the voltage cut-offs
were not reached, so that half of the capacity of the battery was used in each cycle.
The voltage cut-offs were never reached during charging, but were reached during
discharge. Current efficiencies are evaluated by dividing the discharge time by the
charge time of the previous half-cycle.
As shown in Fig. 2, current efficiency starts at about 92% and climbs to about
95% over ,15 standard cycles. Note that these measurements are done near viable
operating current densities for a battery of this kind. Because of this, we believe thisnumber places an upper bound on the irreversible losses in the cell. In any case,
95% is comparable to values seen for other battery systems. For example, ref. 19
reports vanadium bromide batteries with current efficiencies of 50–90%, with largechanges in current efficiency observed for varying membrane conditions. Our system
will probably be less dependent on membrane conditions because we are storingenergy in anions and neutral species as opposed to cations, which Nafion canconduct reasonably well.
In Fig. 2b we illustrate the capacity retention of the battery (that is, the numberof coulombs available for discharge at the nth cycle divided by that available fordischarge at the (n 2 1)th cycle) to be 99.2% on average, which is quite high andprovides direct evidence that our irreversible losses are below 1%. If we attribute allof this loss (the 0.78% capacity fade per cycle) to some loss of redox-active quinone,
it would be equivalent to losing 0.0006634 moles of quinone per cycle. If we attributeall of the loss to bromine crossover (which would react with the hydroquinone andoxidize it back to quinone), this corresponds to a crossover current density of
1.785 mA cm22, which is within the range of the widely varying crossover values
reported in the literature20. Note that these crossover numbers can be very sensitive
to membrane pretreatment conditions. It is also important to mention that, todetermine very accurate current efficiencies, detailed chemical analyses of theelectrolyte are necessary.
Half-cell measurements. These were conducted using a BASi Epsilon EC potentiostat, a BASi Ag/AgCl aqueous reference electrode (RE-5B, 3 M KCl filling solution) and a Pt wire counter electrode. Rotating disk electrode (RDE) measurements
were conducted using a BASi RDE (RDE-2) and a 3 mm diameter glassy carbon
disk electrode. Electrode potentials were converted to the standard hydrogen electrode (SHE) scale using E(SHE) 5 E(Ag/AgCl) 1 0.210 V, where E(SHE) is thepotential versus SHE and E(Ag/AgCl) is the measured potential versus Ag/AgCl.
2,7-Anthraquinone disulphonate disodium salt 98% was purchased from TCI andused as received. 1,8-Dihydroxy-anthraquinone-2,7-disulphonic acid was madeaccording to the literature procedure21. The electrolyte solution was sulphuric acid
(ACS, Sigma) in deionized H2O (18.2 MV cm, Millipore). The Pourbaix diagram(plot of E0 versus pH) shown in Extended Data Fig. 4 was constructed using aqueous
1 mM solutions of AQDS in pH buffers using the following chemicals: sulphuricacid (1 M, pH 0), HSO42/SO422 (0.1 M, pH 1–2), AcOH/AcO2 (0.1 M, pH 2.65–5),
H2PO42/HPO422 (0.1 M, pH 5.3–8), HPO422/PO432 (0.1 M, pH 9.28–11.52), and
KOH (0.1 M, pH 13). The pH of each solution was adjusted with 1 M H2SO4 or
0.1 M KOH solutions and measured with an Oakton pH 11 Series pH meter(Eutech Instruments).
RDE studies. All RDE data represent an average of three runs. Error bars in Extended
Data Figs 2 and 3 indicate standard deviations. Before each run, the glassy carbon
disk working electrode was polished to a mirror shine with 0.05 mm alumina andrinsed with deionized water until a cyclic voltammogram of a solution of 1 mM
AQDS in 1 M H2SO4 showed anodic and cathodic peak voltage separation of 34 to
35 mV (39 mV for DHAQDS) at a sweep rate of 25 mV s21. The electrode was thenrotated at 200, 300, 400, 500, 700, 900, 1,200, 1,600, 2,000, 2,500 and 3,600 r.p.m.
while the voltage was linearly swept from 310 to 60 mV (250 to 2100 for DHAQDS)
at 10 mV s21 (Extended Data Fig. 1). The currents measured at 60 mV (2100 for
DHAQDS) (that is, the diffusion-limited current density) versus the square root of
the rotation rate (v) is plotted in Extended Data Fig. 2. The data were fitted with astraight line, with the slope defined by the Levich equation as 0.620nFACOD2/3n21/6,
where n 5 2, Faraday’s constant F 5 96,485 C mol21, electrode area A 5 0.0707 cm2,
AQDS concentration CO 5 1026 mol cm23, kinematic viscosity n 5 0.01 cm2 s21.
This gives D values of 3.8(1) 3 1026 cm2 s21 for AQDS and 3.19(7) 3 1026 cm2 s21
for DHAQDS. The reciprocal of the current at overpotentials of 13, 18, 23, 28, 33,
38 and 363 mV was plotted versus the reciprocal of the square root of the rotation
rate (Fig. 3b and Extended Data Fig. 2). The data for each potential were fitted witha straight line; the intercept gives the reciprocal of iK, the current in the absence ofmass transport limitations (the extrapolation to infinite rotation rate). A plot of
log10(iK) versus overpotential was linearly fitted with a slope of 62 mV (AQDS) and
68 mV (DHAQDS) defined by the Butler–Volmer equation as 2.3aRT/nF (Extended
Data Fig. 3), where R is the universal gas constant, T is temperature in kelvin and ais the charge transfer coefficient. This gives a 5 0.474(2) for AQDS and 0.43(1) for
DHAQDS. The x-intercept gives the log of the exchange current i0, which is equal
to FACOk0, and gives k0 5 7.2(5) 3 1023 cm s21 for AQDS and 1.56(5) 3 1022 cm s21
for DHAQDS.
Stability studies. AQDS (50 mg) was dissolved in 0.4 ml of D2O, and treated with
100 ml of Br2. The 1H and 13C NMR spectra (Extended Data Figs 5a, b and 6a, b)
were unchanged from the starting material after standing for 20 h at 25 uC. AQDS
(50 mg) was then treated with 1 ml of 2 M HBr and 100 ml of Br2. The reaction washeated to 100 uC for 48 h and evaporated to dryness at that temperature. The
resulting solid was fully dissolved in D2O giving unchanged 1H and 13C NMR(Extended Data Figs 5c and 6c); however, the 1H NMR reference was shifted due toresidual acid. These results imply that bromine crossover will not lead to irreversible destruction of AQDS.
Sulphonation of anthraquinone and electrochemical study. 9,10-Anthraquinone
was treated with H2SO4 (20% SO3) at 170 uC for 2 h according to a literature procedure8. The resulting red solution, containing roughly 37% AQDS, 60% 9,10anthraquinone-2,6-disulphonic acid and 3% 9,10-anthraquinone-2-sulphonicacid, was diluted and filtered. A portion of this solution was further diluted with
1 M H2SO4 to ,1 mM total anthraquinone concentration. The cyclic voltammogram (Extended Data Fig. 8) is similar to that of pure AQDS, though the anodic/
cathodic peak current density is broadened due to the presence of the multiple
sulphonic acid isomers.
Theory and methods. We used a fast and robust theoretical approach to determine the E0 of quinone/hydroquinone couples in aqueous solutions. We employed
an empirical linear correlation of DHf, the heat of formation of hydroquinone at
0 K from the quinone and hydrogen gas, to the measured E0 values22. Following thetreatment of ref. 22, the linear correlation is described as DG 5 2nFE0, where DG
is the difference in total free energy between quinone and hydroquinone, n is thenumber of electrons involved in the reaction and F is the Faraday constant. The
entropy contributions to the total free energies of reaction have been neglected
because the entropies of reduction of quinones are found to be very similar22,23, andthe E0 of the oxidation–reduction system is linearly expressed as (2nF)21DHf 1 b,
where b is a constant. It was also assumed that the reduction of quinones takes
place in a single-step reaction involving a two-electron two-proton process9,24. The
total free energies of molecules were obtained from first-principles quantum
chemical calculations within density functional theory (DFT) at the level of thegeneralized gradient approximation (GGA) using the PBE functional25. The projectoraugmentedwave(PAW)techniqueandaplane-wavebasisset26,27 asimplemented©2014 Macmillan Publishers Limited. All rights reserved

RESEARCH LETTER
in VASP28,29 were employed. The kinetic energy cut-off for the plane-wave basis
was set at 500 eV, which was sufficient to converge the total energies on a scale of
1 meV per atom. To obtain the ground-state structures of molecules in the gas
phase, we considered multiple initial configurations for each molecule and optimized them in a cubic box of 25 Å using C-point sampling. The geometries were
optimized without any symmetry constraints using the conjugate gradient (CG)algorithm, and the convergence was assumed to be complete when the totalremaining forces on the atoms were less than 0.01 eV Å21.
The search for conformational preference through theoretical calculations for eachhydroxylated quinone is crucial because of the significant effects of intramolecularhydrogen bonds on the total free energies of the molecules30. Three-dimensional
conformer structures for each quinone/hydroquinone molecule were generated using
the ChemAxon suite (Marvin 6.1.0 by ChemAxon, http://www.chemaxon.com)
with up to 25 conformers generated per molecule using the Dreiding force field31.
The conformers generated were used as input structures for the DFT geometry
optimization employed for determining DHf, which in turn is used to estimate E0
and G0solv .
To calculate the E0 of the hydroxy-substituted AQDS molecules (Fig. 3c), thecorrelation between DHf and E0 was calibrated from experimental data on six wellcharacterized quinones32. Specifically, we used the experimental values of the aqueous
E0 and computed DHf of 1,2-benzoquinone, 1,4-benzoquione, 1,2-naphthoquinone,
1,4-naphthoquinone, 9,10-anthraquinone and 9,10-phenanthrene33. The training
set ensures that the calibration plot addresses most classes and aspects of quinones,
including two quinones each from 1-ring (benzoquinone), 2-ring (naphthoquinone)and 3-ring (anthraquinone and phenanthrene) structures. In addition, the experimental values of E0 of the training set spanned 0.09 V (9,10-anthraquinone) to0.83 V (1,2-benzoquinone), providing a wide range for E0 (Extended Data Fig. 7).
The linear calibration plot for E0 yields an R2 5 0.97 between the calculated DHf
and E0 (Extended Data Fig. 7).
The G0solv values of the quinones in water were calculated using the Jaguar 8.0
program in the Schrödinger suite 2012 (Jaguar, version 8.0, Schrödinger). The
standard Poisson–Boltzmann solver was employed34,35. In this model, a layer ofcharges on the molecular surface represents the solvent. G0solv was calculated as thedifference between the total energy of the solvated structure and the total energy ofthe molecule in vacuum. A more negative value for G0solv corresponds to a quinonethat is likely to have a higher aqueous solubility. An absolute prediction of thesolubility is not readily available, as the accurate prediction of the most stable formsof molecular crystal structures with DFT remains an open problem36.
Cost calculations. These were done using the following formula: C 5 (3.6 3 103) 3(PM)/(nFE), where C is the cost in US dollars of the compound per kilowatt hour, P
is the cost in US dollars per kilogram, M is the molecular mass of the compound,
F is Faraday’s constant (96,485 C mol21), n is the number of moles of electronstransferred per mole of storage compound (two for the QBFB) and E is the opencircuit voltage (V) of the storage device. In calculating the price for the anthraquinonebromine battery, a price of $4.74 per kilogram (eBioChem) was used for anthraquinone(note that, to get the sulphonated form actually used here, anthraquinone must bereacted with oleum (H2SO4/SO3), which adds a negligible cost at scale; this cost is
not included here). The price of bromine was $1.76 per kilogram, based on estimates from the US Geological Survey17. The cell voltage used to calculate costs here
was 0.858 V. For vanadium, costs were calculated from USGS prices from 201118 of
vanadium pentoxide at $14.37 per kilogram, and the cell voltage used was 1.2 V.
Balance-of-system costs have not been estimated because the technology is too
immature.
19. Skyllas-Kazacos, M., Milne, N. A. & Kazacos, G. C. Membrane properties andbehavior in the Generation 2 vanadium bromide redox flow batteries. Materials
Forum 32, 72–77 (2008).
20. Yeo, R. S. & McBreen, J. Transport properties of Nafion membranes inelectrochemically regenerative hydrogen/halogen cells. J. Electrochem. Soc. 126,
1682–1687 (1979).
21. Haase, J., Baudys, J., Obruba, K. & Panek, J. Dihydroxyanthraquinone disulfonicacids. Czechoslovakian patent 148,977 (1973).
22. Dewar, M. J. S. & Trinajstic, N. Ground states of conjugated molecules-XIV: redox
potentials of quinones. Tetrahedron 25, 4529–4534 (1969).
23. Pullman, B. & Pullman, A. Quantum Biochemistry 475 (Interscience, 1963).
24. Guin, P. S., Das, S. & Mandal, P. C. Electrochemical reduction of quinones indifferent media: a review. Int. J. Electrochem. 816202 (2011).
25. Perdew, J. P., Burke, K. & Ernzerhof, M. Generalized gradient approximation madesimple. Phys. Rev. Lett. 77, 3865–3868 (1996).
26. Blöchl, P. E. Projector augmented-wave method. Phys. Rev. B 50, 17953–17979(1994).
27. Kresse, G. & Joubert, D. From ultrasoft pseudopotentials to the projector
augmented-wave method. Phys. Rev. B 59, 1758–1775 (1999).
28. Kresse, G. & Hafner, J. Ab initio molecular dynamics for liquid metals. Phys. Rev. B
47, 558–561 (1993).
29. Kresse, G. & Furthmuller, J. Efficient iterative schemes for ab initio total-energy
calculations using a plane-wave basis set. Phys. Rev. B 54, 11169–11186 (1996).
30. Qu, R., Liu, H., Feng, M., Yang, X. & Wang, Z. Investigation on intramolecularhydrogen bond and some thermodynamic properties of polyhydroxylated
anthraquinones. J. Chem. Eng. Data 57, 2442–2455 (2012).
31. Mayo, S. L., Olafson, B. D. & Goddard, W. A. III Dreiding: a generic force field formolecular simulations. J. Phys. Chem. 94, 8897–8909 (1990).
32. Wang, J. & Hou, T. Recent advances on aqueous solubility prediction. Comb. Chem.
High Throughput Screen. 14, 328–338 (2011).
33. Johnsson Wass, J. R. T., Ahlberg, E., Panas, I. & Schiffrin, D. J. Quantum
chemical modeling of the reduction of quinones. J. Phys. Chem. A 110, 2005–2020(2006).
34. Tannor, D. J. et al. Accurate first principles calculation of molecular charge
distributions and solvation energies from ab initio quantum mechanics andcontinuum dielectric theory. J. Am. Chem. Soc. 116, 11875–11882 (1994).
35. Marten, B. et al. New model for calculation of solvation free energies: correction ofself-consistent reaction field continuum dielectric theory for short-range
hydrogen-bonding effects. J. Phys. Chem. 100, 11775–11788 (1996).
36. Hongo, K., Watson, M. A., Sánchez-Carrera, R. S., Iitaka, T. & Aspuru-Guzik, A. Failureof conventional density functionals for the prediction of molecular crystal
polymorphism: a quantum Monte Carlo study. J. Phys. Chem. Lett. 1, 1789–1794(2010).

©2014 Macmillan Publishers Limited. All rights reserved

LETTER RESEARCH

0

Current Density (μA cm−2)-200

200 r.p.m.

-400

-600

-800

-1,0003600 r.p.m.
-1,200-100

0100
Potential (mV vs SHE)

200

300

Extended Data Figure 1 | Plot of potential versus current density at
different rotation rates of the RDE. The solution is 1 mM DHAQDS (1 mM
in 1 M H2SO4), using a rotating disk electrode (RDE) of glassy carbon. Rotation
rates are 200, 300, 400, 500, 700, 900, 1,200, 1,600, 2,000, 2,500 and 3,600 r.p.m.

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RESEARCH LETTER
-100

-80

-70

-70

-60

Limiting Current (μA)

Limiting Current (μA)

-80

a

-90

-60
-50
-40
-30

b

-50
-40
-30
-20

-20

-10

-10
0
0

2

4

6

8

10
ω

1/2

12

(Rad/s)

14

16

18

0

20

0

2

4

6

8

1/2

10
12
14
ω1/2 (Rad/s)1/2

16

18

20

-70

c
-60
-50

i−1 (mA−1)-40
η / mV
13
18
23
28
33
38
236

-30
-20
-10

0
0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.18 0.20 0.22 0.24
ω−1/2 (Rad/s)−1/2

Extended Data Figure 2 | Levich and Koutecký–Levich plots obtainedusing the RDE. a, Levich plot (limiting current versus square root ofrotation rate v) of 1 mM AQDS in 1 M H2SO4 (the fitted line has a slope of
4.53(2) mA s1/2 rad21/2, giving D 5 3.8(1) 3 1026 cm2 s21). Data are an averageof three runs; error bars indicate the standard deviation. b, As a but for

DHAQDS in 1 M H2SO4 (slope of 3.94(6) mA s1/2 rad21/2
gives D 5 3.19(7) 3 1026 cm2 s21). c, Koutecky–Levich plot (i21 versus v21/2)
of 1 mM DHAQDS in 1 M H2SO4. The current response, i, is shown forseven different AQDS reduction overpotentials g.

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LETTER RESEARCH
45

45

b

40

40

35

35

30

30

Overpotential (mV)

Overpotential (mV)

a

25
20
15

25
20
15

10

10

5

5

0
-4.4

-4.3

-4.2

-4.1

-4.0

-3.9

-3.8

-3.7

-3.6

0
-4.0

-3.9

log(iK)

Extended Data Figure 3 | Fit of Butler–Volmer equation. Constructedusing the current response in the absence of mass transport at low AQDS
reduction overpotentials; iK is the current extrapolated from the zero-intercept
of Fig. 3b and Extended Data Fig. 2c (infinite rotation rate). Data are an average

-3.8

-3.7
-3.6
log(iK)

-3.5

-3.4

-3.3

of three runs; error bars indicate the standard deviation. a, AQDS: best-fit
line has the equation y 5 62(x 1 4.32). This yields a 5 0.474(2) andk0 5 7.2(5) 3 1023 cm s21. b, DHAQDS: best-fit line is the function
y 5 68(x 1 3.95). This yields a 5 0.43(1) and k0 5 1.56(5) 3 1022 cm s21.

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RESEARCH LETTER300

E0 (mV vs SHE)

200

100

0

-100

-200

-300
0

2

4

6

8

10

12

14

pH

Extended Data Figure 4 | Pourbaix diagram (E0 vs pH) of AQDS. Data are
fitted to three solid lines indicating slopes of 259 mV pH21, 230 mV pH21
and 0 mV pH21, corresponding to two-, one- and zero-proton processes,
respectively. Dashed lines linearly extrapolate the one- and zero-proton
processes to give E0 values of 18 mV (2e2/1H1) and 2296 mV (2e2/0H1).

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LETTER RESEARCH

Extended Data Figure 5 | 1H NMR (500 MHz, D2O) spectra. a, Spectrum of
AQDS: chemical shift d 5 7.99 p.p.m. versus tetramethylsilane (TMS) (doublet(d), coupling constant J 5 2 Hz, 1,8 C–H), 7.79 p.p.m. (doublet of doublets,
J 5 2 and 8 Hz, 4,5 C–H), 7.50 p.p.m. (d, J 5 8 Hz, 3,6 C–H). b, The same

sample, 20 h after addition of Br2. c, 1H NMR of AQDS treated with 2 M HBr
and Br2 and heated to 100 uC for 48 h. The peaks are shifted due to presence oftrace HBr, which shifted the residual solvent peak due to increased acidity.
Coupling constants for each peak are identical to a.

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RESEARCH LETTER

Extended Data Figure 6 | 13C NMR (500 MHz, D2O) spectra. a, AQDS,
d 5 181.50 p.p.m. versus TMS (C 9), 181.30 p.p.m. (C 10), 148.51 p.p.m.
(C 2,7), 133.16 p.p.m. (C 11), 132.40 p.p.m. (C 12), 130.86 p.p.m. (C 3,6),

128.59 p.p.m. (C 4,5), 124.72 p.p.m. (C 1,8). b, The same sample, 24 h after
addition of Br2. c, 13C NMR of AQDS treated with 2 M HBr and Br2 and heated
to 100 uC for 48 h.

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LETTER RESEARCH
0.9
1,2-BQ

R2 = 0.97

0.8
1,4-BQE0 from Literature (V)0.7
0.6
1,2-NQ

1,4-NQ

0.5
9,10-PQ

0.4
0.3
0.2
0.1
0.0
-2.5

9,10-AQ-2.0-1.5
ΔHf (eV)

-1.0

-0.5

Extended Data Figure 7 | Calibration model for DHf and experimental E0.
This shows a linear relationship (red dashed line; R2 5 0.97) between calculated
DHf (this work) and experimental E0 (from the literature) of six quinones
in aqueous solutions: BQ, benzoquinone; NQ, naphthoquinone; AQ,
anthraquinone; and PQ, phenanthraquinone.

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RESEARCH LETTER

Current Density (μA cm−2)

200

100

0

-100

-200

-300
-100

0

100
200
300
Potential (mV vs. SHE)

400

500

Extended Data Figure 8 | AQDS cyclic voltammograms. Black curve,
obtained for a 1 mM solution of AQDS in 1 M H2SO4 on a stationary glassy
carbon working electrode. Red curve, obtained for a crude anthraquinone
sulphonation solution containing a mixture of AQDS, 9,10-anthraquinone-2,6disulphonic acid and 9,10-anthraquinone-2-sulphonic acid diluted to 1 mM
total anthraquinone in 1 M H2SO4.

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LETTER RESEARCH
1.5

Cell Potential (V)

1.2

0.9

0.6

0.3

0.0
0.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

4.0

4.5

5.0

5.5

Time (hrs)

Extended Data Figure 9 | Flow-battery cycling behaviour with HBrelectrolyte on both sides. Data collected by cycling the current at 0.2 A cm22 at
40 uC using a 2 M HBr 1 0.5 M Br2 solution on the positive side and a 2 M
HBr 1 0.1 M AQDS solution on the negative side; cell potential versus time
performance is comparable to data in Fig. 2.

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RESEARCH LETTER
Extended Data Table 1 | AQDS screened by theoretical calculations

ID
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36

-OH
substituted
NonMono-

Di-

Tri-

Tetra-

PentaHexa-

R1

R3

R4

R5

R6

R8

E (V)

0

G solv (kJ mol )

H
OH
H
H
OH
OH
OH
OH
OH
H
H
H
H
OH
OH
OH
OH
OH
OH
OH
OH
H
H
OH
OH
OH
OH
OH
OH
OH
OH
H
OH
OH
OH
OH

H
H
OH
H
OH
H
H
H
H
OH
OH
OH
H
OH
OH
OH
OH
H
H
H
H
OH
OH
OH
OH
OH
OH
OH
OH
H
H
OH
OH
OH
OH
OH

H
H
H
OH
H
OH
H
H
H
OH
H
H
OH
OH
H
H
H
OH
OH
OH
H
OH
OH
OH
OH
OH
H
H
H
OH
OH
OH
OH
OH
OH
OH

H
H
H
H
H
H
OH
H
H
H
OH
H
OH
H
OH
H
H
OH
H
H
OH
OH
H
OH
H
H
OH
OH
H
OH
OH
OH
OH
OH
H
OH

H
H
H
H
H
H
H
OH
H
H
H
OH
H
H
H
OH
H
H
OH
H
OH
H
OH
H
OH
H
OH
H
OH
OH
H
OH
OH
H
OH
OH

H
H
H
H
H
H
H
H
OH
H
H
H
H
H
H
H
OH
H
H
OH
H
H
H
H
H
OH
H
OH
OH
H
OH
H
H
OH
OH
OH

0.222
0.185
0.325
0.108
0.176
0.027
0.122
0.143
0.101
0.153
0.179
0.202
0.000
-0.070
0.083
0.187
0.310
-0.1020.089-0.085-0.048-0.1070.106-0.0980.012-0.222-0.0190.046
0.080-0.259-0.199-0.083-0.252-0.292-0.030-0.077-81.5-81.5-111.7-88.2-110.3
-85.6
-96.7
-85.7
-83.2
-105.4-119.1-112.0-95.6-101.7-116.2-114.3-120.9-91.4-114.0-87.1-102.8-107.8-136.8-109.0-108.4-102.3-132.3-114.6-111.1-99.0-91.9-120.6-117.1-108.3-111.6-121.0The effect of –OH substitution on reduction potential and solvation energy.

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