ll

ArticleExtremely Stable Anthraquinone Negolytes
Synthesized from Common Precursors
Min Wu, Yan Jing, Andrew A.
Wong, ..., Zhijiang Tang, Roy G.
Gordon, Michael J. Aziz
gordon@chemistry.harvard.edu (R.G.G.)
maziz@harvard.edu (M.J.A.)

HIGHLIGHTS
Potentially low-cost andextremely stable anthraquinone
negolytes
No SN2 or SNAr decomposition
reaction in carbon-linked
anthraquinones
Record low capacity fade rate of<1% per year
Molecules exhibit longer lifetime
at high pH than that at mild pH

Two extremely stable anthraquinone negolytes were synthesized from inexpensive
precursors that potentially decrease the mass production cost. The carbon-linked
anthraquinones eliminate SN2 or SNAr side reactions. Pairing with a Fe(CN)63 /4
posolyte, they exhibited an open-circuit voltage of 1.0 V. By operating at pH 14, a
record low capacity fade rate of <1% per year was demonstrated.

Wu et al., Chem 6, 1432–1442
June 11, 2020 ª 2020 Elsevier Inc.
https://doi.org/10.1016/j.chempr.2020.03.021

ll
ArticleExtremely Stable Anthraquinone Negolytes
Synthesized from Common Precursors
Min Wu,1 Yan Jing,2 Andrew A. Wong,1 Eric M. Fell,1 Shijian Jin,1 Zhijiang Tang,1 Roy G. Gordon,1,2,*
and Michael J. Aziz1,3,*

SUMMARY

The Bigger Picture

Synthetic cost and long-term stability remain two of the most challenging barriers for the utilization of redox-active organic molecules
in redox flow batteries for grid-scale energy storage. Starting frompotentially inexpensive 9,10-dihydroanthracene, we developed a
new synthetic approach for two extremely stable anthraquinone negolytes, i.e., 3,30 -(9,10-anthraquinone-diyl)bis(3-methylbutanoicacid) (DPivOHAQ) and 4,40 -(9,10-anthraquinone-diyl)dibutanoic
acid (DBAQ). Pairing with a ferrocyanide posolyte at pH 12, DPivOHAQ and DBAQ can transfer up to 1.4 and 2 M electrons with capacity fade rates of 0.014% per day and 0.0084% per day, respectively,
and exhibit 1.0 V of open-circuit voltage. By adjusting the supporting electrolytes to pH 14, DPivOHAQ exhibited a record low capacity fade rate of <1% per year. We attribute the capacity loss of theseflow batteries primarily to the formation of anthrone, which can besuppressed by increasing the pH of the electrolyte and reversed byexposure to air.

The cost of solar and windelectricity is decreasing so rapidlythat a grid-scale energy storage
technology will become essential.
Aqueous organic redox flow
batteries are a potentially safe,
inexpensive substitute for lithiumion batteries and vanadium flow
batteries for large-scale energy
storage. Here, we report a new
synthetic strategy for two
extremely stable anthraquinone
negolyte (negative electrolyte)
molecules starting from
inexpensive precursors that
potentially decrease the costwhen scaled up. Additionally, wedemonstrate that ananthraquinone negolyte is morestable running at pH 14 than at pH
12, and is expected to be morestable in alkaline solution than inacidic or neutral conditions.
Paired with a Fe(CN)63/4
positive electrolyte, theanthraquinone cell exhibited arecord low capacity fade rate of<1% per year. The new synthetic
strategy for these highly stable
anthraquinone negolytes might
facilitate the commercialization ofanthraquinone-based flow
batteries.

INTRODUCTION
The cost of solar and wind electricity has dropped so precipitously that the main barrier to widespread implementation is their intrinsic intermittency.1–3 A safe, low-cost,large-scale electrical energy storage system could enable grid-scale adoption of renewables. Among the numerous proposed technologies, redox flow batteries (RFBs)
have been recognized as a potentially viable strategy to address the intermittency of
renewable energy.2–4 Compared with conventional stationary rechargeable batteries (e.g., lithium-ion batteries and lead-acid batteries), RFBs use redox-active materials dissolved in liquid supporting electrolytes that are stored in external tanks andseparated from the power generation stack. This separation allows for the decoupling of energy capacity from output power capacity, thereby providing the possibility of low-cost long-duration discharge.2–4
Aqueous RFBs, featuring non-flammable electrolytes, are particularly suitable for storing
massive amounts of electricity. Aqueous vanadium RFBs are the most widely studied andadopted systems, but are hindered by the high cost of vanadium.3–5 In contrast, redoxactive organic molecules comprising earth abundant elements such as carbon (C),hydrogen (H), oxygen (O), and nitrogen (N) have the potential to be inexpensive alternatives to vanadium.6–13 Additionally, the structural diversity and tunability of organics
enable chemists to design organics with essential properties such as high aqueous solubility, high chemical stability, fast kinetics, and appropriate redox potential.6,8–12,14,15
Recently, water-soluble anthraquinones 4,40 -([9,10-anthraquinone-2,6-diyl]dioxy)dibutyrate (2,6-DBEAQ) and (([9,10-dioxo-9,10-dihydroanthracene-2,6-diyl]bis[oxy])1432 Chem 6, 1432–1442, June 11, 2020 ª 2020 Elsevier Inc.

ll

Articlebis[propane-3,1-diyl])bis(phosphonic acid) (2,6-DPPEAQ) in mildly alkaline solutions
have demonstrated extremely low temporal fade rates in flow batteries paired with
K4Fe(CN)6.16–18 These quinones are chemically synthesized from 2,6-dihydroxyanthraquinone (2,6-DHAQ) by industry-compatible methods. However, 2,6-DHAQ
and 2,7-DHAQ are always co-produced and are costly to separate. Furthermore,
our previous research showed that the molecular lifetimes of anthraquinone-based
electrolytes can differ by two orders of magnitude depending on the positions of
their functional groups (e.g., 1,8- and 2,6-anthraquinones).16,19 Therefore, it isimportant to quantify the stabilities of organic molecules with a mixture of isomers.
Additionally, 2,6-DHAQ and 2,7-DHAQ are synthesized from 9,10-anthraquinone2,6-disulfonic acid and 9,10-anthraquinone-2,7-disulfonic acid, respectively, instrong alkaline solution for 35 h at a high temperature (180  C) with a moderate yield,
which is energy-intensive and costly.20 Thus, designing low-cost and chemically stable anthraquinones is of vital importance for the commercialization of aqueousorganic RFBs.
Because of the inherent chemical stability of the parent structure, the addition ofside chains to an anthraquinone core is usually accomplished by a stepwise procedure via anthraquinone derivatives (e.g., hydroxylated anthraquinone or chlorinatedanthraquinone).16,17,19,21 Here, we report a new synthetic route for water-soluble anthraquinones starting from a potentially inexpensive anthracene derivative, 9,10-dihydroanthracene, which can be readily produced from anthracene with a yield ofalmost 100%.22,23 Anthracene, a component of coal tar, is one of the major resources
for large-scale anthraquinone production.24 The first step is a Friedel-Crafts alkylation or acylation to render dihydroanthracene and anthracene water-soluble; thelast step is an oxidation to produce the corresponding redox-active anthraquinones,i.e., 3,30 -(9,10-anthraquinone-diyl)bis(3-methylbutanoic acid) (DPivOHAQ) and 4,40 (9,10-anthraquinone-diyl)dibutanoic acid (DBAQ).
Both molecules exhibit high water solubility and chemical stability at pH 12. The
DBAQ negolyte (negative electrolyte) has a water solubility of 1.0 M, corresponding
to a volumetric capacity of 53.6 Ah L; when paired with potassium ferrocyanide, a fullcell exhibited a capacity fade rate of 0.0084% per day or 3.1% per year. DPivOHAQ
has a solubility of 0.74 M; when paired with potassium ferrocyanide, a full cell exhibited a capacity fade rate of 0.014% per day or 5.1% per year. Additionally, wedemonstrated that the DPivOHAQ negolyte is even more stable in strong base, exhibiting a capacity fade rate of 0.0018% per day or 0.66% per year at pH 14. Furthermore, we demonstrated that the capacity fade is due to formation of anthrone, which
can convert back to anthraquinone through air exposure and can also be suppressedat high pH. Thus, these findings suggest that, through a combination of increased
pH and periodic air exposure, both DPivOHAQ and DBAQ offer the possibility of
decadal lifetimes in aqueous RFBs. In the following section we report methods
and results, first for DPivOHAQ and then for DBAQ.

RESULTS AND DISCUSSION
Figure 1 illustrates the synthetic routes, chemical structure, Pourbaix diagram, andcyclic voltammogram (CV) of DPivOHAQ. Synthesis was achieved by first functionalizing 9,10-dihydroanthracene with the water-soluble group –C(CH3)2CH2COOH, followed by an oxidation step with CrO3. The oxidation step is well known in industry
for anthraquinone synthesis, and it can also be accomplished by other methods,
such as nitric acid and air, as confirmed by the 1H NMR of DPivOHAQ in Figure S3,that could further decrease the costs when scaled up.25,26 However, CrO3 oxidation

1John A. Paulson School of Engineering andApplied Sciences, Harvard University,
Cambridge, MA 02138, USA
2Department of Chemistry and Chemical Biology,

Harvard University, Cambridge, MA 02138, USA
3Lead Contact*Correspondence:
gordon@chemistry.harvard.edu (R.G.G.),
maziz@harvard.edu (M.J.A.)
https://doi.org/10.1016/j.chempr.2020.03.021

Chem 6, 1432–1442, June 11, 2020 1433

ll

Figure 1. DPivOHAQ Synthesis and Electrochemistry(A) Synthetic route for DPivOHAQ.
(B) Pourbaix diagram.
(C) Cyclic voltammograms of 5 mM DPivOHAQ and 10 mM potassium ferrocyanide at pH 12 with a scan rate of 100 mV s –1 .

is a more controllable method on a research scale to get a purer product, therefore,
in this work, we use CrO3 as the oxidant unless specifically mentioned otherwise.
Compared with 2,6-DBEAQ and 2,6-DPPEAQ synthesis, the reported synthetic
approach could be potentially more cost-effective. For example, assuming anthracene is the starting material, both DBEAQ and DPPEAQ require five synthetic steps
in total as illustrated in Scheme S1 with an overall yield around 52% and 59%, respectively; whereas, DPivOHAQ synthesis requires only three steps with an overall yieldof 81%. Moreover, the cost of side chains for DPivOHAQ is slightly lower than that of
DBEAQ and substantially lower than that of DPPEAQ at sub-kg scales, as shown in
Table S1. Therefore, we expect that DPivOHAQ is most likely less expensive to produce than DBEAQ and DPPEAQ at industrial scales.
In addition to potentially lower synthetic costs, DPivOHAQ is functionalized with carbon-linked functional groups, which are chemically more robust than the oxygenlinked side chains in DBEAQ and DPPEAQ, minimizing the opportunity for SN2
and SNAr side reactions to occur.16,17 The thermochemical stability of both oxidized
and reduced forms of DPivOHAQ were evaluated at high temperature (65  C) and instrongly alkaline conditions (pH 14) for eight days. No apparent decomposition was
detected from the 1H NMR as shown in Figure S8, indicating that DPivOHAQ is quitechemically stable. A temperature of 65  C is well above anticipated operating temperatures of RFBs; thus there should be much less decomposition in realapplications.
Synthesis of DPivOHAQ results in a mixture of 2,6- and 2,7- isomers that does not
require further separation prior to use in a battery. The Pourbaix diagram of DPivOHAQ, shown in Figure 1B, suggests that the molecule undergoes a two-proton/

1434 Chem 6, 1432–1442, June 11, 2020

ArticleArticlell

two-electron process below pH 9, a one-proton/two-electron process between pH 9
and 11, and a zero-proton process with a pH-independent potential of approximately 0.48 V versus standard hydrogen electrode (SHE) at pH > 11. Pairing a DPivOHAQ negolyte with potassium ferrocyanide at pH 12 should yield an equilibrium
cell potential of approximately 0.98 V (Figure 1C). Electrochemical kinetics of DPivOHAQ reduction were determined with the rotating disk electrode (RDE) method asshown in Figure S9. The charge transfer coefficient is 0.49, the diffusion coefficientis 2.4 3 106 cm2 s–1, and the kinetic rate constant is 2.5 3 103 cm s–1; the latteris much higher than is typical of inorganic redox-active materials at an uncatalyzed
carbon electrode.27 It should be noted that such a high kinetic rate constant could
be a lower bound to the actual value, because a RDE system exhibits mixed kinetic
and transport limits in the observed mass-transfer limited current density region. The
solubilities of the oxidized forms of anthraquinones in alkaline solutions are typicallylower than those of the reduced forms. The solubility of DPivOHAQ at pH 12 wasdetermined by UV-vis spectrophotometry (Figure S10) to be 0.74 M, corresponding
to a volumetric capacity of 39.7 Ah L–1.
Polarization experiments of a 0.5 M DPivOHAQ-ferrocyanide full cell at pH 12 were
performed at various states of charge. The electrolytes comprised 5 mL of 0.5 M DPivOHAQ (negolyte) at pH 12 (10 mM KOH) and 80 mL of 0.3 M potassium ferrocyanide and 0.1 M potassium ferricyanide (posolyte) at pH 12. The cell was constructedfrom graphite flow plates and AvCarb carbon cloth electrodes, separated by a Fumasep E-620 (K) membrane because of its low permeability to ferricyanide and highionic conductivity.16,17 A peak galvanic power density of 0.34 W cm2 was achievedat 100% state of charge (SOC) (Figure 2A). Careful attention to engineering designand construction should raise the power density even farther.28 The open-circuit
voltage (OCV) increases from 0.95 to 1.08 V as the SOC increases from 10% to100%, and the OCV at 50% SOC of 0.99 V (Figure 2B) is consistent with the voltage
expected from CV. The alternating current area-specific resistance (ASR) of the cell
was determined via high-frequency potentiostatic electrochemical impedance spectroscopy (EIS), and the value was below 0.6 U cm2 across all SOCs (Figure 2B). This isa relatively low alternating current ASR value for RFBs with alkaline electrolytes.16,17
The polarization ASR was determined using the linear region within the voltage
range 0.9–1.1 V (Figures 2A and 2B). The ASR of the membrane (0.54 U cm2 at
50% SOC, determined by high-frequency EIS in the full cell) accounted for around
67% of the ASR of the entire cell (0.81 U cm2 at 50% SOC, DC polarization). The capacity utilization for the negolyte (capacity-limiting side) is approximately 95% at 50
mA cm2 with a high round-trip energy efficiency of 91.5% (Figures 2C and 2D). At areasonable practical operation target of 80% round-trip energy efficiency, the low
value of the ASR permits galvanostatic operation at around 140 mA cm2 with anelectrolytic power density of 0.16 W cm2, a galvanic power density of 0.13 W
cm2, and 92% capacity utilization.
The same 0.5 M DPivOHAQ-ferrocyanide full cell was used for long-term stability
evaluation (Figure 3). The cell was cycled at a constant current density of G0.1 A
cm2, and each galvanostatic half cycle was followed by a potential hold at thevoltage limit (1.3 V for charge, 0.6 V for discharge) until the current density fell below
2 mA cm2 to mitigate the effect of temporal variations in accessible capacity during
full cell cycling caused by drifts in cell resistance.29 The charge-discharge profiles
near the voltage limits (Figures 3B and 3C) are quite steep and are followed by small
subsequent horizontal segments during the potential holds. The horizontal segments end at 95.5% of theoretical capacity of negolyte, but the steepness followed
by a small subsequent horizontal segment of the charge-discharge profiles suggests

Chem 6, 1432–1442, June 11, 2020 1435

llArticleFigure 2. Polarization Measurements of 0.5 M DPivOHAQ/Ferrocyanide Full Cell at pH 12(A) Cell voltage versus current density at room temperature (~20  C) at various SOC.
(B) OCV, high-frequency ASR and polarization ASR versus SOC.
(C) Galvanostatic charge-discharge voltage profiles at various current densities. The vertical dashed line indicates the theoretical capacity.
(D) Capacity utilization, Coulombic efficiency, and round-trip energy efficiency versus current density.

that the electrolyte had around 4.5% inactive material and that the active material isundergoing deep cycling to essentially the full SOC limits. The cell was cycled for690 cycles at 100 mA cm2, which required 15.6 days to complete. The capacity
retention over the 690 cycles was 99.78% with an average Coulombic efficiency
greater than 99.9%, reflecting a capacity fade rate of 0.00031% per cycle or
0.014% per day (Figure 3A), i.e., 5.1% per year. This temporal fade rate is among
the lowest exhibited by full cells in which organic molecules composed the capacity-limiting side: 2,6-DBEAQ fades at 0.04% per day; 2,6-DPPEAQ fades at
0.014% per day.18
After 15.6 days of cycling at pH 12, DPivOHAQ negolyte, in the discharged state,
was exposed to air for 2 h and the pH was adjusted to 14 by dissolving KOH pellets
into the negolyte and posolyte without changing cell materials or setup. As a result,
81% of the lost capacity was recovered, as shown in Figure 3A. Over the additional
16 days at pH 14, the cell exhibited a capacity fade rate of 0.0018% per day, which is
6 times lower than that at pH 12. The charge-discharge voltage profiles (Figures 3B
and 3C) are almost invariant, indicating no apparent change in ohmic resistance and1436 Chem 6, 1432–1442, June 11, 2020

ArticlellFigure 3. Cycling Performance of DPivOHAQ at pH 12 and 14(A) Coulombic efficiency and discharge capacity versus time and cycle number for a negolytelimited DPivOHAQ/K 4 Fe(CN) 6 full cell. Each 100 mA cm 2 half cycle was followed by apotentiostatic hold until the magnitude of the current density fell below 2 mA cm2 . The negolytecomprised 5 mL of 0.5 M DPivOHAQ at pH 12, and the posolyte comprised 80 mL of 0.3 M
potassium ferrocyanide and 0.1 M potassium ferricyanide at pH 12. After approximately 16 days ofcycling, the negolyte was exposed to air and the pH of both negolyte and posolyte were adjusted to
14 before cycling for an additional 16 days. Note that the left y axis represents only 0.54% of thecapacity of the DPivOHAQ negolyte.
(B) Charge-discharge voltage profile of DPivOHAQ from selected cycles at pH 12 in Figure 3A.
(C) Charge-discharge voltage profile of DPivOHAQ from selected cycles at pH 14 in Figure 3A.

good chemical compatibility with cell membrane and other cell components.30
Based on the decomposition study of 2,6-DHAQ,31 we attribute capacity fade toanthrone formation (Scheme 1). Therefore, increasing the hydroxide concentration
should suppress the formation of anthrone. This expectation is consistent with thelower capacity fade rate observed at pH 14 than at pH 12. In general, the disproportionation reaction will generate OH (or consume H+) at pH above the first pKa of theanthrahydroquinone. Therefore, anthrone formation will be disfavored under alkaline conditions relative to acid conditions and will be progressively disfavored as
the pH increases (Figure S11). We interpret the sudden increase in capacity at cycle
691 as the consequence of anthrone being converted back to anthraquinone by boththe pH effect and the effect of exposure to atmospheric O2.
To confirm the major side reaction is the disproportionation of reduced anthraquinone, afully reduced (100% SOC) sample of DPivOHAQ at pH 12 was prepared and stored inan fluorinated ethylene propylene (FEP) vial in a glove box for 238 days, allowing thedisproportionation to reach equilibrium. Indeed, some appreciable side peaks appeared in the 1H NMR spectrum of the reduced DPivOHAQ, whereas upon re-oxidationin air, the 1H NMR spectrum contained no observable decomposition peaks (Figure S12),
indicating that the decomposition compounds can either be re-oxidized back to

Chem 6, 1432–1442, June 11, 2020 1437

llArticleO+ 2 H2O

2

O

O

O

+ 3 OH-+O

Scheme 1. The Reversible Disproportionation Reaction of Anthraquinone Negolyte under
Alkaline Conditions
The pK a of anthrone is reported32 to be 10; therefore, anthrone is depicted in the deprotonatedform at pH values between 12 and 14. 31

DPivOHAQ or are converted to other products with no observable signals above thedetection limit of the NMR instrument. Therefore, high-performance liquid chromatography-mass spectrometry was performed to analyze both the reduced and the reoxidized samples (Figure S13). The anthrone species was detected in the reduced sample but not in the re-oxidized sample, in agreement with the 1H NMR result shown in Figure S12. In the re-oxidized sample after 238 days, approximately 1.24% of the signal corresponded to an anthrone dimer, suggesting a fade rate of 1.90% per year for reduced
DPivOHAQ at pH 12 after aeration.
Goulet et al. showed that, by avoiding high SOC, the anthrone formation rate in 2,6DHAQ decreased substantially.31 They also demonstrated recovery of most of thelost capacity by air exposure. Because DPivOHAQ also decomposes via anthrone
formation, we hypothesize that similar approaches will extend its lifetime significantly. We suggest that a flow cell with decades-long calendar life might be achievable with DPivOHAQ at pH 12.
Another anthraquinone, DBAQ, with higher solubility, was synthesized (Figure 4A)
using a similar strategy. The first step is a Friedel-Crafts acylation, followed by
Wolff-Kishner reduction of the carbonyl groups to methylene. The last step is tooxidize the corresponding anthracene (or 9,10-dihydroanthracene) to the final
anthraquinone form. Compared with the three steps for DPivOHAQ synthesis, four
steps are required for DBAQ synthesis when starting from anthracene. Because of
the low cost of succinic anhydride and given that the Wolff-Kishner reduction iswell-developed in industry, the cost of DBAQ could also be low. Electrochemical kinetics studies of DBAQ were conducted by RDE techniques as shown in Figure S14A.
The diffusion coefficient of the oxidized form of DBAQ was determined by Levich
analysis (Figure S14B) to be 2.5 3 106 cm2 s–1. According to the Koutecký-Levich
equation and Tafel plot, the charge transfer coefficient is 0.50, and the kinetic rate
constant is 2.9 3 103 cm s1, which is still slightly higher than that of DPivOHAQ.
The solubility of DBAQ was determined to be 1.0 M at pH 12, corresponding to a
volumetric capacity of 53.6 Ah L–1 for the negolyte. The reduction potential is0.47 V versus SHE at pH 12. When paired with potassium ferrocyanide, a full cellof approximately 0.97 V can be achieved. To evaluate the stability of a DBAQ negolyte, a full cell was assembled with 5 mL of 0.5 M DBAQ negolyte at pH 12 as the capacity-limiting side and a posolyte comprising 80 mL of 0.3 M K4Fe(CN)6 with 0.1 M
K3Fe(CN)6 at pH 12 as the non-capacity-limiting side. The flow cell was constructedfrom graphite flow plates and carbon paper electrodes, separated by a Fumasep E620 (K) membrane. The cell exhibited 91.7% of its theoretical capacity for the negolyte. It was cycled for 650 cycles at 100 mA cm2, which required 15.5 days to complete. The average capacity fade rate was 0.0084% per day, corresponding to 3.1%
per year. Assuming the capacity fade is primarily due to anthrone formation then,with careful control of the pH and SOC of the DBAQ electrolyte and periodic exposure to air, DBAQ might exhibit an even lower loss rate in real-world applications.

1438 Chem 6, 1432–1442, June 11, 2020

ArticlellFigure 4. DBAQ Synthesis and Electrochemistry(A) Synthetic route for DBAQ.
(B) Cyclic voltammograms of 5 mM DBAQ and 10 mM potassium ferrocyanide at pH 12 with a scan rate of 100 mV s –1 .
(C) Coulombic efficiency and discharge capacity versus time and cycle number for a negolyte-limited DBAQ/K 4 Fe(CN) 6 full cell. The cell was cycledgalvanostatically at 100 mA cm 2 between 0.4 and 1.3 V with SGL 39AA carbon electrodes, and each half cycle was followed by 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 DBAQ at pH 12, and the posolyte comprised 80 mL of
0.3 M potassium ferrocyanide and 0.1 M potassium ferricyanide at pH 12. Note that the left y axis represents only 0.32% of the capacity of the DBAQ
negolyte.

Conclusions
In this report, we have demonstrated a new route to synthesize water-soluble anthraquinones with solubilizing groups attached by carbon-carbon bonds, starting frompotentially inexpensive 9,10-dihydroanthracene. These anthraquinones exhibit
high aqueous solubilities and low capacity fade rates of 0.0084% per day and0.014% per day at pH 12, respectively. We demonstrated in a full cell containing a
DPivOHAQ negolyte that anthrone formation is the major side reaction responsible
for capacity fade and that air exposure can recover most of the lost capacity. Furthermore, by increasing the pH of the negolyte, we demonstrated the suppression of the
DPivOHAQ capacity fade rate to an extremely low value of less than 1% per year. We
expect that the stability of DPivOHAQ and DBAQ can be even further improved withcareful control of the battery operating conditions. We suggest that strategies
combining SOC limit control, precision air exposure, and pH tuning can be extended
to other inexpensive anthraquinone molecules to achieve extremely low capacity
fade rates, paving the way for commercializing anthraquinone-based RFBs to enablegrid-scale energy storage of renewable electricity.

EXPERIMENTAL PROCEDURES
Experiment Materials
9,10-dihydroanthracene (97%), 3,3-dimethylacrylic acid (97%), and anhydrous dichloromethane were purchased from Sigma Aldrich. Anhydrous aluminum chloride(95%) was purchased from Alfa Aesar. All chemicals were used as received.

Chem 6, 1432–1442, June 11, 2020 1439

ll
3,30 -(Anthracene-Diyl)Bis(3-Methylbutanoic Acid) (DPivOHAC)
13.32 g (99.93 mmol) of AlCl3 was suspended in 200 mL of anhydrous CH2Cl2. A
solution of 6.67 g (66.62 mmol) of 3,3-dimethylacrylic acid in 20 mL of anhydrous
CH2Cl2 was added by syringe and the mixture stirred at room temperature (20  C)
for 0.5 h under nitrogen. Subsequently, a solution of 5.00 g (27.74 mmol) of 9,10-dihydroanthracene in 15 mL of anhydrous CH2Cl2 was added to the above mixture
and stirred for 48 h at room temperature (20  C). After that, the solvent was
quenched with 200 mL of 1 M aqueous HCl and stirred overnight. The organic layer
was then removed and the remaining solution was filtered to afford the pale-yellow
product. The yield given was 95%.
3,30 -(9,10-Anthraquinone-Diyl)Bis(3-Methylbutanoic Acid) (DPivOHAQ)
DPivOHAC (6.00 g, 15.85 mM) was dissolved in glacial acetic acid (70 mL). Then, a
CrO3 solution (3.33 g, 33.3 mM) was added to the DPivOHAC and acetic acidmixture. The reaction mixture was heated at 90  C for 1 h. After cooling down toroom temperature (20  C), water was added to precipitate the solid. The compound was purified by dissolution in base followed by addition of acid to afford
the precipitate. The yield given was 85%.
The 1H NMR spectra of 3,30 -(anthracene-diyl)bis(3-methylbutanoic acid) and 3,30 (9,10-anthraquinone-diyl)bis(3-methylbutanoic acid) are shown in Figures S1 andS2. DBAQ-related synthesis uses a similar method, thus is provided in the supporting information.
Full Cell Measurements
Flow battery experiments were conducted with cell hardware from Fuel Cell Tech.
(Albuquerque, NM), assembled into a zero-gap flow cell configuration, similar to aprevious report.17,18 Pyrosealed POCO graphite flow plates with serpentine flow
patterns were used for both electrodes. Each electrode comprised a 5 cm2 geometric surface area covered by one sheet of AvCarb carbon cloth or three sheets of SGL
39AA electrode. For DPivOHAQ-ferrocyanide full cell tests, a Fumasep E-620 (K)membrane was used to serve as the ion-selective membrane between the AvCarb
electrodes. For DBAQ-ferrocyanide full cell tests, a Fumasep E-620 (K) membrane
was used to serve as the ion-selective membrane between the SGL 39AA electrodes.
The outer portion of the space between the electrodes was gasketed by Viton sheetswith the area over the electrodes cut out. Torque applied during cell assembly was
60 lb-in (6.78 N m) on each of eight bolts. The electrolytes were fed into the cellthrough fluorinated ethylene propylene (FEP) tubing at a rate of 60 mL/min,controlled by Cole-Parmer six Masterflex L/S peristaltic pumps. All cells were run inside a nitrogen-filled glove bag. Cell polarization measurements, impedance spectroscopy, and charge-discharge cycling were performed using a Biologic BCS-815
battery cycler. Galvanostatic cycling was performed at G0.1 A cm2 at room temperature (20  C) with voltage limits of 0.6 and 1.3 V. To obtain the polarization
curves, the cell was first charged to the desired SOC and then polarized via linear
sweep voltammetry at a rate of 100 mV s1. 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.

SUPPLEMENTAL INFORMATION
Supplemental Information can be found online at https://doi.org/10.1016/j.chempr.
2020.03.021.

1440 Chem 6, 1432–1442, June 11, 2020

Articlell

ArticleACKNOWLEDGMENTS
This research was supported by U.S. DOE award DE-AC05-76RL01830 through
PNNL subcontract 428977, Innovation Fund Denmark via the Grand Solutions project ‘‘ORBATS’’ file number. 7046-00018B, and NSF grant CBET-1914543. The authors thank Daniel A. Pollack for valuable discussions and for assistance with chemical stability experiments and in reviewing and editing the manuscript.

AUTHOR CONTRIBUTIONS
M.W. designed and synthesized the molecules. M.W. conducted the chemical stability and solubility tests, and M.W., A.A.W., Z.T., and Y.J. performed full cell tests. Y.J.
also conducted DPivOHAQ synthesis. Y.J. and S.J. identified the anthrone sideproduct and performed HPLC tests. E.M.F. performed the RDE test. R.G.G. supervised the molecular synthesis and characterization. M.J.A. supervised the cell assembly, electrolyte composition selection, and performance evaluation. M.W. drafted
the manuscript with input from Y.J. All authors reviewed and edited the manuscript.

DECLARATION OF INTERESTSHarvard University has filed a patent application on the materials described in thispaper.
Received: January 22, 2020
Revised: March 14, 2020
Accepted: March 23, 2020
Published: April 17, 2020

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Chem, Volume 6

Supplemental Information

Extremely Stable Anthraquinone Negolytes
Synthesized from Common Precursors
Min Wu, Yan Jing, Andrew A. Wong, Eric M. Fell, Shijian Jin, Zhijiang Tang, Roy G.
Gordon, and Michael J. Aziz

Supporting information

Contents
Chemical synthesis and characterization ......................................................................... 2
Chemical stability experiments .................................................................................... 10
Cyclic Voltammetry (CV) and Rotating Disk Electrode (RDE) Measurements ................. 11
Solubility tests ........................................................................................................... 12
Thermodynamics of the disproportionation reaction ....................................................... 13
Anthrone detection ..................................................................................................... 171

Experimental section:
9,10-dihydroanthracene (97%), 3,3-dimethylacrylic acid (97%), succinic anhydride,hydrazine monohydrate (98%), anhydrous dichloromethane, and diethylene glycol (99%)
were purchased from Sigma Aldrich, and anhydrous aluminum chloride (95%) was
purchased from Alfa Aesar. All chemicals were used as received.

Chemical synthesis and characterization

Figure S1. 1H NMR spectrum of DPivOHAC in DMSO-d6. Solvent peaks are those that are notintegrated. 1H NMR (500 MHz, DMSO-d6)  11.88 (s, 2H), 8.45 (t, 2H), 7.99 (dd, 2H), 7.91 (d,2H), 7.62 (d, 2H), 2.71 (s, 2H), 1.51 (s, 6H).

2

Figure S2. 1H NMR spectrum of DPivOHAQ in DMSO-d6. Solvent peaks are those that are notintegrated. 1H NMR (500 MHz, DMSO-d6)  11.99 (s, 2H), 8.17 (d, 2H), 8.15 (dd, 2H), 8.0 (dd,2H), 2.75 (s, 2H), 1.46 (s, 6H).

DPivOHAQ synthesized with nitric acid/oxygen method:1
3 g DPivOHAC (7.94 mmol) was taken in a flask containing 60 mL of acetic acid, refluxed
at 95 °C and 12.5 mL of a solution containing conconcentrated nitric acid (70%), acetic
acid and water (1:9:9, by volume) was added dropwise to the well-stirred mixture, while
industrial oxygen was bubbled into the solution. The addition took about 10 min, deepbrown fumes evolved. After the reaction was completed (normally within 2 h) the flask
was cooled and 1M HCl was added to the solution, and filtrate to get the final DPivOHAQ.
Washed with acetone, the final yield is around 80%.

3

Figure S3. 1H NMR spectrum of DPivOHAQ synthesized from nitric acid/oxygen methodin DMSO-d6. Solvent peaks (3.3 H2O, 2.5 DMSO, 1.9 CH3CO2H) are those that are notintegrated. 1H NMR (500 MHz, DMSO-d6)  11.99 (s, 2H), 8.17 (d, 2H), 8.15 (dd, 2H),
8.0 (dd, 2H), 2.75 (s, 2H), 1.46 (s, 6H).

4,4'-(9,10-dihydroanthracene-diyl)bis(4-oxobutanoic acid) (DOBDHAC):
15.50 g (115.89 mmol) of AlCl3 was suspended in ~200 mL of anhydrous CH2Cl2. A
solution of 5.69 g (56.49 mmol) of succinic anhydride was added and the mixture stirred
at 0 C for 0.5 hour under nitrogen. Subsequently, a solution of 5.00 g (27.74 mmol) of
9,10-dihydroanthracene in ~15 mL of anhydrous CH2Cl2 was added to the above mixture
and stirred overnight. Then the suspension was filtered to obtain the red solid, which wasthen washed with ice water twice and filtered to afford a yellow solid (10.5 g). Yield: 99%.
1
H NMR (500 MHz, DMSO-d6)  12.13 (s, 2H), 7.96 (t, 2H), 7.84 (dd, 2H), 7.49 (d, 2H),
4.09 (s, 4H), 3.25 (t, 4H), 2.58 (t, 4H).

4

Figure S4. 1H NMR spectrum of DOBDHAC in DMSO-d6. Solvent peaks are those that are notintegrated.

4,4'-(9,10-dihydroanthracene-diyl)dibutanoic acid (DBDHAC):
To a solution of DOBDHAC (10.5 g, 27.74 mmol) in diethylene glycol (100 mL) at room
temperature was added hydrazine monohydrate (98%, 8 mL, 166 mmol). The reaction
mixture was stirred for half an hour. Then potassium hydroxide was added (8.73 g, 155.59mmol) and stirred at 100 C for 1 hour under nitrogen. Then the temperature was allowed
to rise to 200 C for 4 hours. After cooling to room temperature, 500 mL of water wasadded to the solution, which was then filtered twice to remove insoluble solid. The filtrate
was acidified with 5 M HCl to adjust the pH to 1, and the resulting precipitate was collected
and washed with water (500 mL) twice to afford the light brown solid. Yield:
approximately 95%. 1H NMR indicated that the product was DBDHAC (4,4'-(9,10dihydroanthracene-2,6-diyl)dibutanoic acid) with around 10% percent of DBAC (4,4'(anthracene-diyl)dibutanoic acid). As both dihydroanthracene and anthracene may be
converted to anthraquinone, no further purification was needed at this step. 1H NMR (500
MHz, DMSO-d6) DBHAC:  7.21 (d, 2H), 7.13 (s, 2H), 6.99 (dd, 2H), 3.83 (t, 4H), 2.55
5

(t, 4H), 2.20 (t, 4H), 1.78 (m, 4H). DBAC:  8.43 (t, 2H), 7.99 (dd, 2H), 7.80 (s, 2H), 7.37
(t, 2H), 2.79 (t, 4H), 2.29 (t, 4H), 1.93 (m, 4H).

Figure S5. 1H NMR spectrum of DBDHAC in DMSO-d6. Solvent peaks are those that are notintegrated.

4,4'-(9,10-anthraquinone-diyl)dibutanoic acid (DBAQ):
DBDHAC (10 g, 28.4 mM) was dissolved in 100 mL glacial acetic acid. Then, a CrO3
solution, prepared by dissolving CrO3 (7.1 g, 71 mM) in 4 mL of water, was added to theabove solution. The reaction mixture was allowed to stir at room temperature overnight.
Then 300 mL of water was added to the mixture, which was then filtered to afford theyellow product. The product was dissolved in base and further acidified to afford the yellow
precipitate. The last step is used to remove the residual Cr (III) compound. Yield: 90%. 1H
NMR (500 MHz, DMSO-d6)  8.05 (dd, 2H), 7.94 (t, 2H), 7.70 (dd, 2H), 2.77 (t, 4H), 2.27
(t, 4H), 1.86 (m, 4H).
6

Figure S6. 1H NMR spectrum of DBAQ in DMSO-d6. Solvent peaks are those that are notintegrated.

DBAQ synthesized with the nitric acid/oxygen method:1
3 g DBDHAC (8.52 mmol) was taken in a flask containing 60 mL of acetic acid, refluxed
at 95 °C and 13.5 mL of a solution containing conconcentrated nitric acid (70%), acetic
acid and water (1:9:9, by volume) was added dropwise to the well-stirred mixture, while
industrial oxygen was bubbled into the solution. The addition took about 10 min, deepbrown fumes evolved. After the reaction was completed (normally within 2 h) the flask
was cooled and 1M HCl was added to the solution, and filtrate to get the final DBAQ.
Yield: 97%.

7

Figure S7. 1H NMR spectrum of DBAQ synthesized from the nitric acid/oxygen method in
DMSO-d6. Solvent peaks (3.3 H2O, 2.5 DMSO) and small impurities are those that are notintegrated. 1H NMR (500 MHz, DMSO-d6)  8.05 (dd, 2H), 7.94 (t, 2H), 7.70 (dd, 2H), 2.77 (t,4H), 2.27 (t, 4H), 1.86 (m, 4H).
Note that the reaction condition is not optimized, and the impurities in the final compound isconsider to come from the anthracene core rather than from the functionalized chain.

8Comparison of synthetic conditions for DPivOHAQ, DBAQ, DBEAQ, and DPPEAQ

Scheme S1. Comparison of the synthetic routes for DBEAQ,2-5 DPPEAQ,2-4,6 DPivOHAQ,7 andDBAQ7 from anthracene.

Table S1. Lab-scale cost for the side chain resources and starting materials of four stable
anthraquinones from Sigma-Aldrich in Nov. 2019.
2,6-DHAQDBEAQDPPEAQ
9,10-dihydroanthracene
DPivOHAQ
DBAQ
ethyl 4-bromobutyrate diethyl-3-bromopropylphosphonate
3,3-dimethylacrylic acid succinic anhydride
side chaincost
5 g: $105
250 g: $138
5mL: $91
5 g: $41.1
500 g: $152
500 g: $41
unit price ($/mol) 5044.4
107.7
3498
1481
30.4
8.2

Note that even though the lab-scale cost of 2,6-DHAQ is high, the mass production costof 2,6-DHAQ is predicted to be as low as $2.4/kg by Borealis Technology Solutions
9LLC.8 As the lab-scale cost of 9,10-dihydroanthracene is approximately one third of thatof 2,6-DHAQ, we suppose the cost of 9,10-dihydroanthracene could be inexpensive if
produced in large-scale.

Chemical stability experiments
Samples of DPivOHAQ at 0.1 M concentration and at pH 14 were stored in fluorinatedethylene propylene bottles and heated in an oven at 65 C for 8 days. The extent of
decomposition was determined by 1H NMR, with peak integrals measured relative to aninternal standard of NaCH3SO3 prepared at 10 mM concentration in D2O. All samples werediluted in this deuterated solvent containing the internal standard at a fixed ratio of 1:5.

Figure S8. 1H NMR spectra (500 MHz, 1 M KOD in D2O with 10 mM NaCH3SO3 internal
standard) of (a) the oxidized form of DPivOHAQ in pH 14 aqueous solution (1 M KOH); (b)DPivOHAQ treated at 65 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 theoxidized form; (c) DPivOHAQ treated at 65 C for 8 days at 0.1 M concentration in pH 14
10aqueous solution (1 M KOH) in the oxidized form. No apparent decomposition was detected; theintegration did not appreciably change.

Cyclic Voltammetry (CV) and Rotating Disk Electrode (RDE) Measurements
Glassy carbon was used as the working electrode for all three-electrode CV tests. RDE
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 a 3 M NaCl solution), and a graphite counterelectrode. The diffusion coefficient of the oxidized form of DPivOHAQ was calculatedusing the Levich equation, which relates the mass-transport-limited current to the numberof electrons transferred (n), the area of the electrode (A), and the concentration of redoxactive species in the electrolyte (C) by plotting the mass-transport-limited current against
the square root of the rotation rate (Figure S9b) with the following parameters: n = 2, F =96,485 Coulombs/mol, A = 0.196 cm2 , C = 5 mM, kinematic viscosity of 1 M KCl =0.89 10-6 m2/s.9 The resulting value of the diffusion coefficient for the oxidized form of
DPivOHAQ is 2.4 × 10-6 cm2/s.

Figure S9. DPivOHAQ reduction kinetics in RDE. a) Linear sweep voltammograms of 5 mM
DPivOHAQ in 1 M KCl at pH 12 on a glassy carbon electrode at rotation rates between 400 and
112000 rpm. b) Levich plot (limiting current versus square root of rotation rate in rad/s) of 5 mM
DPivOHAQ in 1 M KCl at pH 12. Limiting current is taken as the current at -0.65 V in (a). The
slope yields a diffusion coefficient for the oxidized form of DPivOHAQ of 2.39 × 10-6 cm2 s-1. c)
Koutecky-Levich plot (reciprocal current versus inverse square root of rotation rate in rad/s) of 5mM DPivOHAQ in 1 M KCl at pH12. d) Fitted Tafel plot of 5 mM DPivOHAQ in 1 M KCl at pH
12. The charge transfer coefficient is calculated to be 0.49, and the rate constant is calculated to be
2.48 × 10-3 cm s-1.

Solubility tests
For the solubility of K2DBAQ, we prepared a 1.0 M solution in pH 12 KOH and found thatit remained fully dissolved. We did not determine an upper limit. Therefore we report thesolubility as 1.0 M.
The solubility limit of DPivOHAQ was measured in the oxidized form by adding thepotassium salt of DPivOHAQ (prepared by reacting DPivOHAQ with potassium hydroxide
in water) until no further solid could be dissolved. The mixture was adjusted to pH 12.
After the suspension was filtered through a nylon 0.45 μm syringe filter, a saturated
solution of DPivOHAQ at pH 12 was obtained. The saturated solution was then diluted
20,000 times while maintaining a pH of 12, and the concentration was evaluated by UV–
Vis spectrophotometry (Agilent Cary 60 spectrophotometer). The concentration was
calculated according to a pre-calibrated absorbance concentration curve of known
concentrations of DPivOHAQ at pH 12. The resulting value of the solubility of the oxidized
form of DPivOHAQ at pH 12 is 0.74 M.
Solubility of the oxygen-sensitive reduced form of DPivOHAQ was not measured, but was
expected to be higher than the oxidized form because: (1) no precipitation was observed
after full electrochemical reduction of DPivOHAQ; (2) we expect the two extra negative
charges on the reduced form to make quinone-quinone interactions unfavorable andincrease its solubility;10and (3) anthraquinone is known to be insoluble in base, whereas
the reduced form of anthraquinone is readily dissolved in alkaline solution and is often
called soluble anthraquinone.11

12

Figure S10. a) UV–Vis spectra of DPivOHAQ at different concentrations; b) the absorbance at
287.9 nm versus the concentration; a least-squares linear fit to the data was performed to generate
the calibration curve utilized in this work. The saturated sample was diluted 20,000 times andcalculated to be 0.74 M at pH 12.

Thermodynamics of the disproportionation reaction
When the pH is between 12 and 14, both 9,10-dihydroxyanthracene (reduced
anthraquinone) and anthrone (pKa = 10)12 are deprotonated, and the disproportionation
reaction proceeds as follows:O+ 2 H2O

2OOO+ 3 OH -+OThis disproportionation reaction is a combination of the following two half reactions:

13

O

O

half reaction 1

+ 2eO

O

O

O
+ 2 H2O + 2e-+ 3 OH-half reaction 2

O

Within this pH range, the redox potential E1 for half reaction 1 is pH independent,
whereas the redox potential E2 for the half reaction 2 decreases by approximately 89 mV
per pH unit increase based on Nernst equation. The Gibbs free energy change ∆G per
molecule of anthrone formed by the disproportionation reaction (half reaction 2 minus
half reaction 1) is given by: G   2  F  ( E 2  E1 ) / N A

assuming reactants and products (other than OH-) have unit activities and the activity of
OH- is consistent with the pH.
As pH increases, the Gibbs free energy change ∆G increases, and the disproportionation
reaction is less prone to happen. This observation is equivalent to the statement that theoverall reaction within this pH range involves OH-, so the activity of OH- factors into thereaction quotient and equilibrium constant. By Le Chatelier’s principle, the reaction will
be disfavored as the pH increases.

When the pH is below the pKa1 (typically approximately 7) of 9,10-dihydroxyanthracene,
the disproportionation reaction becomes:

The two corresponding two half reactions are

14

The redox potential E3 for half reaction 3 decreases by approximately 59 mV per pH unit
increase based on Nernst equation, and so does the redox potential E4.
The Gibbs free energy change ∆G′ per molecule of anthrone formed by thedisproportionation reaction (half reaction 4 minus half reaction 3) is given by:
 G '   2  F  ( E 4  E3 ) / N A

assuming reactants and products (other than H+) have unit activities and the activity of H+
is consistent with the pH.
As pH (below pKa1) increases, the Gibbs free energy change ∆G′ is constant. Thisobservation is equivalent to the statement that the overall reaction within this pH range
does not involve H+, so the activity of H+ does not factor into the reaction quotient andequilibrium constant.

If we draw a theoretical Pourbaix diagram for the two half reactions of the anthroneforming disproportionation reaction according to Nernst equation (assuming that the pKa
of anthrone lies between the pKa1 and pKa2 of hydroquinone), we obtain:

15

Figure S11. Representative Pourbaix diagram of anthraquinone, anthrahydroquinone, and anthrone.
The pKa1 and pKa2 belong to the 9,10-dihydroxyanthracene (anthrahydroquinone). The pKa(reported to be approximately 10) belongs to anthrone.12 Here we assume that the pKa of anthrone
lies between the pKa1 and pKa2 of anthrahydroquinone.

This Pourbaix diagram illustrates that the Gibbs free energy change ∆G for thedisproportionation reaction is larger at alkaline pH than at neutral or acidic pH. Therefore,
anthrone formation is disfavored at alkaline pH relative to neutral or acidic pH. In realapplications with a reasonable concentration of quinone, even starting at pH 7, the pH will
increase to (or above) the pKa2 of 9,10-dihydroxyanthraquinone upon quinone reduction;
therefore, it is still possible to get a relatively stable anthraquinone negolyte starting at pH
7.

16

Anthrone detection

Figure S12. 1H NMR spectra of the reduced and aerated DPivOHAQ in DMSO-d6. Sample
preparation: 5 mL of 0.1 M DPivOHAQ was fully charged to ~100% SOC and stored in a brownvial in a glovebox for 238 days. After that, an aliquot was taken and acidified with concentrated
HCl to obtain the protonated and reduced DPivOHAQ precipitate after clear solution was removed.
The protonated and reduced DPivOHAQ was re-dissolved in DMSO-d6 and stored in a J. Young
NMR tube for 1H NMR spectrometry. The aerated sample was prepared by exposing the same
reduced sample to air and shaking the NMR tube for several minutes before collecting the NMR
spectrum. The 1H NMR spectrum of the reduced sample revealed, in addition to peaks from reduced
DPivOHAQ and a small amount of the oxidized form, some appreciable peaks highlighted with red
circles, corresponding to signals from unknown compounds. However, these signals disappeared
when the same NMR sample was aerated, indicating that the unknown compounds can either beoxidized back to DPivOHAQ or convert into other products with no observable signals above thedetection limit of the NMR instrument.

17

Figure S13. LC-MS results from the reduced DPivOHAQ samples stored for 238 days. a) The basepeak chromatogram of the aerated sample. b) The extracted-ion chromatogram for the anthrone
derivative from the aerated sample. c) The extracted-ion chromatogram for the anthrone dimer
derivative from the aerated sample. d) The base peak chromatogram of the reduced sample. e) The
extracted-ion chromatogram for the anthrone derivative from the reduced sample. f) The extractedion chromatogram for the anthrone dimer derivative from the reduced sample. In the reduced
sample, both anthrone and anthrone dimer forms were detected. After aeration, no anthrone form
is detected. The anthrone dimer accounts for 1.24% in the aerated sample, corresponding a faderate of 1.90% per year for the reduced DPivOHAQ after aeration. Sample preparation: 5 mL of 0.1
M DPivOHAQ was fully charged to ~100% SOC and stored in a brown vial in a glovebox for 238
days. After that, aliquots were taken and acidified with concentrated HCl to obtain protonated
DPivOHAQ precipitate after clear solution was removed. The protonated and reduced DPivOHAQ
was re-dissolved in DMSO. The resulting solution was further diluted to the desired concentration(10-20 M) by acetonitrile/water co-solvents (V/V=1:1). One sample was stored in a glovebox in
the reduced form prior to the LC-MS measurement, whereas another sample was intentionallyaerated. After that, the two samples were immediately subjected to the LC-MS experiment.

18

Figure S14. RDE test for DBAQ. a) Linear sweep voltammograms of 5 mM DBAQ in 1 M KCl at
pH 12 on a glassy carbon electrode at rotation rates between 400 and 2000 rpm. b) Levich plot
(limiting current versus square root of rotation rate in rad/s) of 5 mM DBAQ in 1 M KCl at pH 12.
Limiting current is taken as the current at -0.65 V in (a). The slope yields a diffusion coefficient forthe oxidized form of DBAQ of 2.54 × 10-6 cm2 s -1. c) Koutecky-Levich plot (reciprocal current
versus inverse square root of rotation rate in rad/s) of 5 mM DBAQ in 1 M KCl at pH 12. d) Fitted
Tafel plot of 5 mM DBAQ in 1 M KCl at pH 12. The charge transfer coefficient is calculated to be
0.50, and the rate constant is calculated to be 2.87 × 10-3 cm s-1.

19

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