Article pubs.acs.org/cm
Concentration-Dependent Dimerization of Anthraquinone Disulfonic Acid and Its Impact on Charge Storage Thomas J. Carney,†,‡ Steven J. Collins,§ Jeffrey S. Moore,†,∥ and Fikile R. Brushett*,†,§ †
Joint Center for Energy Storage Research, 9700 S. Cass Avenue, Argonne, Illinois 60439, United States Department of Materials Science and Engineering and §Department of Chemical Engineering, Massachusetts Institute of Technology, 77 Massachusetts Avenue, Cambridge, Massachusetts 02139, United States ∥ School of Chemical Sciences and Beckman Institute for Advanced Science and Technology, University of Illinois at Urbana−Champaign, Urbana, Illinois 61801, United States ‡
S Supporting Information *
ABSTRACT: 9,10-Anthraquinone-2,7-disulfonic acid (AQDS) is considered a benchmark active material for aqueous organic redox flow batteries. At low concentration, AQDS demonstrates two-electron transfer at near ideal electrochemical reversibility; however, at higher concentration, AQDS displays more complex behavior presumably due to the emergence of intermolecular reactions. Here, we systematically examine the electrochemical and physical properties of AQDS solutions using a suite of electrochemical, analytical, and spectroscopic techniques. Depending on the AQDS pretreatment, concentration, solution pH, and electrolyte composition, coupled chemical and electrochemical reactions lead to different charge storage capabilities. To elucidate the underlying cause of these differences, we performed various pretreatments of AQDS, examined chemical speciation by NMR, and investigated the corresponding electrochemical properties through cyclic voltammetry and bulk electrolysis. In all cases, reversible intermolecular dimerization was detected at solution concentrations greater than 10 mM. Moreover, we found that the charge state of the formed dimers was dependent on the AQDS pretreatment and the solution pH. Under acidic conditions, 1.5 electrons per molecule of AQDS were reversibly accessible, whereas under buffered mild-alkaline conditions, only one electron per molecule of AQDS was accessible. Because of insufficient proton concentration, AQDS did not cycle reversibly in unbuffered neutral electrolyte. Even when employing chemical oxidants during a chemical titration, charge storage of two electrons per molecule could not be realized. We hypothesize that adduct formation between AQDS and CO2, along with solution pH, play important roles in the charge accessibility.
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design space.12 Moreover, although most transition metal salts can only reversibly store one electron, several organic families undergo two-electron redox chemistry, hence doubling the charge storage capacity of the electrolyte. Although many new organic redox couples have been proposed, to date, quinones have been the most successful. Historically, quinones have been used in a variety of industrial applications including hydrogen peroxide production,13 dyes,14 and medicine.15 Recently, quinones have found applications as sensors to understand better biological electron transport processes.16,17 Accordingly, significant work has been undertaken in the past decade to understand the fundamental electrochemistry of quinones.18−20 In one such case, Batchleor-McAuley et al. reported the electrochemical properties of 9,10-anthraquinone-2,6-disulfonic acid (AQDS-2,6) at low concentrations (1 mM), and developed a scheme of squares mechanistic pathway at various
INTRODUCTION Energy storage is expected to play a significant role in meeting emerging needs in the evolving electric power sector, including facilitating the integration of intermittent renewable sources and enhancing the efficiency of nonrenewable energy processes.1 Redox flow batteries (RFBs) are rechargeable electrochemical devices that are well-suited for multihour energy storage and offer several key advantages over enclosed battery technologies (e.g., lithium (Li)-ion, lead-acid) including independent scaling of power and energy, long service life, improved safety, and simplified manufacturing.2 Current stateof-the-art systems, however, have yet to achieve wide-scale commercial success due to technical and economic challenges. This, in turn, has spurred research into alternative chemistries with improved properties. Recently, a new class of redox couples based on organic materials has emerged.3−11 Organic materials are of interest for flow battery applications as electrochemical (e.g., redox potential) and physical properties (e.g., solubility) are tunable through functionalization, thereby expanding the molecular © 2017 American Chemical Society
Received: February 14, 2017 Revised: May 19, 2017 Published: May 19, 2017 4801
DOI: 10.1021/acs.chemmater.7b00616 Chem. Mater. 2017, 29, 4801−4810
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Chemistry of Materials pH.16 In another, Quan et al. examined 1,4 benzoquinone at 1 mM in numerous electrolytes to deduce the effects of hydrogen bonding in aqueous quinone electrochemistry.17 Both Quan et al. and Batchleor-McAuley et al. found a two-electron, twoproton transfer in electrolytes where [H+] was greater than the concentration of the active species and a two-electron transfer without proton transfer when the [H+] was less than that of the active species.16,17 Building on these fundamental studies, researchers have sought to implement quinones in RFBs.3−6,21−27 9,10Anthraquinone-2,7-disulfonic acid (AQDS) is perhaps the most promising and widely reported derivative due to its good charge retention (days), low crossover rates (through Nafion), high solubility in relevant aqueous electrolytes (>1 M), and low redox potential (ca. 200 mV vs RHE).3,16 Huskinson et al. first coupled AQDS, as a negative electrolyte, and bromine (Br2), as a positive electrolyte, to demonstrate a metal-free organic−inorganic aqueous flow battery with a peak power density of 600 mW cm−2 at 1.3 A cm−2 operating at 40 °C.3 This work was extended by Yang et al., who coupled 9,10anthraquinone-2,6-disulfonic acid (AQDS-2,6), an isomer of AQDS, with 1,2-dihydrobenzoquinone (Tiron) to fabricate an all organic aqueous flow battery with a peak power density of ca. 50 mW cm−2 at 100 mA cm−2 operating at rt.6 In these reports, the anthraquinone disulfonic acid derivatives were characterized, via cyclic voltammetry, at low concentration (ca. 1 mM), and then employed in a flow cell at high concentration (ca. 0.5−1 M) for discharge/charge cycling. Interestingly, although the characteristic two-electron transfer behavior was observed at low concentration, the reported flow cells did not demonstrate the expected charge storage capacity. Specifically, assuming two electron transfer, Huskinson et al. and Yang et al. were only able to access 50% and 33% of the theoretical charge storage capacity of their respective full cells.3,6 We hypothesize that this discrepancy is due to the changes in the electrochemical and chemical behavior of active materials as a function of concentration. However, this is difficult to confirm unambiguously in an operating flow cell as multiple factors contribute to deviations from theoretical capacity including incomplete polarization due to selected operating conditions, loss of active species concentration due to membrane crossover or leaking from the experimental apparatus, and molecular decomposition due to undesirable (electro)chemical side reactions. To mitigate these confounding variables, we investigated the concentration-dependent properties of solution-phase AQDS, in isolation, using a combination of electrochemical, analytical, and spectroscopic techniques. These studies not only reveal concentration-dependent intermolecular dimerization above 10 mM but also highlight the important role of AQDS pretreatment and solution pH on the charge state of the formed dimers. These findings have direct ramifications on accessible electrons per AQDS and thus the charge storage capabilities of any resulting redox electrolyte.
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colored solution. A saturated solution of NaCl was added to the reaction mixture and the resulting solution was slowly concentrated. The sample, hereafter referred to as AQDS-H2SO4, precipitated out as a mustard-yellow product, which was filtered, washed with DI water and dichloromethane (Sigma-Aldrich, 99.5%), and dried under vacuum at 100 °C for 24 h. No impurity peaks were observed during NMR analysis of AQDS-H2SO4 and thus no further purification steps were undertaken (Figure S4). In the second pretreatment, AQDS was flushed three times through a column containing Amberlyst 15H ionexchange resin (Sigma-Aldrich) to replace the sodium cation with a proton.3,6 This solution was then evaporated to recover dark-yellow, solid product, hereafter referred to as AQDS-IER, which was dried under vacuum at 100 °C for 24 h. No impurity peaks were observed during NMR analysis of AQDS-IER (Figure S3) and thus no further purification steps were undertaken. For comparison, two additional anthraquinone derivatives were investigated. 9,10-Anthraquinone-2,6disulfonic acid disodium salt (AQDS-2,6; 98%) was used as received from Santa Cruz Biotechnology and 9,10-anthraquinone-2-sulfonic acid sodium salt monohydrate (AQDS-2; 97%) was used as received from Sigma-Aldrich. Sodium sulfate decahydrate (99%, Alfa Aesar) was purified via conventional techniques28 by dissolving sodium sulfate in water at 30 °C (1.1 mL/g), cooling to 0 °C, collecting the precipitant, and drying the recovered salt under vacuum at room temperature. Sodium carbonate decahydrate (99+%, Alfa Aesar) and sodium hydrogen carbonate (99%, Alfa Aesar) were used as received. Electrochemical Characterization. Cyclic Voltammetry. Cyclic voltammetry (CV) experiments were performed in a three-electrode cell composed of a glassy carbon rod encased in polychlorotrifluoroethylene (PCTFE) as a working electrode (CH Instruments), a platinum coil as a counter electrode (CH Instruments), and a silver/ silver chloride electrode (Ag/AgCl, BASi) as the reference electrode filled with 3 M NaCl (0.209 vs SHE). The working electrodes were prepared immediately prior to each CV by thoroughly polishing on 0.05 μm alumina powder and sonicating in deionized water for 2 min. All experiments were conducted at rt. Argon gas (Airgas UHP 300, 99.999%) was bubbled through a 1/16 in. diameter PTFE tube into the electrolyte for at least 10 min prior to the start of the experiment, and then moved above the degassed solution during the experiment to create an inert blanket. The ohmic drop was measured to be less than 5 Ω, which translated to an iR drop of less than 1 mV for measured currents in the 10−100 μA range. Scan rates of 10, 50, and 100 mV s−1 were used, with potential limits 500 mV above and below the redox potential. At each rate, five consecutive scans were taken as a preliminary measure of stability. The potential was either referenced against the measured experimental Ag/AgCl potential or translated to the RHE scale by measuring the pH of the solution. For all CVs, the open circuit potential was more positive than the redox potential and thus we started with a reductive sweep. Large Volume Bulk Electrolysis. Constant current electrolysis was conducted using 50 mL of electrolyte with various concentrations of AQDS and supporting salt. The solution was cycled in a commercial three-electrode cell (BASi) with a reticulated vitreous carbon (RVC) working electrode, a platinum mesh counter electrode isolated by a glass frit (Ace Glass, Porosity E, 4−8 μm pore size), and a Ag/AgCl reference electrode. Currents were chosen based on a C-rate of C/1.5 or C/3, which corresponds to a current necessary to discharge the cell in 1.5 or 3 h, respectively. To calculate the C-rate, in a separate experiment, the capacity of the material (1 mol e− per mol AQDS, 1.5 mol e− per mol AQDS) was determined by measuring the capacity after one cycle and then iteratively adjusting the current until the correct discharge time was achieved. This ensures the active material at a constant concentration is polarized for the same duration. The voltage limits were set 500 mV above and below the redox potential for consistency with the CV experiments. The solution was stirred at 1,000 rpm to mitigate mass transfer limitations and resulting overpotentials. The collection efficiency of the cell was determined by reducing and reoxidizing a 50 mM solution of potassium hexacyanoferrate(III) (K3Fe(CN)6; 99.0% Sigma-Aldrich) in 0.5 M Li2SO4 (99% Sigma-Aldrich, recrystallized 3 times from DI water) at C/3. A Coulombic efficiency of 99.9% was measured at a capacity of 1
EXPERIMENTAL SECTION
Materials Synthesis. 9,10-Anthraquinone-2,7-disulfonic acid disodium salt (AQDS), red powder, was used as received from Combi-Blocks, 95% (Sigma-Aldrich). As a comparison, AQDS was recrystallized three times from deionized H2O and dried overnight under vacuum at 100 °C. No difference in the electrochemical performance, nuclear magnetic resonance (NMR) spectrum (Figure S2), and color was observed between recrystallized and as-received AQDS. In addition, two pretreatments were investigated. In the first pretreatment, AQDS was dissolved in 0.5 M H2SO4, yielding an amber 4802
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Chemistry of Materials Scheme 1. Summary of Pretreatments for AQDS Examined in This Study
Figure 1. 1H NMR spectrum (400.00 MHz) of 5 mg of AQDS, AQDS-IER, or AQDS-H2SO4 in 1 mL of (a) D2O and (b) d6-DMSO vs TMS. The full NMR spectra is in the Supporting Information (Figures S2−4). (c) 1H NMR peak position of AQDS (noted by the different symbols) in D2O plotted against the log of concentration and fitted to a dimer model (Table S1). AQDS-IER and AQDS-H2SO4 lie on the same trend line (Figure S6d). mol e− per mol K3Fe(CN)6. Argon gas was bubbled through the
Small Volume Bulk Electrolysis. Constant current electrolysis was conducted using 8 mL of electrolyte (4 mL in each chamber), with various concentrations of AQDS and supporting salt, placed in both
solution at least 50 min prior to and throughout the entire experiment. 4803
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Chemistry of Materials sides of a custom electrolysis cell separated by a central glass frit (P5, Adams and Chittenden, 1.0−1.6 μm pore size) (Figure S1).29 A reticulated vitreous carbon (RVC) working electrode and a Ag/AgCl reference electrode were inserted into one compartment of the cell, whereas a RVC counter electrode was inserted in other compartment. Current was chosen based on a 1 C rate, which corresponds to a discharge time of 1 h. The potential limits were set 200 mV above and below the redox potential to ensure full access of the solution capacity. Because the ohmic drop in the small volume bulk electrolysis cell was lower than the large volume bulk electrolysis cell, full capacity is accessed within a smaller potential window. The solutions on both sides were stirred at 600 rpm to mitigate mass transfer limitations. Argon was bubbled through the solution on the working electrode side for at least 5 min prior to and throughout the entire experiment. The RVC electrodes used were constructed by manually cutting a 1 cm × 4 cm piece of 45 ppi RVC (Aerospace Corporation). The top of the RVC electrode was then wrapped with platinum wire as a connector, which was fixed in place with heat-shrink tubing. Again employing K3Fe(CN)6 as a model one electron compound, the small BE cell showed identical behavior to the large BE cell, in terms of both collection efficiency and accessed capacity. Ag/AgCl Calibration. A hydrogen reference electrode (Hydroflex, EDAQ) was placed into a buffered solution (pH = 1.68, Oakton) and allowed to soak overnight. A Ag/AgCl electrode (BASi) was placed into the solution and allowed to approach equilibrium for at least 30 min. The potential between the two electrodes was monitored until the drift was less than 1 mV. pH Meter Calibration. An ORION Star A215 Meter was used to measure the pH of the solutions during the electrochemical experiments. The meter was calibrated using buffer solutions (Oakton) with pH = 12.46, 10.01, 7.00, 4.01, and 1.68. Materials Characterization. Nuclear Magnetic Resonance Spectroscopy. 1H NMR and 13C NMR spectra were acquired using a Bruker Avance 400. NMR solvents, D2O (D 99.96%) and deuterated DMSO (C2D6OS, D 99.96%), were used as received from Cambridge Isotope Laboratories. NMR spectra were aligned using the residual solvent peak and reported vs 4,4-dimethyl-4-silapentane-1-sulfonic acid (DSS) for D2O and tetramethylsilane (TMS) for DMSO.30 The temperature of the NMR tube was measured by the peak separation between the HDO peak and the DSS peak.30 Thermogravimetric Analysis. TGA spectra were collected using a TA Instruments SDT Q600. Approximately 10 mg of the particular AQDS derivative was added to an alumina pan and set to ramp at 5 °C/min to 1300 °C under nitrogen (99.999%, Airgas). The mass of the active material in the derivative was determined by measuring the mass fraction at 150 °C, where a plateau existed for all derivatives once residual water and organic impurities had been removed. Mass Spectrometry. Mass spectra were obtained using a Bruker APEXIV 4.7t FT-ICR-MS (Fourier-Transform Ion Cyclotron Resonance Mass Spectrometer). The Bruker APOLLO ESI (electro spray ion) source, with negative ions, was used to obtain an exact mass (error < ±0.003 amu). Chemical Titration. Na2S2O4 (85%, Sigma-Aldrich) was used as received to chemically reduce AQDS in a basic environment. An aliquot of the solution was transferred using Schlenk line needle techniques to a degassed acidic solution. A reticulated vitreous carbon (RVC) working electrode and a Ag/AgCl reference electrode were placed in solution to monitor the open circuit potential. The solution was titrated with FeCl3 (99.99% Sigma-Aldrich) as a chemical oxidant using a syringe pump (Harvard Apparatus, PHD Ultra). KMnO4 (99.0% Sigma-Aldrich) was also used as a titrant and calibrated against Na2C2O4 (99.5% certified by BAM, according to ISO 17025, certificate U9-1013).
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after this pretreatment, which suggests a difference in chemical structure.3,4,6 AQDS-H2SO4 was chosen to compare the electrochemical and chemical properties between pristine AQDS and AQDS that had been exposed to an acidic environment commonly seen in full cell experiments. We performed 1H NMR and 13C NMR on the three preparations and studied their chemical structure, determining that the AQDS self-associates into a dimer at concentrations above 10 mM. We observed the three derivatives to be partially reduced in solution due to the formation of a CO2 adduct. The electrochemical properties of the derivatives were examined using cyclic voltammetry and bulk electrolysis, to determine how the formation and nature of the dimers impact chargestorage capabilities. Additional spectroscopic and analytical techniques were used to support key observations and hypotheses. Chemical Characterization of AQDS Derivatives. To understand the chemical structure of the various AQDS derivatives, we acquired 1H NMR and 13C NMR spectra in deuterated H2O (D2O) and DMSO (d6-DMSO). In D2O, we observed three chemically unique hydrogens (Figure 1a) that match the predicted pattern by symmetry, simulation, and literature (Table S2).31 The 13C NMR spectrum of the derivatives displayed eight peaks that match the predicted spectrum and confirm the presence of AQDS (Figure S5 and Table S3). We also observed the presence of small multiplets between δ 8 and 9 ppm, which are indicative of possible impurities or alternative states of ionization or association. In d6-DMSO, we observed the same three chemically unique hydrogens (Figure S2−4), and an additional broad peak at δ 6.00 ppm (Figure 1b) for AQDS-IER. This peak’s appearance in d6-DMSO, absence in D2O, broadness, and position suggests that it represents an easily exchangeable, phenol-type proton.31 Given that the chemical shift of an aryl sulfonic acid is 11−12 ppm, we dismissed the possibility of assigning this peak to the sulfonic acid group.31 Both AQDS and AQDS-H2SO4 show the same 1H NMR and 13C NMR spectrum in both D2O and DMSO indicating that, in solution, they have a similar chemical structure. During the preparation of the NMR solutions, we observed a strong concentration dependence of the three 1H NMR peak positions (Figure 1c). As the concentration is increased, the peaks shift upfield, indicating greater shielding of the atomic nuclei.32 To probe systematically this dependence, we prepared solutions from 0.1 to 750 mmol/kg. At low concentrations, approaching an infinite dilution limit, 0.1 to 10 mmol/kg (ca. 0.11 to 11.1 mM assuming no volume change), the peak positions are constant, indicating no intermolecular reactions. As the concentration is increased to 10 mmol/kg (11.1 mM), the peak positions reach an inflection point indicating the onset of intermolecular reactions and formation of an adduct in solution. The formation of the adduct in solution occurs instantaneously upon addition of AQDS to the solution as manifested by the weighted average chemical shift of the of the monomer and polyaggregate chemical shift.33 We fit this dependence to two models: (1) a dimer model,33,34 where two ADQS molecules associate with a single equilibrium constant, and (2) an isodesmic model,35−37 where multiple AQDS molecules can associate to form dimers, trimers, and any other polyaggregates, all with the same equilibrium constant (Table S1). The two models were fitted to each of three NMR peaks (noted by the different symbols) and the equilibrium constants and residuals were averaged (Figure
RESULTS
In this study, we examined the chemical and electrochemical properties of three AQDS derivatives: AQDS, AQDS-IER, and AQDS-H2SO4 (Scheme 1). AQDS-IER was chosen because prior studies had reported that increased solubility was achieved 4804
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Chemistry of Materials S6). The dimer model yielded an equilibrium constant of 1.1 × 10−3 kg/mmol with a residual of 3.5 × 10−3 ppm and the isodesmic model yielded an equilibrium constant of 1.4 × 10−3 kg/mmol with a residual of 5.0 × 10−3 ppm. NMR clearly indicates a self-association of the AQDS molecules, but as the equilibrium constant between the two models is close, NMR cannot distinguish between a dimer-only association or aggregation involving higher-ordered species. Further studies of colligative properties are needed to determine the exact solution structure. Upon heating the NMR solution to 40 °C, the equilibrium constants were unchanged for both the dimer and isodesmic model (Table S4), revealing a negligible heat of reaction for the dimer formation. The equilibrium constant at 60 °C could not be determined as the hydrogen−deuterium exchange between the AQDS protons and the deuterated water was too rapid to obtain an NMR spectrum. In addition, we examined the concentration dependence in an acidic electrolyte (0.5 M D2SO4) and found AQDS also self-associates into a dimer with an equilibrium constant of 5.5 × 10−3 kg/mmol, indicating a slightly greater driving force for dimer formation in acidic electrolyte as compared to neutral electrolyte (Table S5). Prior X-ray crystallography analysis of AQDS-like molecules supports the finding of solution-phase self-dimerization. Specifically, Gamag et al. found π stacking of the AQDS molecules with 6-fold coordination of the sodium cations, and hydrogen bonding of the water, quinone oxygen atoms, and sulfonate groups.38 In addition, our mass spectrometry (MS) analysis of the m and m+n peaks shows a 0.5 m/z unit splitting for the monoanion peak, confirming dimerization of AQDS (Figure S7 and Table S6). Surprisingly, when AQDS is dissolved in an acidic electrolyte, regardless of chemical supplier (Combi-Chem, TCI, or Santa Cruz Biotechnology), we observed CO2 gas evolution, determined via gas chromatography (Figure S8). Previous research on quinones has shown that CO2 is bound as a 1:1 quinone adduct.39−45 To determine if anthraquinone-CO2 adducts were present in our AQDS powder, we performed a high concentration 13C NMR experiment and found an additional peak present in the range for a carbonyl carbonate (Figure S9).42 In addition, acid titration of AQDS revealed an initial buffering of the solution at 9.73 followed by another plateau at 6.51. No plateaus should be observed in this range as there are no weak acidic groups on oxidized AQDS. The pKa value of the hydroxyanthraquinone species is estimated at 9.71, whereas 6.51 is close to the pKa of carbonic acid, 6.35 (Figure S10).46 Thus, we conclude that the protons in an acidic electrolyte displace the CO2 adduct present in our AQDS powder to form a hydroxyanthraquinone species and carbonate in solution (Figure S11). At pH below the carbonic acid equilibrium, 6.35, CO2 gas is evolved. Without sparging of an inert gas, this hydroxyanthraquinone species is likely quickly oxidized back to the quinone species by ambient oxygen, as in our process to synthesize AQDS-H2SO4 (Scheme 1). When AQDS is passed through the ion-exchange resin, however, the hydroxyanthraquinone is stabilized in the form of a semianthraquinone-anthraquinone dimer. The nature of the dimer and its electrochemical properties will be discussed in the following sections. AQDS-IER’s additional phenol-type proton is spectroscopic evidence of this dimer (Figure 1b). Moreover, titration of AQDS-IER with 0.1 M NaOH confirms that there is one proton per two molecules of AQDS (Figure S12). Prior research provides further support for the formation of partially reduced dimers.47−50 Spectrophotometric and potentiometric
measurements by Broadbent and Melanson for AQDS-2 in acetate buffer solution (pH = 4.65), have shown the formation of a green hydroxyanthraquinone−anthraquinone dimer.49 In our work, during the electrochemical reduction of AQDS (detailed in the next section), we also observed the solution color changing from yellow to green. Electrochemical studies by Beck47,48 and others on anthraquinone derivatives in nonaqueous and aqueous electrolytes have also suggested the formation of a hydroxyanthraquinone−anthraquinone dimer.50 Electrochemical Investigation of AQDS Derivatives. To characterize the impact of these pretreatments on the electrochemical properties of the AQDS derivatives, cyclic voltammetry (CV) was performed in an acidic electrolyte (0.5 M H2SO4, pH = 0.32), an unbuffered neutral electrolyte (0.5 M Na2SO4, pH = 5.32), and a buffered mild-alkaline electrolyte (0.5 M Na2CO3:NaHCO3, pH = 9.58) (Figure S12). Our results are in agreement with prior literature in terms of redox potential, voltammogram shape, and electrochemical reversibility.3,4,16 We observed AQDS to have a redox potential of 220 mV vs RHE in an acidic pH and 300 mV vs RHE in buffered mild-alkaline electrolyte. As shown in Figure 2, the concentration of AQDS was found to affect the kinetics of the redox event, as manifested by the
Figure 2. Peak to peak separation (ΔEpp) of the oxidative and reductive peaks during cyclic voltammetry at varying concentrations of AQDS in different electrolytes averaged over five cycles (standard deviation < 4 mV, Table S7). For all experiments, the scan rate is 20 mV/s.
concentration-dependent peak to peak separation (ΔEpp) between the oxidative and reductive peaks (Figure S13). All CVs were iR corrected to eliminate the effects of uncompensated resistance, which are expected to scale with concentration. No shifts in redox potential (averaged between the oxidation and reduction peaks) were observed with increasing concentration except in the unbuffered neutral solution. In unbuffered neutral and buffered mild-alkaline electrolytes, the ΔEpp remains constant between 1 and 50 mM indicating, on the CV time scale, the redox kinetics appear independent of concentration. In contrast, as the concentration of AQDS is increased from 1 mM to 50 mM in the acidic electrolytes, ΔEpp increases by ca. 100 mV, suggesting the rise of intermolecular reactions (Figure 2). Increasing the acid concentration decreases ΔEpp to such an extent that for a high 4805
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Chemistry of Materials
Figure 3. (a) Bulk electrolysis of 10 mM AQDS in 0.5 M Na2CO3:NaHCO3 at pH = 9.58. (b) Bulk electrolysis of 10 mM of AQDS in 0.1, 0.5, or 1.0 M H2SO4. (c) Bulk electrolysis of 10 mM of AQDS, AQDS-IER, and AQDS-H2SO4 in 0.5 M H2SO4. (d) Bulk electrolysis of 10 mM of AQDS in 0.5 M H2SO4 for 20 cycles at 1 C. All experiments were conducted with argon continuously bubbling in solution.
S15). Although we also performed electrolysis experiments at 50 mM (Figure S16), we did not perform experiments at AQDS concentrations below 10 mM as the electrochemical capacitance from the porous RVC electrode and high dielectric constant of water lead to difficulty in separating the capacitive and faradaic contributions to the observed current. In 0.5 M Na2CO3:NaHCO3, we observed reversible cycling performance with an average redox potential of −475 mV vs Ag/AgCl (315 mV vs RHE). However, as reduced AQDS is an oxygen scavenger, the charge capacity and Coulombic efficiency was extremely sensitive to oxygen content in solution.51 Accordingly, we chose a cycling rate of 2 C to minimize the amount of time the reduced product existed in solution. The consistent discharge capacity access of 1 mol e−/mol AQDS for 20 cycles at 2 C further confirms reversible oxidation and reduction of AQDS and eliminates the possibility of a deleterious side reaction causing decay of the active species (Figure 3a). The additional capacity seen in the first cycle is likely due to scavenging of the solubilized oxygen already present in the solution. Regardless of the AQDS concentration and rate, in 0.5 M Na2SO4, AQDS did not cycle. Although a reduction plateau was observed at ca. −500 mV vs Ag/AgCl
acid concentration of 1 M H2SO4 and low AQDS concentration of 1 mM, a ΔEpp of 36 mV was measured, reflecting near ideal Nernstian behavior (ΔEpp = 29.5 mV) and aligning with previous reports.3 For comparison, we also examined the effect of changing the number and position of sulfonate groups on anthraquinone (Figure S14). At a 1 mM concentration in 0.5 M H2SO4, AQDS-2 and AQDS-2,6 displayed a ΔEpp of 40 and 187 mV, respectively, which although lower, are in qualitative agreement with Yang et al., who reported 49 mV for AQDS-2 and 211 mV for AQDS-2,6 at a 1 mM concentration in 1 M H2SO4.4 As the concentration was increased to 50 mM, the ΔEpp increased to 339 mV for AQDS-2,6, whereas for AQDS-2, 50 mM exceeded the solubility limit in the acidic electrolyte. The solubility limits of AQDS-2 and AQDS-2,6 in 0.5 M Na2SO4 are below 1 mM and hence CV analysis was not performed. Energy Storage and Cycling Performance of AQDS Derivatives. Bulk electrolysis was performed on the various AQDS derivatives to probe their practical charge storage capacity, charge retention, and stability. As these materials are hydroscopic, thermogravimetric analysis, to determine mass, was performed before the electrolysis experiments (Figure 4806
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Chemistry of Materials
dimerization are likely close to that of the self-association of pristine AQDS. As a final confirmation that the 1.5 mol e−/mol AQDS capacity is not an artifact of our experimental apparatus or electrochemical testing, we performed a chemical titration experiment. First, 0.5 M of AQDS in 2 M NaOH was reduced with Na2S2O4 with a 20% molar excess to ensure complete twoelectron reduction of the AQDS, yielding a hydroxyanthraquinone form of AQDS.48,52 An aliquot of solution was then transferred using Schlenk line needle techniques to a degassed 3 M H2SO4 solution to create an overall solution concentration of 50 mM. The large excess of 3 M H2SO4 deactivates the excess Na2S2O4 and neutralizes the NaOH in the aliquot. Finally, the solution was titrated with 0.1 M FeCl3 as a chemical oxidant (Figure 4).
(300 mV vs RHE), no capacity was retained upon oxidation (Figure S16d). Upon varying the pH in the solutions, we found that if the [H+] in solution is less than the AQDS concentration, AQDS does not cycle reversibly (Figure S17), as a proton is needed to structurally stabilize the electron transfer. We then examined the effect of acid concentration on the electrolysis of AQDS (Figure 3b) to see if the number of electrons stored was dependent on the [H+] in solution. Regardless of acid concentration used, a capacity of 1.3 mol e−/ mol AQDS was accessed.3,4,6 As the acid concentration is increased, the redox potential shifts to higher potentials, consistent with the Nernstian relationship, and the overpotential is reduced, consistent with an increase in overall solution conductivity (Figure S18). The slight variations in charge/discharge capacities in Figure 3b are likely associated with either AQDS redox shuttling between the working and counter electrodes, or reaction with oxygen generated at the counter electrode. The fifth bulk electrolysis cycle for each of the AQDS derivatives in 0.5 M H2SO4 (Figure 3c) indicated that AQDSH2SO4 and AQDS have the same capacity, whereas AQDS-IER has a 17% greater capacity (ca. 1.5 mol e−/mol AQDS). Moreover, the redox potential, reduction, and oxidation plateaus overlap for all of the derivatives, signaling the same active species is operating in solution, consistent with the CV analysis (Figure S13). We attribute this difference in capacity between 1.3 and 1.5 mol e−/mol AQDS to differences in the fraction of AQDS that is bound to CO2 and the additional cations that must be present for charge neutrality. For example, if one CO2 molecule and one sodium cation are bound to AQDS, the increase in molar mass extends the capacity from 1.3 to 1.5 mol e−/mol AQDS. For ease of comparison, we assumed in Figure 3c each derivative to have the same molar mass of 412.30 g/mol. For AQDS and AQDS-H2SO4, gas bubbles are observed indicating bound CO2, whereas for AQDS-IER, no bubbles are observed; consequently, the number of “active” moles varies from 1.3 to 1.5 mol e−/mol AQDS for these solutions. Regardless of the electrochemical driving force applied to the system, we could not completely oxidize AQDS beyond 1.5 mol e−/mol AQDS warranting additional studies on how to achieve the theoretical 2 mol e−/ mol AQDS. As shown in Figure 3d, in acidic conditions AQDS can reversibly cycle at 1 C for 20 cycles and access a 1.3 mol e−/ mol AQDS. No decay in the Coulombic efficiency or capacity was detected, which supports observations of stable cycling at the cell level.3,6 Importantly, the same capacity per unit mass and cycle stability was found across all concentrations tested (10−50 mM) as well as when the charge/discharge rate was decreased to C/3, and when a different cell geometry was used, confirming this result is independent of electrolysis conditions or experimental apparatus (Figure S16). In addition, we performed a single reductive electrolysis step on 10 mM AQDS in 0.5 M D2SO4 and analyzed the reduced solution with 1H NMR (Figure S19). No additional peaks were observed and the chemical shifts of reduced species were close to the chemical shifts of the neutral species, suggesting that although the constituents of the dimer after reduction is different, transforming from semianthraquinone-anthraquinone (neutral form) to hydroxyanthraquinone species (reduced form), the chemical bonding and equilibrium constant for
Figure 4. Chemical titration of 50 mM of reduced AQDS in 3 M H2SO4 with 0.1 M FeCl3.
We observed a plateau with a total capacity of 1.5 mol e−/ mol AQDS at the redox potential for AQDS in acid, 17 mV vs Ag/AgCl (17 mV vs RHE), indicating chemical oxidation of AQDS (Figure 4). Overall, this experiment is a chemical analogy to reducing electrochemically the AQDS-containing solution 20% beyond its theoretical capacity, and then measuring the capacity retained upon oxidation. The same capacity of 1.5 mol e−/mol AQDS was obtained for KMnO4, which was calibrated against a BAM certified sodium oxalate (Figure S20). Chemical titration confirms the electrochemical results (Figure 3) that only 1.5 mol e−/mol AQDS can be accessed upon oxidizing the reduced form of AQDS (Figure 4).
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DISCUSSION When the AQDS-CO2 adduct is placed in an aqueous solution, the carboxyl group is quickly hydrolyzed to a hydroxyl group (Scheme 2). If the pH of the aqueous solution is below 6.35, the carbonate anion will be expelled as CO2 gas. Two semianthraquinone molecules will disproportionate to form a fully oxidized AQDS and a hydroxyanthraquinone species and subsequently dimerize to a hydroxyanthraquinone−anthraquinone dimer. If the solution is exposed to ambient oxygen, this dimer will fully oxidize to the anthraquinone AQDS form, whereas if kept under an inert atmosphere in a buffered mildalkaline electrolyte, the dimer is stabilized. In an acid 4807
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Scheme 2. AQDS Monomer with a CO2 Adduct Will Be Hydrolyzed and Dimerize with a Varying Level of Reduction Depending on the Electrolytea
a
AQDS in the buffered mild-alkaline and acid electrolyte is not exposed to the atmosphere due to consistent sparging with argon.
environment, the hydroxyanthraquinone−anthraquinone dimer is partially oxidized to a semianthraquinone−anthraquinone dimer. Surprisingly, even upon sparging with pure oxygen the dimer is not fully oxidized and remains a semianthraquinoneanthraquinone dimer suggesting conventional methods cannot be used to completely oxidize the AQDS dimer. Our results demonstrate two electrochemical regimes of operation exist for AQDS: buffered mild-alkaline electrolytes and acidic electrolytes. In buffered mild-alkaline electrolytes, AQDS displays a redox potential of 304 mV vs RHE, a maximum capacity of 1.0 mol e−/mol AQDS, and no concentration dependence, at least in the range tested. In acidic electrolytes, AQDS displays a redox potential of 220 mV vs RHE, a maximum capacity of 1.5 mol e−/mol AQDS, and a concentration-dependent electrochemical behavior. This bimodal electrochemical behavior stems from the difference in the solution-phase chemical structure of AQDS in these two environments. In neutral electrolytes, AQDS cycles between the dimer hydroxyanthraquinone−anthraquinone state (Scheme 3a) with two available sites and the hydroxyanthraquinone, as only one electron per molecule of AQDS is observed. Moreover, AQDS does not show concentration dependent interactions with other species in solution and is a relatively stable unit, warranting additional experimental investigation. In acidic electrolytes, electrolysis experiments conclude this species exists as a dimer of AQDS and a semianthraquinone species, formed by intermolecular
Scheme 3. (a) Hydroxyanthraquinone−Anthraquinone Dimer in Neutral to Basic pH Reversibly Stores Two Electrons To Give an Overall Energy Density of One Electron Per Molecule of AQDS; (b) Semianthraquinone− Anthraquinone Dimer in Acidic pH Reversibly Stores Three Electrons To Give an Overall Capacity of 1.5 Electrons Per Molecule of AQDS
hydrogen bonding (Scheme 3b). The partially reduced dimer has three available sites, hence 1.5 mol e−/mol AQDS. This intrinsic charge storage capacity is also confirmed with chemical titration (Figure 4). 4808
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It is worth noting that AQDS stability is another factor to consider when selecting electrolyte environment as RFBs are expected to operate for multiple years over a range of electrochemical potentials and current regimes. Although a higher charge storage capacity is accessed in an acidic solution, prior work by Beck has shown that when AQDS is fully reduced to hydroxyanthraquinone, this species undergoes an irreversible acid-catalyzed disproportionation to anthraquinone and anthrone with a second-order rate constant of 3 dm3 mol−1 h−1 (1 M H2SO4) to 6.5 dm3 mol−1 h−1 (5 M H2SO4).48 When operated in a lower capacity, neutral pH solution, however, this decay mechanism should be mitigated although further stability studies are warranted.
AUTHOR INFORMATION
Corresponding Author
*Professor Fikile R. Brushett:
[email protected]. ORCID
Thomas J. Carney: 0000-0002-9250-1659 Fikile R. Brushett: 0000-0002-7361-6637 Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This work was funded by the Joint Center for Energy Storage Research (JCESR) (DE-AC02-06CH11357), an Energy Innovation Hub funded by the United States Department of Energy, Office of Science, Basic Energy Sciences. This research was conducted with Government support under and awarded by DoD, Air Force Office of Scientific Research, National Defense Science and Engineering Graduate (NDSEG) Fellowship, 32 CFR 168a. We acknowledge Jarrod D. Milshtein, Jeffrey A. Kowalski, John L. Barton, Dr. Liang Su of MIT and Dr. Anthony Burrell of the National Renewable Energy Laboratory for fruitful discussions.
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CONCLUSIONS The electrochemical reversibility and number of electrons transferred is divided into two regimes for AQDS: an acidic pH regime, where AQDS associates into a semianthraquinone− anthraquinone dimer, and a buffered mild-alkaline pH regime, where AQDS associates into a hydroxyanthraquinone− anthraquinone dimer. The partial reduction of the dimer arises from CO2 bound to the AQDS that is displaced in aqueous solution, leading to a semianthraquinone species. In the acidic regime, the semianthraquinone species is partially oxidized to a semianthraquinone−anthraquinone dimer, whereas in the buffered mild-alkaline regime, the species is stabilized into a hydroxyanthraquinone−anthraquinone dimer. In the acid regime, the redox behavior is highly dependent on concentration and the number of electrons stored is 1.5 per molecule of AQDS. For the buffered mild-alkaline regime, the redox behavior is independent of concentration and the number of electrons stored is one per molecule of AQDS. AQDS does not cycle in an unbuffered neutral pH. We found no pathway to store two electrons per molecule of AQDS, even upon employing chemical oxidants during a chemical titration. Two electrons cannot be reversibly stored by AQDS as it will undergo intermolecular dimerization and cycle between the dimer and hydroxyanthraquinone species. Determining the optimum operating conditions for an aqueous organic RFB comprising AQDS is a trade-off between stability and capacity. Although in an acidic solution a higher capacity is accessed, AQDS is prone to irreversible decay, whereas in a lower capacity, neutral pH solution, AQDS’s decay mechanism is mitigated. As the deviation from the ideal storage capacity of two electrons per molecule of AQDS stems from the self-dimerization, which prevents the full oxidation of the molecule, future research should aim to prevent this selfdimerization. Specifically adding functional groups that sterically hinder the dimerization while not greatly affecting the solubility and the redox behavior could provide a pathway to storing two electrons per molecule of AQDS.
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Article
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ABBREVIATIONS AQDS, 9,10-Anthraquinone-2,7-disulfonic acid; AQDS-IER, AQDS subject to ion-exchange resin; AQDS-H2SO4, AQDS subject to H2SO4 treatment
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ASSOCIATED CONTENT
S Supporting Information *
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.chemmater.7b00616. Optical photographs, extended NMR spectrum, mass spectrometry analysis, gas chromatography, pH titration, cyclic voltammetry, bulk electrolysis, thermogravimetric analysis, and chemical titration (PDF) 4809
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