Title : Alkaline Quinone Flow Battery

Storage of photovoltaic and wind electricity in batteries could solve the mismatch problem between the intermittent supply of these renewable resources and variable demand. Flow batteries permit more economical long-duration discharge than solid-electrode batteries by using liquid electrolytes stored outside of the battery. We report an alkaline flow battery based on redox-active organic molecules that are composed entirely of earth-abundant elements and are non-toxic, non-flammable, and safe for use in residential and commercial environments. The battery operates efficiently with high power density near room temperature. These results demonstrate the stability and performance of redox-active organic molecules in alkaline flow batteries, potentially enabling cost-effective stationary storage of renewable energy. Main Text: The cost of photovoltaic (PV) and wind electricity has dropped so much that one of the largest barriers to getting the vast majority of our electricity from these renewable sources is their intermittency (1–3). Batteries provide a means to store electrical energy; however, traditional, enclosed batteries maintain discharge at peak power for far too short a duration to adequately regulate wind or solar power output (1, 2). In contrast, flow batteries can independently scale the power and energy components of the system by storing the electro-active species outside the battery container itself (3–5). In a flow battery, the power is generated in a device resembling a fuel cell, which contains electrodes separated by an ion-permeable membrane. Liquid solutions of redox-active species are pumped into the cell where they can be charged and discharged, before being returned to storage in an external storage tank. Scaling the amount of energy to be stored thus involves simply making larger tanks (Fig 1A). Existing flow batteries are based on metal ions in acidic solution but there are challenges with corrosivity, hydrogen evolution, kinetics, materials cost and abundance, and efficiency that thus far have prevented large-scale commercialization. The use of anthraquinones in an acidic aqueous flow battery can dramatically reduce battery costs (6, 7); however, the use of bromine in the other half of the system precludes deployment in residential communities due to toxicity concerns. We demonstrate that quinone-based flow batteries can be adapted to alkaline solutions, where hydroxylated anthraquinones are highly soluble and bromine can be replaced with the non-toxic ferricyanide ion (8, 9) — a food additive (10). Functionalization of 9,10-anthraquinone (AQ) with electron-donating groups such as OH has been shown to lower the reduction potential and expand the battery voltage (6). In alkaline solution, these OH groups are deprotonated to provide solubility and greater electron donation capability, which results in an increase in the open circuit voltage of 47% over the previously reported system. Because functionalization away from the ketone group provides molecules with the highest solubility (11, 12), we initially targeted commercially available 2,6-dihydroxyanthraquinone (2,6-DHAQ), which we find exhibits a room temperature solubility > 0.6 M in 1 M KOH. This system can achieve power densities in excess of 0.45 W cm 2 at room temperature and 0.7 W cm 2 at 45 °C. The use of alkaline electrolyte exploits pH as a parameter to shift the thermodynamic potentials of proton-dependent reactions to more negative values. In acid solutions, AQ undergoes a two-electron two-proton reduction at a single potential, which shifts to more negative values as the pH increases (6). When the pH exceeds 12 the reduction potential of 2,6DHAQ becomes pH-independent because the reduced species is generated in its fullydeprotonated form (Fig. S1). In contrast with the pH-dependent electrochemical behavior of quinones (negative terminal), the ferro/ferricyanide redox couple (positive terminal) has a pHindependent redox potential. This contrasting pH-dependence can be exploited through the development of low reduction potential quinones at high pH. The cyclic voltammograms (CVs) of 2,6-DHAQ and ferro/ferricyanide predict an equilibrium cell potential of 1.2 V upon combination of these two half-reactions (Fig. 1B). A quantitative analysis of the CV of 2,6DHAQ at pH 14 (Fig. S2) revealed redox behavior consistent with two one-electron reductions at potentials separated by only 0.06 V with rapid kinetic rate similar to quinones in acid (6). This behavior raises interesting questions about the relationship between quinone redox and hydrogen bonding (13). Cell testing was performed at 20 °C using solutions of 0.5 M 2,6-DHAQ dipotassium salt and 0.4 M K4Fe(CN)6 both dissolved in 1 M KOH. These solutions were pumped through a flow cell constructed from graphite flow plates and carbon paper electrodes, which were separated by a Nafion membrane. A charging current of 0.1 A cm was applied to charge the cell, and polarization curves were measured at 10%, 50%, and 100% states of quinone charge (SOC). The open-circuit voltage (OCV) is 1.2 V at 50% SOC; its dependence on SOC is shown in Fig. 2A. The polarization curves (Fig. 2B) show no sign of redox kinetic limitations and exhibit a peak galvanic power density exceeding 0.4 W cm . The cell was cycled at a constant current density of ±0.1 A cm for 100 cycles (Fig. 3A). The current efficiency exceeded 99%, with a stable round-trip energy efficiency of 84%. A 0.1% loss in capacity per cycle was observed during cycling, which appears to be a continuous loss of electrolyte over the 100 cycles. Three possible loss mechanisms were explored: chemical decomposition, electrolyte crossover through the membrane, and leakage from the pumping system. Chemical and electrochemical stability studies showed that the negative electrolyte is stable. 10 mM of 2,6-DHAQ was heated in 5 M KOH solution at 100 °C for 30 days and was characterized by proton NMR. Cycled negative electrolyte was also collected and characterized by the same method; both studies showed no degradation product at the sensitivity level of 1% (Fig S3). Membrane crossover contamination has been a common challenge in acid-based redox flow batteries where most electro-active molecules are either neutral or positive and tend to migrate through proton-conductive membranes (5). In this alkaline system, however, all the electro-active molecules remain negatively charged in all charge states, leading to a dramatic decrease in the degree of crossover during cell cycling. Cyclic voltammetry of the ferro/ferricyanide electrolyte collected at the end of cycling showed no evidence of the presence of 2,6-DHAQ. This observation places an upper limit on crossover of 0.8% of the DHAQ, implying a crossover current density less than 2.5 μA cm (Fig. S4). Finally, hydraulic leakage was investigated because an apparent but unquantifiable small decrease in fluid levels was observed in the reservoirs. After cell cycling, the cell was washed with DI water until no coloration of eluent could be observed. The cell was then dissembled; coloration was found on the gaskets indicating the likely site of electrolyte leakage (Fig. S5). This source of capacity loss—the equivalent to roughly 8 drops in our system—is expected to become negligible as system size is scaled up. By increasing the temperature to 45 °C, the peak galvanic power density increases from 0.45 W cm to approximately 0.7 W cm (Fig. 2C), as the cell area-specific resistance (ASR) decreases from about 0.878 to 0.560 , estimated from the linear parts of the polarization curves in Fig. 2. The majority of this ASR decrease comes from a change in the high frequency area-specific resistance (rhf) measured by electrochemical impedance spectroscopy (Fig. S6). In both cases, the rhf contributes more than 70% of the ASR and is indeed the limiting factor to the cell current and power outputs. The rhf is dominated by the resistance of the membrane, which is an order of magnitude higher than the resistance of the same membrane in a pH 0 acid solution (14). The sluggish kinetics of the hydrogen evolution reaction in alkaline solution on carbon electrodes results in a larger practical stability window in base rather than in acid (Fig. S7). Consequently, quinones with substantially more negative reduction potentials are feasible as negative electrolyte materials. Preliminary investigations into the synthesis of different hydroxysubstituted anthraquinones suggest further increases in cell potential are possible. Selfcondensation reactions of substituted benzene yield 2,3,6,7-tetrahydroxy-AQ (THAQ) (15) and 1,5-dimethyl-2,6-DHAQ (15-DMAQ) (Fig. S8 and S9). The CV’s of these species in 1 M KOH suggest cell potentials vs. ferri/ferrocyanide approaching 1.35 V (Fig. 4A and B), which exceeds that of many aqueous rechargeable batteries (Fig. 4C). The cyanide ions in both ferroand ferricyanide are bound too tightly to be released under most conditions. Consequently, it is non-toxic in both oxidized and reduced forms, and is even permitted for use as a food additive (10). The use of ferrocyanide offers significant advantages over bromine because it is non-volatile and non-corrosive, allowing simpler and less expensive materials of construction. In addition, these triand tetra-anionic organometallic molecules exhibit low crossover rates through cation-exchange membranes. The results reported herein highlight the ability of hydroxy-substituted anthraquinone and ferrocyanide to function as stable flow battery electrolytes in alkaline solution. The use of organic and organometallic coordination complexes in base, rather than aqueous metal ions in acid, resolves serious cost, corrosion, and safety concerns of previous flow battery chemistries. Alkaline flow batteries can compensate for higher membrane resistance with higher voltage, leading to performance similar to their acidic counterparts. In addition, quinone-ferrocyanide alkaline chemistry avoids the membrane crossover, corrosivity, toxicity and regulations associated with bromine. This reduced corrosivity can lead to a substantially lower materials cost because many components can be made of inexpensive polyolefin or PVC plastics. References and Notes: 1. B. Dunn, H. Kamath, J.-M. Tarascon, Electrical Energy Storage for the Grid: A Battery of Choices. Science. 334, 928–935 (2011). 2. Z. Yang et al., Electrochemical Energy Storage for Green Grid. Chem. Rev. 111, 3577– 3613 (2011). 3. T. Nguyen, R. F. Savinell, Flow Batteries. Electrochem. Soc. Interface. 19, 54–56 (2010). 4. D. Biello, Solar Wars. Sci. Am. 311, 66–71 (2014). 5. M. Skyllas-Kazacos, M. H. Chakrabarti, S. A. Hajimolana, F. S. Mjalli, M. Saleem, Progress in Flow Battery Research and Development. J. Electrochem. Soc. 158, R55–R79 (2011). 6. B. Huskinson et al., A metal-free organic–inorganic aqueous flow battery. Nature. 505, 195–198 (2014). 7. R. M. Darling, K. G. Gallagher, J. A. Kowalski, S. Ha, F. R. Brushett, Pathways to low-cost electrochemical energy storage: a comparison of aqueous and nonaqueous flow batteries. Energy Env. Sci. 7, 3459–3477 (2014). 8. G. G. I. Joseph, A. J. Gotcher, G. Sikha, G. J. Wilson, High performance flow battery (2011), (available at http://www.google.com/patents/US20110244277). 9. J. R. Goldstein, Novel flow battery and usage thereof (Google Patents, 2014; http://www.google.com/patents/US20150048777). 10. “Seventeenth Report of the Joint FAO/WHO Expert Committee on Food Additives. Report No. 539,” Wld Hlth Org. techn. Rep. Ser. (539, World Health Organization, 1974). 11. H. Pal, T. Mukherjee, J. P. Mittal, Pulse radiolytic one-electron reduction of 2-hydroxy-and 2, 6-dihydroxy-9, 10-anthraquinones. J Chem Soc Faraday Trans. 90, 711–716 (1994). 12. S. Er, C. Suh, M. P. Marshak, A. Aspuru-Guzik, Computational design of molecules for an all-quinone redox flow battery. Chem. Sci. (2014), doi:10.1039/C4SC03030C. 13. M. Quan, D. Sanchez, M. F. Wasylkiw, D. K. Smith, Voltammetry of Quinones in Unbuffered Aqueous So Bonding in the Aqueous Electrochemistry of Quinones. J. Am. Chem. Soc. 129, 12847– 12856 (2007). 14. Q. Chen, M. R. Gerhardt, L. Hartle, M. J. Aziz, A Quinone-Bromide Flow Battery with 1 W/cm2 Power Density. J. Electrochem. Soc. 163, A5010–A5013 (2016). 15. T. S. Balaban, A. Eichhöfer, M. J. Krische, J.-M. Lehn, Hierarchic Supramolecular Interactions within Assemblies in Solution and in the Crystal of 2, 3, 6, 7-Tetrasubstituted 5, 5 -(Anthracene-9, 10-diyl) bis [pyrimidin-2-amines]. Helv. Chim. Acta. 89, 333–351 (2006). 16. D. Linden, T. B. Reddy, Handbook of batteries (McGraw-Hill, 2002). 17. A. J. Bard, L. R. Faulkner, Electrochemical Methods: Fundamentals and Applications (Wiley, 2000) 18. L. B. Hitchcock, J. S. McIlhenny, Viscosity and Density of Pure Alkaline Solutions and Their Mixtures. Ind. Eng. Chem. 27, 461–466 (1935). 19. K. B. Oldham, J. C. Myland, Modelling cyclic voltammetry without digital simulation. Electrochimica Acta. 56, 10612–10625 (2011). 20. Q. H. Liu et al., High Performance Vanadium Redox Flow Batteries with Optimized Electrode Configuration and Membrane Selection. J. Electrochem. Soc. 159, A1246–A1252 (2012). 21. A. J. Esswein, J. Goeltz, D. Amadeo, High solubility iron hexacyanides (2014), US20140051003 22. H. Behre, F. Mueller-Hauck, J. Scherer, G. Schroeder, Method for producing 3-hydroxy-2methylbenzoic acid (2004), WO2003080542A3. Acknowledgments: This work was funded by the U.S. DOE ARPA-E award # DE-AR0000348. Methods, along with any additional Extended Data display items and Source Data, are available in the online version of the paper; references unique to these sections appear only in the online paper. M.P.M., K.L., R.G.G. and M.J.A. formulated the project. K.L., S.B.K., and D.H. synthesized, analyzed and purified the compounds. K.L., L.T. and S.B.K. collected and analysed the NMR and MS data. K.L. A.V. L.E. measured solubility. K.L., Q.C. L.E. and M.R.G. collected and analysed the electrochemical data. K.L., Q.C., M.R.G., M.P.M., R.G.G. and M.J.A. wrote the paper, and all authors contributed to revising the paper. Supplementary Materials: Materials and Methods Supplementary Text Fig. S1 to S9 References (17-22) Fig. 1 Cyclic Voltammetry of Electrolyte and Cell Schematic. (A) Cyclic voltammogram of 2 mM 2,6-DHAQ (dark cyan curve) and ferrocyanide (gold curve) scanned at 100 mV/s on glassy carbon electrode; arrows indicate scan direction. Dotted line represents cyclic voltammogram of 1 M KOH background scanned at 100 mV/s on graphite foil electrode. (B) Schematic of cell in charge mode. Cartoon on top of the cell represents sources of electrical energy from wind and solar. Curved arrows indicate direction of electron flow and white arrows indicate electrolyte solution flow. Blue arrow indicates migration of cations across the membrane. Essential components of electrochemical cells are labeled with color-coded lines and text. The molecular structure of oxidized and reduced species are shown on corresponding reservoirs. Fig. 2 Cell Performance. (A) Cell OCP vs. SOC. All potentials were taken when cell voltage stabilized to within ± 1 mV. 100% SOC was reached by potentiostatic holding at 1.5 V until the current decreases to below 20 mA/cm. (B) and (C), Cell voltage & power density vs. current density at 20 C and 45 C, respectively, at 10%, 50%, and ~100% SOC. Electrolyte composition: At 20 °C, 0.5 M 2,6-DHAQ and 0.4 M ferrocyanide were used in negative electrolyte and positive electrolyte, respectively. At 45 °C, both concentrations were doubled. In both cases, potassium hydroxide content started at 1 M for both sides in the fully discharged state. Fig. 3 Cell Cycling Performance. (A) Representative voltage vs. time curves during 100 charge-discharge cycles at 0.1 A/cm, recorded between the 10 and 19 cycles. (B) Capacity retention, current efficiency and energy efficiency values of 100 cycles. Normalized capacity is evaluated based on the capacity of the first charge and discharge cycle. Fig. 4 Molecular Structure and Cyclic Voltammetry of 2,6-DHAQ Derivatives. (A) and (B) Molecular structures and cyclic voltammogram of 2,3,6,7-THAQ (pink) and 1,5-DMAQ (olive), respectively, plotted along with ferrocyanide (orange curve) scanned at 100 mV/s on glassy carbon electrode. Both 2,6-DHAQ derivatives/ferrocyanide couples showed higher equilibrium potential than 2,6-DHAQ/ferrocyanide. (C), Selected aqueous secondary batteries showing voltage and flow status. Literature data from (3) and (16).