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Positive electrode half-cell reaction 2ClO2- ↔ 2ClO2 + 2e- (Eq. 2) Full cell reaction 2H2O + 2ClO2- ↔ H2 + 2OH- + 2ClO2 (Eq. 3) After this formation step, the cell is in the fully charged state with Zn and ClO2 present at negative and positive electrodes respectively, and may be operated reversibly, with the subsequent half-cell and full- cell reactions being: Negative electrode half-cell reaction Positive electrode half-cell reaction Full cell reaction Zn + 2OH- ↔ Zn(OH)2 + 2e- 2ClO2 + 2e- ↔ 2ClO2- Zn + 2OH- + 2ClO2 ↔ Zn(OH)2 + 2ClO2- (Eq. 4) (Eq. 5) (Eq. 6) Assuming a balanced cell with equal capacity at the negative and positive electrodes, an alkaline electrolyte containing NaClO2 as the sole source of working chlorite ions, and that the ClO2 formed upon charge is present as a pure liquid, the theoretical energy density may be calculated. This energy density is primarily limited by the solubility of NaClO2 in the electrolyte, which we independently determined to be 2 M and 3 M at 0°C and 20°C, respectively. The corresponding energy density values are 77 Wh/L and 106 Wh/L, and the chemical cost of stored energy (cost of Zn and electrolyte components divided by the stored energy) is 7.3 US$/kWh. This chemical cost is significantly lower than for Li-ion, Ni-Cd and NiMH batteries (30-80 US$/kWh) as well as vanadium redox flow batteries (100 US$/kWh), and is comparable to that for primary Zn/MnO2 and high temperature rechargeable Na/NiCl2 batteries.[14] Initial full-cell experiments conducted in three-electrode beaker-type cells led to a custom-built flat cell design shown in Figure 2a, used to obtain the results below. We recognized two likely parasitic reactions, and adapted our cell design to reduce their occurrence. One reaction is that between Zn metal and ClO2-, forming solid Zn(OH)2 and water.[15] Our prototype cell uses a zinc sheet negative electrode and vitreous carbon foam positive electrode, separated by a Nafion 117 membrane which was pre-soaked in NaOH solution to provide Na+ ion conductivity while significantly blocking the crossover of chlorite ions, thereby mitigating this side reaction. A second possible side reaction is the disproportionation reaction ClO2 + H2O ↔ ClO2- + ClO3- + 2H+.[16] To minimize this reaction, one possible mitigation is to allow the ClO2 to exsolve as a separate liquid phase, taking advantage of its low solubility in water, thereby reducing contact. A second approach to separating ClO2 from water is to provide a separate, immiscible, organic phase in which the ClO2 is soluble, as an accumulator phase. This approach decreases energy density, but has the advantage of lowering the vapor pressure of ClO2. Several organic liquids having significant solubility for ClO2 are known [17]; here we selected a synthetic saturated hydrocarbon solvent with a similar range of carbon number to that of diesel fuel, a ClO2 solubility of ~4 g/L, and which segregates as a surface layer on the aqueous electrolyte due to its lower density (Figure 2a). Half-cell experiments (Supplementary Information, Figure S2) conducted using the vitreous carbon positive electrode with a nickel mesh negative electrode (instead of zinc as in the full cell) showed the efficacy of this approach. As the amount of organic phase was increased, so did the reversible capacity. For the proportions of each material used in the full cell (Figure 2a), two-thirds of the ClO2 produced will reside in the accumulator phase at equilibrium, and is separated from the aqueous electrolyte in which disproportionation may occur. The distribution of ClO2 between the accumulator phase and the aqueous electrolyte could be further varied through the choice of organic phase and the relative amounts of these two phases. Results obtained from a representative Zn-ClO2 full cell tested at 0°C appear in Figure 2b, plotted as cell voltage vs capacity for a capacity-limited cycling regimen wherein galvanostatic charging at a current 5PDF Image | Reversible Chlorite Chlorine Dioxide Anion Redox Storage
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