Question #44089

To study redox precipitation constraints, solid silver hypochlorite, $\ce{AgClO(s)}$, is added to an aqueous solution. The hypochlorite ion undergoes disproportionation to chlorate and chloride ions, superseding simple solubility equilibrium. \ Given the following standard reduction potentials and solubility product constants at $298\text{ K}$: \ $E^\circ_{\ce{ClO^-/Cl^-}} = 0.89\text{ V}$ \ $E^\circ_{\ce{ClO_3^-/ClO^-}} = 0.62\text{ V}$ \ $K_{sp}(\ce{AgClO}) = 10^{-4}$ \ $K_{sp}(\ce{AgCl}) = 10^{-10}$ \ Assume $\frac{2.303RT}{F} = 0.06\text{ V}$. \ Determine the equilibrium constant ($K_{eq}$) for the following overall transformation: \ $3\ce{AgClO(s)} \rightleftharpoons 2\ce{AgCl(s)} + \ce{Ag+(aq)} + \ce{ClO_3^{-}(aq)}$