Question #44751

A $1.0 \text{ L}$ aqueous solution of $\ce{H2O2}$ (initial concentration $1.0 \text{ M}$) is subjected to electrolysis at $298 \text{ K}$ using inert electrodes and a constant current of $193 \text{ A}$. Simultaneously, in the presence of a homogeneous catalyst, $\ce{H2O2}$ undergoes a pseudo-first-order disproportionation into $\ce{H2O(l)}$ and $\ce{O2(g)}$ with a rate constant $k = 10^{-3} \text{ s}^{-1}$. The anodic electrolysis reaction strictly oxidizes $\ce{H2O2}$ to $\ce{O2(g)}$ and $\ce{H+(aq)}$. Assuming no volume change and $100\%$ current efficiency, what is the total moles of $\ce{O2(g)}$ evolved from the solution by the time $\ce{H2O2}$ is completely depleted? (Given: $F = 96500 \text{ C mol}^{-1}$)