Question #44738

An aqueous solution of $1.0\text{ M } \ce{NaCl}$, buffered at $\text{pH} = 7.0$, is electrolyzed at $298\text{ K}$ using smooth Platinum electrodes. A constant current of $10.0\text{ A}$ is passed for $965\text{ s}$.

Given standard reduction potentials at $298\text{ K}$: $E^\circ_{\ce{Cl2(g)}/\ce{Cl^- (aq)}} = +1.36\text{ V}$ $E^\circ_{\ce{O2(g)}, \ce{H^+ (aq)}/\ce{H2O(l)}} = +1.23\text{ V}$ $E^\circ_{\ce{H^+ (aq)}/\ce{H2(g)}} = 0.00\text{ V}$

Kinetic overpotentials ($\eta$) on Pt under these conditions: $\eta_{\ce{O2}} = 0.65\text{ V}$ $\eta_{\ce{Cl2}} = 0.05\text{ V}$ $\eta_{\ce{H2}} = 0.00\text{ V}$

Assume $\frac{2.303 RT}{F} = 0.06\text{ V}$ and $1\text{ F} = 96500\text{ C mol}^{-1}$. All gases are at $1\text{ bar}$.

Calculate the total minimum electrical energy required to drive the cell and produce the major anodic product. (Assume 100% current efficiency).