A student electroplates pure copper onto an electrode using an aqueous solution of \(\text{CuSO}_4\). A constant current of \(9.65\text{ A}\) is passed through the electrolytic cell for \(2{,}000\text{ s}\).

\[ \text{Cu}^{2+}\text{(aq)} + 2\,e^- \rightarrow \text{Cu(s)} \]

Assuming that all current is consumed in the reduction of \(\text{Cu}^{2+}\text{(aq)}\) ions, what mass of \(\text{Cu(s)}\) is deposited on the electrode? (Faraday's constant \(F = 96{,}500\text{ C/mol } e^-\); molar mass of \(\text{Cu} = 63.55\text{ g/mol}\))