A student constructs a concentration cell at \(298\text{ K}\) using two identical \(\text{Cu(s)}\) electrodes. Compartment 1 contains \(100\text{ mL}\) of \(0.010\text{ M }\text{CuSO}_4\text{(aq)}\) and serves as the anode, while compartment 2 contains \(100\text{ mL}\) of \(1.0\text{ M }\text{CuSO}_4\text{(aq)}\) and serves as the cathode. The overall spontaneous cell reaction is represented below.

\[\text{Cu}^{2+}\text{(aq, compartment 2)} \rightarrow \text{Cu}^{2+}\text{(aq, compartment 1)}\quad E_{\text{cell}} > 0\text{ V}\]

The student then adds excess concentrated \(\text{NH}_3\text{(aq)}\) to compartment 2, causing the solution to turn dark blue due to the formation of the tetraamminecopper(II) complex ion according to the following equilibrium:

\[\text{Cu}^{2+}\text{(aq)} + 4\text{NH}_3\text{(aq)} \rightleftharpoons [\text{Cu}(\text{NH}_3)_4]^{2+}\text{(aq)} \quad K_f \gg 1\]

Which of the following best predicts the effect of adding excess \(\text{NH}_3\text{(aq)}\) to compartment 2 on the cell potential, \(E_{\text{cell}}\), and provides the correct justification?