In the given cell reaction, zinc ($\text{Zn}$) is oxidized to $\text{Zn}^{2+}$ and acts as the anode, while silver oxide ($\text{Ag}_2\text{O}$) is reduced to silver ($\text{Ag}$) and acts as the cathode. The standard reduction potentials are given as: $E^\circ_{\text{cathode}} (\text{Ag}_2\text{O}/\text{Ag}) = 0.34 \text{ V}$ $E^\circ_{\text{anode}} (\text{Zn}^{2+}/\text{Zn}) = -0.76 \text{ V}$

The standard cell potential ($E^\circ_{\text{cell}}$) is calculated using the formula: $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$ $E^\circ_{\text{cell}} = 0.34 \text{ V} - (-0.76 \text{ V})$ $E^\circ_{\text{cell}} = 0.34 \text{ V} + 0.76 \text{ V} = 1.10 \text{ V}$