The overall cell reaction is given as: $\text{Fe} + 2\text{Fe}^{3+} \rightarrow 3\text{Fe}^{2+}$.

This reaction can be split into two half-reactions:

1. Oxidation at Anode: $\text{Fe} \rightarrow \text{Fe}^{2+} + 2e^-$ (Standard reduction potential, $E^\circ_{\text{anode}} = E^\circ_{\text{Fe}^{2+}/\text{Fe}} = -0.441 \text{ V}$)
2. Reduction at Cathode: $2\text{Fe}^{3+} + 2e^- \rightarrow 2\text{Fe}^{2+}$ (Standard reduction potential, $E^\circ_{\text{cathode}} = E^\circ_{\text{Fe}^{3+}/\text{Fe}^{2+}} = 0.771 \text{ V}$)

The standard emf of the cell ($E^\circ_{\text{cell}}$) is calculated using the formula: $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$ $E^\circ_{\text{cell}} = 0.771 \text{ V} - (-0.441 \text{ V})$ $E^\circ_{\text{cell}} = 0.771 \text{ V} + 0.441 \text{ V} = 1.212 \text{ V}$

Note: Standard electrode potential is an intensive property . Therefore, multiplying the stoichiometric coefficients of the reduction half-reaction by 2 to balance the electrons does not change the value of its $E^\circ$.