Define the term standard cell potential, $E^\circ_{\text{cell}}$.
An electrochemical cell is arranged to determine $E^\circ_{\text{cell}}$ for a cell made from an $\text{Fe}^{3+}/\text{Fe}^{2+}$ half-cell and a $\text{Cl}_2/\text{Cl}^-$ half-cell. Draw a labelled diagram of this electrochemical cell. Include every necessary substance. It is not necessary to state the conditions used.
The cell reaction for the electrochemical cell in (b) is given below. $\text{Cl}_2 + 2\text{Fe}^{2+} \rightarrow 2\text{Fe}^{3+} + 2\text{Cl}^-$ $E^\circ_{\text{cell}} = +0.59\,\text{V}$ Calculate $\Delta G^\circ$, in $\text{kJ mol}^{-1}$, for this cell reaction.
A second experiment is carried out with the same electrochemical cell. In this experiment, the $\text{Fe}^{2+}$ concentration is $0.15\,\text{mol dm}^{-3}$. All other concentrations stay at their standard values. The Nernst equation is given below. $E = E^\circ + (0.059/z) \log\left(\frac{\text{oxidised species}}{\text{reduced species}}\right)$ Use the Nernst equation to calculate the electrode potential, $E$, for the $\text{Fe}^{3+}/\text{Fe}^{2+}$ half-cell in this experiment. $[E^\circ : \text{Fe}^{3+}/\text{Fe}^{2+} = +0.77\,\text{V}]$
Use the result from (d)(i) to calculate $E_{cell}$ for this electrochemical cell.