An electrochemical cell is set up to measure E°cell of a cell consisting of an Fe3+ /Fe2+ half-cell and a Cl2/Cl- half-cell. Include all necessary substances. It is not necessary to state conditions used. [Important values, constants and standards] [Periodic Table]

1. (a)Define standard cell potential, E°cell.[1]
2. (b)Draw a labelled diagram of this electrochemical cell.[3]
3. (c)The cell reaction for the electrochemical cell in (b) is shown. Cl2 + 2Fe2+ → 2Fe3+ + 2Cl- E°cell = +0.59V Calculate ΔG°, in kJ mol−1, for this cell reaction.[2]
4. (d)(i)Another experiment is set up using the same electrochemical cell. In this experiment the Fe2+ concentration is 0.15moldm¯³. All other concentrations remain at their standard values. The Nernst equation is shown. E = E° + (0.059/z) log [oxidised species] / [reduced species] Use the Nernst equation to calculate the electrode potential, E, for the Fe3+ /Fe2+ half-cell in this experiment. [E° : Fe3+/Fe2+ = +0.77V] E =[1]
5. (d)(ii)Use your answer to (d)(i) to calculate Ecell for this electrochemical cell. Ecell =[1]

This question uses a diagram from the original paper, which isn't shown here. Open the full paper from the past papers archive to see it.