The reaction taking place in an electrochemical cell under standard conditions is Fe²⁺(aq) + Ag⁺(aq) Fe³⁺(aq) + Ag(s)

a. Write two half-equations for this reaction. For each, state whether oxidation or reduction is occurring.

b. The standard electrode potential for the half-cell containing Fe²⁺(aq) and Fe³⁺(aq) is +0.77V.

i. Use the relationship E = E^θ + 0.059/Z  log₁₀ [oxidized form]/[reduced form] to calculate the electrode potential at 298 K if the concentration of Fe²⁺(aq) is 0.02 mol dm³ and the concentration of Fe³⁺(aq) is 0.1 mol dm^–3.

ii. Use the relationship above to explain why the standard electrode potential for the half cell containing Fe²⁺(aq) and Fe³⁺(aq) is always +0.77 V if there are equimolar concentrations of Fe²⁺(aq) and Fe³⁺(aq).

c. The standard electrode potential for the half-cell containing Ag⁺(aq) and Ag(s) is +0.80 V. Calculate the electrode potential at 298K if the concentration of Ag⁺(aq) is 0.05 mol dm^–3.

d. Use your results to parts b i and c to predict whether the reaction Fe²⁺(aq) + Ag⁺(aq) → Fe³⁺(aq) + Ag(s) is likely to occur at the concentrations Fe²⁺(aq) 0.05 mol dm^–3, Fe³⁺(aq) 0.1 mol dm^–3 and Ag+(aq) 0.05 mol dm^–3. Explain your answer.