Gce Model Questions Vol1 2017 Chemistry Question 35

In acidic solution the redox reaction between permanganate ion and ferrous ion occurs: \(\mathrm{MnO_4^- + Fe^{2+} \rightarrow Mn^{2+} + Fe^{3+}}\). Standard reduction potentials are \(E^\circ_{\mathrm{MnO_4^- / Mn^{2+}}}=+1.51\ \text{V}\) for \(\mathrm{MnO_4^- + 8H^+ +5e^- \rightarrow Mn^{2+} +4H_2O}\) and \(E^\circ_{\mathrm{Fe^{3+} / Fe^{2+}}}=+0.77\ \text{V}\) for \(\mathrm{Fe^{3+} + e^- \rightarrow Fe^{2+}}\). \(\text{(a)}\) Write the balanced overall redox equation. \(\text{(b)}\) Calculate the standard cell potential \(E^\circ_{cell}\). \(\text{(c)}\) If the reaction is carried out at 25 °C with concentrations \([\mathrm{MnO_4^-}]=0.010\ \text{M},\ [\mathrm{Fe^{2+}}]=0.10\ \text{M},\ [\mathrm{Mn^{2+}}]=0.050\ \text{M},\ [\mathrm{Fe^{3+}}]=0.020\ \text{M}\) (acid concentration is 1 M), determine the cell potential using the Nernst equation.