In acidic solution the redox reaction $\ce{MnO4^- + Fe^{2+} - Mn^{2+} + Fe^{3+}}$ occurs. Standard potentials are $E^{\circ}_{\ce{MnO4^-/Mn^{2+}}}=+1.51\ \text{V}$ and $E^{\circ}_{\ce{Fe^{3+}/Fe^{2+}}}=+0.77\ \text{V}$. The solution contains 0.010 M $\ce{MnO4^-}$ and 0.10 M $\ce{Fe^{2+}}$. (i) Calculate the cell potential at 25 °C using the Nernst equation, assuming product activities are unity. (ii) State which species is reduced and which is oxidised.

Question

In acidic solution the redox reaction $\ce{MnO4^- + Fe^{2+} -> Mn^{2+} + Fe^{3+}}$ occurs. Standard potentials are $E^{\circ}_{\ce{MnO4^-/Mn^{2+}}}=+1.51\ \text{V}$ and $E^{\circ}_{\ce{Fe^{3+}/Fe^{2+}}}=+0.77\ \text{V}$. The solution contains 0.010 M $\ce{MnO4^-}$ and 0.10 M $\ce{Fe^{2+}}$. (i) Calculate the cell potential at 25 °C using the Nernst equation, assuming product activities are unity. (ii) State which species is reduced and which is oxidised.

Accepted Answer

E = 0.66 V; MnO4⁻ is reduced, Fe²⁺ is oxidised

Neco 2020 Chemistry Question 50

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