Question

At 298 K, the standard electrode potentials of Cu⁺²/Cu, Zn⁺²/Zn, Fe⁺²/Fe and Ag⁺/Ag are 0.34 V, −0.76 V, −0.44 V and −0.80 V respectively.

On the basis of standard electrode potential, predict which of the following reaction cannot occur?

Options

• \[\ce{CuSO4_{(aq)} + Fe_{(s)} -> FeSO4_{(aq)} + Cu_{(s)}}\]

• \[\ce{FeSO4_{(aq)} + Zn_{(s)} -> ZnSO4_{(aq)} + Fe_{(s)}}\]

• \[\ce{2CuSO4_{(aq)} + 2Ag_{(s)} -> 2Cu_{(s)} + Ag2SO4_{(aq)}}\]

• \[\ce{CuSO4_{(aq)} + Zn_{(s)} -> ZnSO4_{(aq)} + Cu_{(s)}}\]

Answer 1

\[\ce{2CuSO4_{(aq)} + 2Ag_{(s)} -> 2Cu_{(s)} + Ag2SO4_{(aq)}}\]

Explanation:

A redox reaction occurs if the metal with lower E° (stronger reducing agent) gets oxidised.

\[\ce{E^{\circ}_{{Cu^{2+}/{Cu}}}}\] = +0.34 V

\[\ce{E^{\circ}_{{Ag^{+}/{Ag}}}}\] = −0.80 V

Since Ag has a more negative E°, it should not reduce Cu²⁺ (which has a higher E°).

Hence, this reaction cannot occur.

So, reaction \[\ce{2CuSO4_{(aq)} + 2Ag_{(s)} -> 2Cu_{(s)} + Ag2SO4_{(aq)}}\] is not feasible based on standard electrode potentials.