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Question
Using the standard electrode potential, predict if the reaction between the following is feasible:
Ag(s) and \[\ce{Fe^{3+}_{( aq)}}\]
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Solution
A reaction is possible if the value of \[\ce{E^{\circ}_{cell}}\] is feasible.
The reaction is as follows:
\[\ce{Ag_{(s)} + Fe^{3+}_{( aq)} -> Ag^{+}_{( aq)} + Fe^{2+}_{( aq)}}\]
According to this, the cell will be as follows:
\[\ce{Ag_{(s)} | Ag^{+}_{( aq)} || Fe^{3+}_{( aq)} | Fe^{2+}_{( aq)}}\]
∴ \[\ce{E^{\circ}_{cell} = E^{\circ}_{Fe^{3+}/Fe^{2+}} - E^{\circ}_{Ag^+/Ag}}\]
= 0.77 − 0.80
= −0.03 V
Since the value of \[\ce{E^{\circ}_{cell}}\] is negative, the reaction is not feasible.
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