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A galvanic cell is set up using the Ag+/Ag and Fe3+/Fe2+ half-cells for the reaction: Fe2+ + Ag+ -> Fe3+ + Ag. If the concentrations of Fe2+ and Fe3+ are equal, find the value of [Ag+] that makes the cell voltage exactly zero. Given: E_cell(Ag+/Ag) = 0.80 V, E_cell(Fe3+/Fe2+) = 0.77 V, and 2.303RT/F = 0.06.
- antilog(-0.5)
- antilog(1.03)
- antilog(0.1)
- antilog(-0.02)
Correct answer: antilog(-0.5)
Solution
When E_cell = 0, both half-cell potentials must be equal. With [Fe3+]=[Fe2+], the Fe half-cell sits at its standard potential 0.77 V, so the Ag half-cell must also equal 0.77 V. Nernst for Ag+/Ag gives 0.80 + 0.06*log[Ag+] = 0.77, yielding log[Ag+] = -0.5, so [Ag+] = antilog(-0.5).
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