For \( \text{NO}_3^- \) to oxidize a metal, \( \text{NO}_3^- \) must itself be reduced. The given reduction half-reaction is:
\( \text{NO}_3^- + 4\text{H}^+ + 3e^- \rightarrow \text{NO} + 2\text{H}_2\text{O} \qquad E^0 = 0.97 \text{ V} \)
For a metal to be oxidized by \( \text{NO}_3^- \), the overall cell potential (\( E^0_{cell} \)) for the redox reaction must be positive. The cell potential is calculated as:
\( E^0_{cell} = E^0_{reduction} - E^0_{oxidation} \)
In our case, \( E^0_{reduction} \) is 0.97 V (for the \( \text{NO}_3^- \) reduction). For the cell potential to be positive, \( E^0_{oxidation} \) must be less than 0.97 V.
Since we are given standard reduction potentials, we need to reverse the sign to obtain the oxidation potentials for each metal:
Comparing these values to 0.97 V, we find that Fe, Ag, and Au have oxidation potentials less than 0.97 V. Therefore, these three metals can be oxidized by \( \text{NO}_3^- \) in aqueous solution.
Vanadium (V) has an oxidation potential greater than 0.97 V, so it will not be oxidized by \( \text{NO}_3^- \). So, the number of metal(s) which will be oxidized by \( \text{NO}_3^- \) in aqueous solution is 3.
(b.)Calculate the potential for half-cell containing 0.01 M K\(_2\)Cr\(_2\)O\(_7\)(aq), 0.01 M Cr\(^{3+}\)(aq), and 1.0 x 10\(^{-4}\) M H\(^+\)(aq).