Table 1 shows some electrode half-equations and their standard electrode potentials.
Table 1
| Electrode half-equation | EθE^\thetaEθ / V |
|---|---|
| VO2+(aq)+2H+(aq)+e−→VO2+(aq)+H2O(l)\text{VO}_2^+\text{(aq)} + 2\text{H}^+\text{(aq)} + \text{e}^- \rightarrow \text{VO}^{2+}\text{(aq)} + \text{H}_2\text{O}\text{(l)}VO2+(aq)+2H+(aq)+e−→VO2+(aq)+H2O(l) | +1.00+1.00+1.00 |
| NO3−(aq)+4H+(aq)+3e−→NO(g)+2H2O(l)\text{NO}_3^-\text{(aq)} + 4\text{H}^+\text{(aq)} + 3\text{e}^- \rightarrow \text{NO}\text{(g)} + 2\text{H}_2\text{O}\text{(l)}NO3−(aq)+4H+(aq)+3e−→NO(g)+2H2O(l) | +0.96+0.96+0.96 |
| Fe3+(aq)+e−→Fe2+(aq)\text{Fe}^{3+}\text{(aq)} + \text{e}^- \rightarrow \text{Fe}^{2+}\text{(aq)}Fe3+(aq)+e−→Fe2+(aq) | +0.77+0.77+0.77 |
| Cu2+(aq)+2e−→Cu(s)\text{Cu}^{2+}\text{(aq)} + 2\text{e}^- \rightarrow \text{Cu}\text{(s)}Cu2+(aq)+2e−→Cu(s) | +0.34+0.34+0.34 |
| 2H+(aq)+2e−→H2(g)2\text{H}^+\text{(aq)} + 2\text{e}^- \rightarrow \text{H}_2\text{(g)}2H+(aq)+2e−→H2(g) | 0.000.000.00 |
| V3+(aq)+e−→V2+(aq)\text{V}^{3+}\text{(aq)} + \text{e}^- \rightarrow \text{V}^{2+}\text{(aq)}V3+(aq)+e−→V2+(aq) | −0.26-0.26−0.26 |
| Cr3+(aq)+3e−→Cr(s)\text{Cr}^{3+}\text{(aq)} + 3\text{e}^- \rightarrow \text{Cr}\text{(s)}Cr3+(aq)+3e−→Cr(s) | −0.74-0.74−0.74 |
Deduce the oxidation state of vanadium in VO2+\text{VO}_2^+VO2+ and in VO2+\text{VO}^{2+}VO2+.
State the weakest reducing agent in Table 1.
Write the conventional representation of the cell that has an EMF of +1.03 V+1.03\text{ V}+1.03 V.
Use data from Table 1 to identify an acid that will oxidise copper metal. Explain your choice, suggest a possible overall equation for the reaction, and calculate the EMF of the cell that has the same overall reaction.
