Table 1 shows some electrode half-equations and their standard electrode potentials.
| Electrode half-equation | EθE^\thetaEθ / V |
|---|---|
| Cl2(g)+2e−→2Cl−(aq)\text{Cl}_2\text{(g)} + 2\text{e}^- \rightarrow 2\text{Cl}^-\text{(aq)}Cl2(g)+2e−→2Cl−(aq) | +1.36+1.36+1.36 |
| 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 |
| Ag+(aq)+e−→Ag(s)\text{Ag}^+\text{(aq)} + \text{e}^- \rightarrow \text{Ag}\text{(s)}Ag+(aq)+e−→Ag(s) | +0.80+0.80+0.80 |
| 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 |
| SO42−(aq)+4H+(aq)+2e−→SO2(g)+2H2O(l)\text{SO}_4^{2-}\text{(aq)} + 4\text{H}^+\text{(aq)} + 2\text{e}^- \rightarrow \text{SO}_2\text{(g)} + 2\text{H}_2\text{O}\text{(l)}SO42−(aq)+4H+(aq)+2e−→SO2(g)+2H2O(l) | +0.17+0.17+0.17 |
| 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 |
| Zn2+(aq)+2e−→Zn(s)\text{Zn}^{2+}\text{(aq)} + 2\text{e}^- \rightarrow \text{Zn}\text{(s)}Zn2+(aq)+2e−→Zn(s) | −0.76-0.76−0.76 |
Deduce the oxidation state of sulfur in SO42−\text{SO}_4^{2-}SO42− and in SO2\text{SO}_2SO2.
State the weakest reducing agent in Table 1.
Write the conventional representation of the cell that has an EMF of +1.53 V+1.53\text{ V}+1.53 V.
Use data from Table 1 to identify an acid that will oxidise silver. Explain your choice, suggest a possible overall equation for the reaction, and calculate the EMF of the cell that has the same overall reaction.
