$$\begin{align} \ce{O2 + 4H+ + 4e- \;&<=> 2H2O}\quad\quad &&E^\circ = +1.229\ \mathrm{V}\\ \ce{S2O8^2- + 2e- \;&<=> 2SO4^2-}\quad\quad &&E^\circ = +2.01\ \mathrm{V} \end{align}$$

From the above, I know that an aqueous solution of a sulfate salt undergoing electrolysis will not oxidise to peroxydisulfate ions over water being oxidised to oxygen at the anode unless it is at extremely high concentrations.

However, what if the solution is molten? Will the product at the anode be peroxydisulfate?

Also, can nitrate ions be oxidised? Perhaps to peroxynitrate or peroxynitrite?

  • 2
    $\begingroup$ Peroxysulfates are industrially created in aqueous electrolysis. Nothing extreme is needed, just concentrated solution and high density current. You didn't think about overpotential here, it seems. $\endgroup$
    – Mithoron
    Commented Oct 19, 2017 at 14:08

1 Answer 1


Not only they can, but they must get oxidized. There is no other option. This is the only way the electric current can pass between the electrode and the solution.

Whether or not the products will be some peroxyanions is another question. At high temperatures typical of molten salts, I don't think they will survive. Hence I would expect $\ce{O2}$ + disulfate in the first case and $\ce{O2 + NO2}$ in the second.

  • $\begingroup$ The proposed anodic products are interesting for another reason. They would seem to suggest that nitrogen is reduced from +5 to +4 at the anode! Might we have to abandon our usual oxidation state convention, which is geared to a water-based world, for nitrate solvent systems? $\endgroup$ Commented Oct 19, 2017 at 13:26
  • 1
    $\begingroup$ Nitrogen would get reduced because it has nowhere else to go. There are no binitrates (unlike bisulfates), and $\ce{N2O5}$ is not particularly stable. This has little to do with electrolysis. $\endgroup$ Commented Oct 19, 2017 at 13:30
  • $\begingroup$ Note the autoionization of dinitrogen tetroxide given by en.wikipedia.org/wiki/Molecular_autoionization. Oxygen seems to transfer with one negative charge. Might the effective oxidation state of oxygen in nitrogen oxide/nitrate solvent systems be $-1$ instead of $-2$? It would give a more parsimonious description of the anodic reaction of nitrate ion in nitrate solvents. $\endgroup$ Commented Oct 19, 2017 at 14:33
  • 1
    $\begingroup$ Whatever. Oxidation states are human invention, not a fact of nature; they are all arbitrary anyway. $\endgroup$ Commented Oct 19, 2017 at 14:37

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