Q.iii. Write the electrode reactions during electrolysis of molten KCl.

Electrolysis is the electrical-energy-driven decomposition of a molten ionic compound or an aqueous electrolyte solution: in an electrolytic cell, oxidation occurs at the anode (species loses electrons, pulled out by the battery) and reduction at the cathode (species gains electrons supplied by the battery), the battery acting as an 'electron pump'. For molten NaCl, $Cl^-$ is oxidised to $Cl_2$ gas at the anode and $Na^+$ reduced to molten Na at the cathode: $2Na^+(l)+2Cl^-(l)\rightarrow2Na(l)+Cl_2(g)$, a non-spontaneous change forced by the applied voltage. For AQUEOUS NaCl, water competes at both electrodes: at the cathode, water ($E^0=-0.83V$) is reduced in preference to $Na^+$ ($E^0=-2.71V$), giving $H_2$ gas instead of Na metal; at the anode, $Cl^-$ is oxidised to $Cl_2$ gas rather than water to $O_2$, despite water's more favourable standard potential, because of a kinetic effect called overvoltage. The net result, $2Cl^-(aq)+2H_2O(l)\rightarrow Cl_2(g)+H_2(g)+2OH^-(aq)$, converts NaCl solution into NaOH solution — the industrial route by which NaOH is manufactured from NaCl.