Electrolysis of Water

The decomposition of pure water into hydrogen and oxygen is a thermodynamically unfavorable process, requiring external energy in the form of electricity. Pure water is a poor conductor of electricity, so an electrolyte, such as a small amount of a soluble salt or an acid like sulfuric acid, is typically added to increase conductivity. The overall balanced equation for the reaction is [latex]2H_2O(l) \rightarrow 2H_2(g) + O_2(g)[/latex].

The process occurs in an electrolytic cell. At the negatively charged cathode, a reduction reaction takes place: [latex]2H_2O(l) + 2e^- \rightarrow H_2(g) + 2OH^-(aq)[/latex]. At the positively charged anode, an oxidation reaction occurs: [latex]2H_2O(l) \rightarrow O_2(g) + 4H^+(aq) + 4e^-[/latex]. The production of hydrogen gas at the cathode is exactly twice the volume of oxygen gas produced at the anode, a direct consequence of the stoichiometry of water. This 2:1 volume ratio is a classic demonstration in chemistry classrooms, often performed using a Hofmann voltameter.

The minimum voltage required for electrolysis to occur, known as the decomposition potential, is 1.23 V at standard conditions. However, in practice, a higher voltage, called overpotential, is needed to overcome various activation barriers. The efficiency of water electrolysis is a key factor in the viability of a hydrogen-based economy, with significant research focused on developing more effective and cheaper catalysts for the anode and cathode to reduce the overpotential and energy consumption.