Electrolysis

Two electrodes made from an inert metal (such as platinum, stainless steel or iridium) are connected to a power source and placed in water.

When the current is switched on HYDROGEN will form on the CATHODE electrode (negatively charged electrode – electrons released into the water)

A REDUCTION REACTION takes place at the negatively charged cathode. Electrons (e−) from the cathode are given to hydrogen cations to form hydrogen:

Reduction at cathode: 2H⁺_(aq) + 2e⁻ → H_{2 (g)}

OXYGEN will form on the ANODE electrode (the positively charged electrode).

An OXIDATION REACTION takes place at the positively charged ANODE, releasing oxygen gas releasing electrons to the anode to complete the circuit:

Oxidation at anode: 2 H₂O_(l) → O_{2 (g)} + 4 H⁺ _(aq) + 4e⁻

Adding the two half reactions :-

Cathode (reduction): 2 H₂O (l) + 2e⁻ → H_{2 (g)} + 2 OH ^– _(aq)

Anode (oxidation): 4 OH ^– _(aq) → O_{2 (g)} + 2 H₂O _(l) + 4 e⁻

This results in the overall decomposition of water into oxygen and hydrogen:

Overall reaction: 2 H₂O _(l) → 2 H_{2 (g)} + O_{2 (g)}

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Viewing the time-lapse video it is possible to identify that the amount of Hydrogen gas produced in the right hand chamber is twice as much as the amount of Oxygen produced in the left hand chamber. This gives greater meaning to the formula of water as H2O where there are two parts of Hydrogen to 1 Part of Oxygen.

This can be visualized in the following animation.:

Pure water is a poor conductor of electricity and will require more energy to bring about electrolysis.