Electrolysis of water
How is the "overall reaction" of electrolysis of water explained?
Electrochemistry
Water can be converted to its component elemental gasses, H2 and O2 through the application of an external voltage. Water doesn't decompose into hydrogen and oxygen spontaneously as the Gibbs free energy for the process at standard conditions is about 474.4 kJ. The decomposition of water into hydrogen and oxygen can be performed in an electrolytic cell. In it, a pair of inert electrodes usually made of platinum immersed in water act as anode and cathode in the electrolytic process. The electrolysis starts with the application of an external voltage between the electrodes. This process will not occur except at extremely high voltages without an electrolyte such as sodium chloride or sulfuric acid (most used 0.1 M). Bubbles from the gases will be seen near both electrodes. The following half reactions describe the process mentioned above: "Anode (oxidation)": 2 H2O(l) → O2(g) + 4 H+(aq) + 4 e– "Cathode (reduction)": 2 H2O(g) + 2 e– → H2(g) + 2 OH–(aq) "Overall reaction": 2 H2O(l) → 2 H2(g) + O2(g) Although strong acids may be used in the apparatus, the reaction will not net consume the acid. While this reaction will work at any conductive electrode at a sufficiently large potential, platinum catalyzes both hydrogen and oxygen formation, allowing for relatively mild voltages (~2 V depending on the pH).