Electrolysis of water is using electricity to split water into oxygen () and hydrogen () gas by electrolysis. Hydrogen gas released in this way can be used as hydrogen fuel, but must be kept apart from the oxygen as the mixture would be extremely explosive. Separately pressurised into convenient 'tanks' or 'gas bottles', hydrogen can be used for oxyhydrogen welding and other applications, as the hydrogen / oxygen flame can reach approximately 2,800°C.
Water electrolysis requires a minimum potential difference of 1.23 volts, although at that voltage external heat is also required. Typically 1.5 volts is required. Electrolysis is rare in industrial applications since hydrogen can be produced less expensively from fossil fuels.
History
In 1789, Jan Rudolph Deiman and Adriaan Paets van Troostwijk used an electrostatic machine to make electricity that was discharged on gold electrodes in a Leyden jar. In 1800 Alessandro Volta invented the voltaic pile, while a few weeks later English scientists William Nicholson and Anthony Carlisle used it to electrolyse water. In 1806 Humphry Davy reported the results of extensive distilled water electrolysis experiments, concluding that nitric acid was produced at the anode from dissolved atmospheric nitrogen. He used a high voltage battery and non-reactive electrodes and vessels such as gold electrode cones that doubled as vessels bridged by damp asbestos. Zénobe Gramme invented the Gramme machine in 1869, making electrolysis a cheap method for hydrogen production. A method of industrial synthesis of hydrogen and oxygen through electrolysis was developed by Dmitry Lachinov in 1888.
Principles
A DC electrical power source is connected to two electrodes, or two plates (typically made from an inert metal such as platinum or iridium) that are placed in the water. Hydrogen appears at the cathode (where electrons enter the water), and oxygen at the anode. Assuming ideal faradaic efficiency, the amount of hydrogen generated is twice the amount of oxygen, and both are proportional to the total electrical charge conducted by the solution. However, in many cells competing side reactions occur, resulting in additional products and less than ideal faradaic efficiency.
Electrolysis of pure water requires excess energy in the form of overpotential to overcome various activation barriers. Without the excess energy, electrolysis occurs slowly or not at all. This is in part due to the limited self-ionization of water.
Pure water has an electrical conductivity about one-millionth that of seawater.
Many electrolytic cells lack requisite electrocatalysts. Efficiency is increased through the addition of an electrolyte (such as a salt, an acid or a base) and electrocatalysts.
Equations
In pure water at the negatively charged cathode, a reduction reaction takes place, with electrons (e−) from the cathode being given to hydrogen cations to form hydrogen gas. At the positively charged anode, an oxidation reaction occurs, generating oxygen gas and giving electrons to the anode to complete the circuit.
The two half-reactions, reduction and oxidation, are coupled to form a balanced system. In order to balance each half-reaction, the water needs to be acidic or basic. In the presence of acid, the equations are:
In the presence of base, the equations are:
Combining either half reaction pair yields the same overall decomposition of water into oxygen and hydrogen:
The number of hydrogen molecules produced is thus twice the number of oxygen molecules, in keeping with the facts that both hydrogen and oxygen are diatomic molecules and water molecules contain twice as many hydrogen atoms as oxygen atoms. Assuming equal temperature and pressure for both gases, volume is proportional to moles, so twice as large a volume of hydrogen gas is produced as oxygen gas. The number of electrons pushed through the water is twice the number of generated hydrogen molecules and four times the number of generated oxygen molecules.
Thermodynamics
The decomposition of pure water into hydrogen and oxygen at standard temperature and pressure is not favorable in thermodynamic terms.