Electrolysis of water
Electrolysis of water is using electricity to split water into oxygen (O
2) and hydrogen (H
2) 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.[1]
Principles[edit]
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.[5] 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.[6] 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.
Efficiency is increased through the addition of an electrolyte (such as a salt, an acid or a base) and electrocatalysts.
The decomposition of pure water into hydrogen and oxygen at standard temperature and pressure is not favorable in thermodynamic terms.
Thus, the standard potential of the water electrolysis cell (Eocell = Eocathode − Eoanode) is −1.229 V at 25 °C at pH 0 ([H+] = 1.0 M). At 25 °C with pH 7 ([H+] = 1.0×10−7 M), the potential is unchanged based on the Nernst equation. The thermodynamic standard cell potential can be obtained from standard-state free energy calculations to find ΔG° and then using the equation: ΔG°= −n F E° (where E° is the cell potential and F the Faraday constant, 96,485 C/mol). For two water molecules electrolysed and hence two hydrogen molecules formed, n = 4, and
However, calculations regarding individual electrode equilibrium potentials requires corrections to account for the activity coefficients.[7] In practice when an electrochemical cell is "driven" toward completion by applying reasonable potential, it is kinetically controlled. Therefore, activation energy, ion mobility (diffusion) and concentration, wire resistance, surface hindrance including bubble formation (blocks electrode area), and entropy, require greater potential to overcome. The amount of increase in required potential is termed the overpotential.