Автопротолиз воды и равновесие ионов гидроксония и гидроксид-ионов

Autoprotolysis or exchange of a proton between two water molecules

The self ionization of water (also autoionization of water, autoprotolysis of water, autodissociation of water, or simply dissociation of water) is an ionization reaction in pure water or in an aqueous solution, in which a water molecule, H2O, deprotonates (loses the nucleus of one of its hydrogen atoms) to become a hydroxide ion, OH−. The hydrogen nucleus, H+, immediately protonates another water molecule to form a hydronium cation, H3O+. It is an example of autoprotolysis, and exemplifies the amphoteric nature of water.

The self ionization of water was first proposed in 1884 by Svante Arrhenius as part of the theory of ionic dissociation which he proposed to explain the conductivity of electrolytes including water. Arrhenius wrote the self ionization as H2O <=> H+ + OH At that time, nothing was yet known of atomic structure or subatomic particles, so he had no reason to consider the formation of an H+ ion from a hydrogen atom on electrolysis as any less likely than, say, the formation of a Na+ ion from a sodium atom. In 1923 Johannes Nicolaus Brønsted and Martin Lowry proposed that the self ionization of water actually involves two water molecules: H2O + H2O <=> H3O+ + OH By this time the electron and the nucleus had been discovered and Rutherford had shown that a nucleus is very much smaller than an atom. This would include a bare ion H+ which would correspond to a proton with zero electrons. Brønsted and Lowry proposed that this ion does not exist free in solution, but always attaches itself to a water (or other solvent) molecule to form the hydronium ion H3O+ (or other protonated solvent). Later spectroscopic evidence has shown that many protons are actually hydrated by more than one water molecule. The most descriptive notation for the hydrated ion is H+(aq), where aq (for aqueous) indicates an indefinite or variable number of water molecules. However the notations H+ and H3O+ are still also used extensively because of their historical importance. This article mostly represents the hydrated proton as H3O+, corresponding to hydration by a single water molecule.

In water–heavy water mixtures equilibria several species are involved: H2O, HDO, D2O, H3O+, D3O+, H2DO+, HD2O+, HO−, DO−.

Water molecules dissociate into equal amounts of H3O+ and OH−, so their concentrations are almost exactly 1.00 at 25 °C and 0.1 MPa. A solution in which the H3O+ and OH− concentrations equal each other is considered a neutral solution. In general, the pH of the neutral point is numerically equal to pKw. Pure water is neutral, but most water samples contain impurities. If an impurity is an acid or base, this will affect the concentrations of hydronium ion and hydroxide ion. Water samples that are exposed to air will absorb some carbon dioxide to form carbonic acid (H2CO3) and the concentration of H3O+ will increase due to the reaction H2CO3 + H2O = HCO3− + H3O+. The concentration of OH− will decrease in such a way that the product [H3O+][OH−] remains constant for fixed temperature and pressure. Thus these water samples will be slightly acidic. If a pH of exactly 7.0 is required, it must be maintained with an appropriate buffer solution.