Hydroxide Ion
The hydroxide ion is a natural constituent of water, because of the self-ionization reaction:
- H+ + OH− H2O
The equilibrium constant for this reaction, defined as
- Kw =
has a value close to 10−14 at 25 °C, so the concentration of hydroxide ions in pure water is close to 10−7 mol dm−3, in order to satisfy the equal charge constraint. The pH of a solution is equal to the decimal cologarithm of the hydrogen cation concentration; the pH of pure water is close to 7 at ambient temperatures. The concentration of hydroxide ions can be expressed in terms of pOH, which is close to 14 – pH, so pOH of pure water is also close to 7. Addition of a base to water will reduce the hydrogen cation concentration and therefore increase the hydroxide ion concentration (increase pH, decrease pOH) even if the base does not itself contain hydroxide. For example, ammonia solutions have a pH greater than 7 due to the reaction NH3 + H+ NH4+, which results in a decrease in hydrogen cation concentration and an increase in hydroxide ion concentration. pOH can be kept at a nearly constant value with various buffer solutions.
In aqueous solution the hydroxide ion is a base in the Brønsted–Lowry sense as it can accept a proton from a Brønsted–Lowry acid to form a water molecule. It can also act as a Lewis base by donating a pair of electrons to a Lewis acid. In aqueous solution both hydrogen and hydroxide ions are strongly solvated, with hydrogen bonds between oxygen and hydrogen atoms. Indeed, the bihydroxide ion, H3O2−, has been characterized in the solid state. This compound is centrosymmetric and has a very short hydrogen bond (114.5 pm) that is similar to the length in the bifluoride ion, HF2− (114 pm). In aqueous solution the hydroxide ion forms strong hydrogen bonds with water molecules. A consequence of this is that concentrated solutions of sodium hydroxide have high viscosity due to the formation of an extended network of hydrogen bonds as in hydrogen fluoride solutions.
In solution, exposed to air, the hydroxide ion reacts rapidly with atmospheric carbon dioxide, acting as an acid, to form, initially, the bicarbonate ion.
- OH− + CO2 HCO3−
The equilibrium constant for this reaction can be specified either as a reaction with dissolved carbon dioxide or as a reaction with carbon dioxide gas (see carbonic acid for values and details). At neutral or acid pH, the reaction is slow, but is catalyzed by the enzyme carbonic anhydrase, which effectively creates hydroxide ions at the active site.
Solutions containing the hydroxide ion attack glass. In this case, the silicates in glass are acting as acids. Basic hydroxides, whether solids or in solution, are stored in air-tight plastic containers.
The hydroxide ion can function as a typical electron-pair donor ligand, forming such complexes as –. It is also often found in mixed-ligand complexes of the type z+, where L is a ligand. The hydroxide ion often serves as a bridging ligand, donating one pair of electrons to each of the atoms being bridged. As illustrated by 3+, metal hydroxides are often written in a simplified format. It can even act as a 3 electron-pair donor, as in the tetramer 4).
When bound to a strongly electron-withdrawing metal centre, hydroxide ligands tend to ionises into oxide ligands. For example, the bichromate ion, written as –, dissociates according to
- – 2– + H+
with a pKa of about 5.9.
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