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Space-filling representation of the hydroxide ion Ball-and-stick model of the nitrate ion
Molecular formula OH
Molar mass 17.008 g mol-1
Infobox references

Hydroxide is a diatomic anion with chemical formula OH. It consists of an oxygen and a hydrogen atom held together by a covalent bond, and carries a negative electric charge. It is an important but usually minor constituent of water. It functions as a base, a ligand, a nucleophile and a catalyst. The hydroxide ion forms salts, some of which dissociate in aqueous solution, liberating solvated hydroxide ions. Sodium hydroxide is a multi-million-ton per annum commodity chemical. A hydroxide attached to a strongly electropositive center may itself dissociate,[citation needed] liberating a hydrogen cation (H+), making the parent compound an acid.

The corresponding electrically neutral compound •HO is the hydroxyl radical. The corresponding covalently-bound group -OH of atoms is the hydroxyl group. Hydroxide ion and hydroxyl group are nucleophiles and can act as a catalyst in organic chemistry.

Many inorganic substances which bear the word "hydroxide" in their names are not ionic compounds of the hydroxide ion, but covalent compounds which contain hydroxyl groups.

Hydroxide ion[edit]


H+ + OH is in equilibrium with H2O

The equilibrium constant for this reaction, defined as



Schematic representation of the bihydroxide ion[2]

In aqueous solution[3] the hydroxide ion is a base in the Brønsted–Lowry sense as it can accept a proton[note 4] 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).[2] 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.


OH + CO2 is in equilibrium with 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.



When bound to a strongly electron-withdrawing metal centre, hydroxide ligands tend to ionises into oxide ligands. For example, the bichromate ion [HCrO4] dissociates according to

[O3CrO-H] is in equilibrium with [CrO4]2− + H+

with a pKa of about 5.9.[5]

Vibrational spectra[edit]

The infrared spectra of compounds containing the OH functional group have strong absorption bands in the region centered around 3500 cm−1.[6] The high frequency of molecular vibration is a consequence of the small mass of the hydrogen atom as compared to the mass of the oxygen atom and this makes detection of hydroxyl groups by infrared spectroscopy relatively easy. A band due to an OH group tends to be sharp. However, the band width increases when the OH group is involved in hydrogen bonding. A water molecule has an HOH bending mode at about 1600 cm−1, so the absence of this band can be used to distinguish an OH group from a water molecule.

When the OH group is bound to a metal ion in a coordination complex, an M−OH bending mode can be observed. For example, in [Sn(OH)6]2− it occurs at 1065 cm−1. The bending mode for a bridging hydroxide tends to be at a lower frequency as in [(bipyridine)Cu(OH)2Cu(bipyridine)]2+ (955 cm−1).[7] M−OH stretching vibrations occur below about 600 cm−1. For example, the tetrahedral ion [Zn(OH)4]2− has bands at 470 cm−1 (Raman-active, polarized) and 420 cm−1 (infrared). The same ion has an (OH)Zn(OH) bending vibration at 300 cm−1.[8]







Inorganic hydroxides[edit]

Alkali metals[edit]

Aside from NaOH and KOH, which enjoy very large scale applications, the hydroxides of the other alkali metals also are useful. Lithium hydroxide is a strong base, with a pKb of -0.36.[13] Lithium hydroxide is used in breathing gas purification systems for spacecraft, submarines, and rebreathers to remove carbon dioxide from exhaled gas.[14]

2 LiOH + CO2 → Li2CO3 + H2O


Alkaline earth metals[edit]

Trimeric hydrolysis product of beryllium dication.[note 5]
Beryllium hydrolysis as a function of pH
Water molecules attached to Be are omitted

Beryllium hydroxide Be(OH)2 is amphoteric.[16] The hydroxide itself is insoluble in water, with a solubility product log K*sp of −11.7. Addition of acid gives soluble hydrolysis products, including the trimeric ion [Be3(OH)3(H2O)6]3+, which has OH groups bridging between pairs of beryllium ions making a 6-membered ring.[17] At very low pH the aqua ion [Be(H2O)4]2+ is formed. Addition of hydroxide to Be(OH)2 gives the soluble tetrahydroxo anion [Be(OH)4]2−.

The solubility in water of the other hydroxides in this group increases with increasing atomic number.[18] Magnesium hydroxide Mg(OH)2 is a weak base but calcium hydroxide is a strong base as are the hydroxides of the heavier alkaline earths, strontium hydroxide and barium hydroxide. A solution/suspension of calcium hydroxide is known as limewater and can be used to test for the weak acid carbon dioxide. The reaction Ca(OH)2 + CO2 is in equilibrium with Ca2+ + [HCO3] + OH illustrates the strong basicity of calcium hydroxide. Soda lime, which is a mixture of NaOH and Ca(OH)2, is used as a CO2 absorbent.

Boron group elements[edit]






Aluminium hydroxide Al(OH)3 is amphoteric and dissolves in alkaline solution.[16]

Al(OH)3 (solid) + OH (aq) is in equilibrium with [Al(OH)4] (aq)

In the Bayer process[20] for the production of pure aluminium oxide from bauxite minerals this equilibrium is manipulated by careful control of temperature and alkali concentration. In the first phase, aluminium dissolves in hot alkaline solution as [Al(OH)4] but other hydroxides usually present in the mineral, such as iron hydroxides, do not dissolve because they are not amphoteric. After removal of the insolubles, the so-called red mud, pure aluminium hydroxide is made to precipitate by reducing the temperature and adding water to the extract, which, by diluting the alkali, lowers the pH of the solution. Basic aluminium hydroxide AlO(OH), which may be present in bauxite, is also amphoteric.


Carbon group elements[edit]


C(OH)4 → HCO3 + H3O+
HCO3 + H+ is in equilibrium with H2CO3

Carbon dioxide is also known as carbonic anhydride, meaning that it forms by dehydration of carbonic acid H2CO3 (OC(OH)2).[22]

Silicic acid is the name given to a variety of compounds with a generic formula [SiOx(OH)4−2x]n.[23][24] Orthosilicic acid has been identified in very dilute aqueous solution. It is a weak acid with pKa1 = 9.84, pKa2 = 13.2 at 25 °C. It is usually written as H4SiO4 but the formula SiO2(OH)2 is generally accepted .[5] Other silicic acids such as metasilicic acid (H2SiO3), disilicic acid (H2Si2O5), and pyrosilicic acid (H6Si2O7) have been characterized. These acids also have hydroxide groups attached to the silicon; the formulas suggest that these acids are protonated forms of polyoxyanions.

Few hydroxo complexes of germanium have been characterized. Tin(II) hydroxide Sn(OH)2 was prepared in anhydrous media. When tin(II) oxide is treated with alkali the pyramidal hydroxo complex Sn(OH)3 is formed. When solutions containing this ion are acidified the ion [Sn3(OH)4]2+ is formed together with some basic hydroxo complexes. The structure of [Sn3(OH)4]2+ has a triangle of tin atoms connected by bridging hydroxide groups.[25] Tin(IV) hydroxide is unknown but can be regarded as the hypothetical acid from which stannates, with a formula [Sn(OH)6]2−, are derived by reaction with the (Lewis) basic hydroxide ion.[26]


Other main-group elements[edit]

Telluric acid.svg
Orthoperiodic acid.png
Phosphorous acid Phosphoric acid Sulfuric acid Telluric acid orthoperiodic acid

In the higher oxidation states of the elements in groups 5, 6 and 7 there are oxoacids in which the central atom is attached to oxide ions and hydroxide ions. Examples include phosphoric acid H3PO4, and sulfuric acid H2SO4. In these compounds one or more hydroxide groups can dissociate with the liberation of hydrogen cations as in a standard Brønsted–Lowry acid. Many oxoacids of sulfur are known and all feature OH groups that can dissociate.[28]


Ortho-periodic acid[note 6] can lose all its protons, eventually forming the periodate ion [IO4]. It can also be protonated in strongly acidic conditions to give the octahedral ion [I(OH)6]+, completing the isoelectronic series, [E(OH)6]z, E = Sn, Sb, Te, I; z = −2, −1, 0, +1. Other acids of iodine(VII) that contain hydroxide groups are known, in particular in salts such as the mesoperiodate ion that occurs in K4[I2O8(OH)2]·8H2O.[30]

As is common outside of the alkali metals, hydroxides of the elements in lower oxidation states are complicated. For example, phosphorous acid H3PO3 predominantly has the structure OP(H)(OH)2, in equilibrium with a small amount of P(OH)3.[31][32]

The oxoacids of chlorine, bromine and iodine have the formula O(n−1)/2A(OH) where n is the oxidation number: +1, +3 or +5, and A = Cl, Br or I. The only oxoacid of fluorine is F(OH). When these acids are neutralized the hydrogen atom is removed from the hydroxide group.[33]

Transition and post-transition metals[edit]

The hydroxides of the transition metals and post-transition metals usually have the metal in the +2 (M = Mn, Fe, Co, Ni, Cu, Zn) or +3 (M = Fe, Ru, Rh, Ir) oxidation state. None are soluble in water, and many are poorly defined. One complicating feature of the hydroxides is their tendency to undergo further condensation to the oxides, a process called olation. Hydroxides of metals in the +1 oxidation state are also poorly defined or unstable. For example, silver hydroxide Ag(OH) decomposes spontaneously to the oxide (Ag2O). Copper(I) and gold(I) hydroxides are also unstable, although stable adducts of CuOH and AuOH are known.[34] The polymeric compounds M(OH)2 and M(OH)3 are in general prepared by increasing the pH of an aqueous solutions of the corresponding metal cations until the hydroxide precipitates out of solution. On the converse, the hydroxides dissolve in acidic solution. Zinc hydroxide Zn(OH)2 is amphoteric, forming the zincate ion Zn(OH)42− in strongly alkaline solution.[16]

Numerous mixed ligand complexes of these metals with the hydroxide ion exist. In fact these are in general better defined than the simpler derivatives. Many can be made by deprotonation of the corresponding metal aquo complex.


Vanadic acid H3VO4 shows similarities with phosphoric acid H3PO4 though it has a much more complex vanadate oxoanion chemistry. Chromic acid H2CrO4, has similarities with sulfuric acid H2SO4; for example, both form acid salts A+[HMO4]. Some metals, e.g. V, Cr, Nb, Ta, Mo, W, tend to exist in high oxidation states. Rather than forming hydroxides in aqueous solution, they convert to oxo clusters by the process of olation, forming polyoxometalates.[35]

Basic salts containing hydroxide[edit]


The mineral malachite is a typical example of a basic carbonate. The formula, Cu2CO3(OH)2 shows that it is half-way between copper carbonate and copper hydroxide. Indeed, in the past the formula was written as CuCO3·Cu(OH)2. The crystal structure is made up of copper, carbonate and hydroxide ions.[36] The mineral atacamite is an example of a basic chloride. It has the formula, Cu2Cl(OH)3. In this case the composition is nearer to that of the hydroxide than that of the chloride CuCl2·3Cu(OH)2.[37] Copper forms hydroxy phosphate (libethenite), arsenate (olivenite), sulfate (brochantite) and nitrate compounds. White lead is a basic lead carbonate, (PbCO3)2·Pb(OH)2, which has been used as a white pigment because of its opaque quality, though its use is now restricted because it can be a source for lead poisoning.[36]

Structural chemistry[edit]

The hydroxide ion appears to rotate freely in crystals of the heavier alkali metal hydroxides at higher temperatures so as to present itself as a spherical ion, with an effective ionic radius of about 153 pm.[38] Thus, the high-temperature forms of KOH and NaOH have the sodium chloride structure,[39] which gradually freezes in a monocinically distorted sodium chloride structure at temperatures below about 300 °C. The OH groups still rotate even at room temperature around their symmetry axes and, therefore, cannot be detected by X-ray diffraction.[40] The room-temperature form of NaOH has the thallium iodide structure. LiOH, however, has a layered structure, made up of tetrahedral Li(OH)4 and (OH)Li4 units.[38] This is consistent with the weakly basic character of LiOH in solution, indicating that the Li-OH bond has much covalent character.

The hydroxide ion displays cylindrical symmetry in hydroxides of divalent metals Ca, Cd, Mn, Fe, and Co. For example, magnesium hydroxide Mg(OH)2 (brucite) crystallizes with the cadmium iodide layer structure, with a kind of close-packing of magnesium and hydroxide ions.[38][41]


A number of mixed hydroxides are known with stoichiometry A3MIII(OH)6, A2MIV(OH)6 and AMV(OH)6. As the formula suggests these substances contain M(OH)6 octahedral structural units.[45] Layered double hydroxides may be represented by the formula [Mz+1−xM3+x(OH)2]q+(Xn−)q/n·yH2O. Most commonly, z = 2, and M2+ = Ca2+, Mg2+, Mn2+, Fe2+, Co2+, Ni2+, Cu2+ or Zn2+; hence q = x.

In organic reactions[edit]


Base catalysis[edit]

The hydroxide ion may act as a base catalyst.[46] The base abstracts a proton from a weak acid to give an intermediate that goes on to react with another reagent. Common substrates for proton abstraction are alcohols, phenols, amines and carbon acids. The pKa value for dissociation of a C–H bond is extremely high, but the pKa alpha hydrogens of a carbonyl compound are about 3 log units lower. Typical pKa values are 16.7 for acetaldehyde and 19 for acetone.[47] Dissociation can occur in the presence of a suitable base.

RC(O)CH2R′ + B is in equilibrium with RC(O)CHR′ + BH+



OH + EtOH is in equilibrium with EtO + H2O

to produce the ethoxide ion. The pKa for self-dissociation of ethanol is about 16 so the alkoxide ion is a strong enough base[48] The addition of an alcohol to an aldehyde to form a hemiacetal is an example of a reaction that can be catalyzed by the presence of hydroxide. Hydroxide can also act as a Lewis-base catalyst.[49]

As a nucleophilic reagent[edit]


The hydroxide ion is intermediate in nucleophilicity between the fluoride ion F, and the amide ion NH2.[50] The hydrolysis of an ester





  6. ^ The name is not derived from "period", but from "iodine": per-iodic acid (compare iodic acid, perchloric acid), and it is thus pronounced per-iodic /ˌpɜr.ˈɒdɨk/ PURR-eye-OD-ik, and not as /ˌpɪərɪˈɒdɨk/ PEER-ee-OD-ik.


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