Hydroxide is a diatomic anion with chemical formula OH−. It consists
of an oxygen and 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
Sodium hydroxide is a multi-million-ton per annum commodity
chemical. A hydroxide attached to a strongly electropositive center
may itself ionize, 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.
1.1 Vibrational spectra
3.1 Alkali metals
3.2 Alkaline earth metals
3.3 Boron group elements
3.4 Carbon group elements
3.5 Other main-group elements
3.6 Transition and post-transition metals
4 Basic salts containing hydroxide
5 Structural chemistry
6 In organic reactions
6.1 Base catalysis
6.2 As a nucleophilic reagent
The hydroxide ion is a natural part of water, because of the
H3O+ + OH− ⇌ 2H2O
The equilibrium constant for this reaction, defined as
Kw = [H+][OH−][note 1]
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;[note 2] 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,[note 3] 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+ ⇌ NH+
4, 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.
Schematic representation of the bihydroxide ion
In aqueous solution 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 H
2 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 HF−
2 (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
In solution, exposed to air, the hydroxide ion reacts rapidly with
atmospheric carbon dioxide, acting as an acid, to form, initially, the
OH− + CO2 ⇌ HCO−
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
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 airtight plastic containers.
The hydroxide ion can function as a typical electron-pair donor
ligand, forming such complexes as [Al(OH)4]−. It is also often found
in mixed-ligand complexes of the type [MLx(OH)y]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 [Pb2(OH)]3+, metal hydroxides are often written in a
simplified format. It can even act as a 3-electron-pair donor, as in
the tetramer [PtMe3(OH)]4.
When bound to a strongly electron-withdrawing metal centre, hydroxide
ligands tend to ionise into oxide ligands. For example, the bichromate
ion [HCrO4]− dissociates according to
[O3CrO–H]− ⇌ [CrO4]2− + H+
with a pKa of about 5.9.
The infrared spectra of compounds containing the OH functional group
have strong absorption bands in the region centered around
3500 cm−1. 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). 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 a (HO)–Zn–(OH) bending vibration at 300 cm−1.
Sodium hydroxide solutions, also known as lye and caustic soda, are
used in the manufacture of pulp and paper, textiles, drinking water,
soaps and detergents, and as a drain cleaner. Worldwide production in
2004 was approximately 60 million tonnes. The principal method
of manufacture is the chlor-alkali process.
Solutions containing the hydroxide ion are generated when a salt of a
weak acid is dissolved in water.
Sodium carbonate is used as an
alkali, for example, by virtue of the hydrolysis reaction
3 + H2O ⇌ HCO−
3 + OH−; (pKa2 = 10.33 at 25 °C and
zero ionic strength)
Although the base strength of sodium carbonate solutions is lower than
a concentrated sodium hydroxide solution, it has the advantage of
being a solid. It is also manufactured on a vast scale (42 million
tonnes in 2005) by the Solvay process. An example of the use of
sodium carbonate as an alkali is when washing soda (another name for
sodium carbonate) acts on insoluble esters, such as triglycerides,
commonly known as fats, to hydrolyze them and make them soluble.
Bauxite, a basic hydroxide of aluminium, is the principal ore from
which the metal is manufactured. Similarly, goethite (α-FeO(OH))
and lepidocrocite (γ-FeO(OH)), basic hydroxides of iron, are among
the principal ores used for the manufacture of metallic iron.
Numerous other uses can be found in the articles on individual
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. Lithium
hydroxide is used in breathing gas purification systems for
spacecraft, submarines, and rebreathers to remove carbon dioxide from
2 LiOH + CO2 → Li2CO3 + H2O
The hydroxide of lithium is preferred to that of sodium because of its
lower mass. Sodium hydroxide, potassium hydroxide and the hydroxides
of the other alkali metals are also strong bases.
Alkaline earth metals
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. 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. 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. Magnesium hydroxide
Mg(OH)2 is a strong base (up to the limit of its solubility, which is
very low in pure water), as are the hydroxides of the heavier alkaline
earths: calcium hydroxide, 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 ⇌ Ca2+ + HCO−
3 + OH− illustrates the basicity of calcium hydroxide. Soda lime,
which is a mixture of the strong bases NaOH and KOH with Ca(OH)2, is
used as a CO2 absorbent.
Boron group elements
Aluminium hydrolysis as a function of pH.
Water molecules attached to
Al are omitted
The simplest hydroxide of boron B(OH)3, known as boric acid, is an
acid. Unlike the hydroxides of the alkali and alkaline earth
hydroxides, it does not dissociate in aqueous solution. Instead, it
reacts with water molecules acting as a Lewis acid, releasing protons.
B(OH)3 + H2O ⇌ B(OH)−
4 + H+
A variety of oxyanions of boron are known, which, in the protonated
form, contain hydroxide groups.
Aluminium hydroxide Al(OH)3 is amphoteric and dissolves in alkaline
Al(OH)3 (solid) + OH− (aq) ⇌ Al(OH)−
In the Bayer process 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.
In mildly acidic solutions the hydroxo complexes formed by aluminium
are somewhat different from those of boron, reflecting the greater
size of Al(III) vs. B(III). The concentration of the species
[Al13(OH)32]7+ is very dependent on the total aluminium concentration.
Various other hydroxo complexes are found in crystalline compounds.
Perhaps the most important is the basic hydroxide AlO(OH), a polymeric
material known by the names of the mineral forms boehmite or diaspore,
depending on crystal structure. Gallium hydroxide, indium
hydroxide and thallium(III) hydroxides are also amphoteric.
Thallium(I) hydroxide is a strong base.
Carbon group elements
Carbon forms no simple hydroxides. The hypothetical compound C(OH)4
(orthocarbonic acid or methanetetrol) is unstable in aqueous
C(OH)4 → HCO−
3 + H3O+
3 + H+ ⇌ H2CO3
Carbon dioxide is also known as carbonic anhydride, meaning that it
forms by dehydration of carbonic acid H2CO3 (OC(OH)2).
Silicic acid is the name given to a variety of compounds with a
generic formula [SiOx(OH)4−2x]n. 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 Si(OH)4 is generally
accepted.[dubious – discuss] 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.
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
Hydrolysis of Pb2+ in aqueous solution is accompanied by the formation
of various hydroxo-containing complexes, some of which are insoluble.
The basic hydroxo complex [Pb6O(OH)6]4+ is a cluster of six lead
centres with metal–metal bonds surrounding a central oxide ion. The
six hydroxide groups lie on the faces of the two external Pb4
tetrahedra. In strongly alkaline solutions soluble plumbate ions are
formed, including [Pb(OH)6]2−.
Other main-group elements
In the higher oxidation states of the pnictogens, chalcogens,
halogens, and noble gases 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
Telluric acid is often written with the formula H2TeO4·2H2O but is
better described structurally as Te(OH)6.
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.
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.
The oxoacids of chlorine, bromine and iodine have the formula
On−1/2A(OH) where n is the oxidation number: +1, +3, +5 or +7, and A
= Cl, Br or I. The only oxoacid of fluorine is F(OH), hypofluorous
acid. When these acids are neutralized the hydrogen atom is removed
from the hydroxide group.
Transition and post-transition metals
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. 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
Zinc hydroxide Zn(OH)2 is amphoteric, forming the zincate
ion Zn(OH)42− in strongly alkaline solution.
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.
LnM(OH2) + B ⇌ LnM(OH) + BH+ (L = ligand, B = base)
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.
Basic salts containing hydroxide
In some cases the products of partial hydrolysis of metal ion,
described above, can be found in crystalline compounds. A striking
example is found with zirconium(IV). Because of the high oxidation
state, salts of Zr4+ are extensively hydrolyzed in water even at low
pH. The compound originally formulated as ZrOCl2·8H2O was found to be
the chloride salt of a tetrameric cation [Zr4(OH)8(H2O)16]8+ in which
there is a square of Zr4+ ions with two hydroxide groups bridging
between Zr atoms on each side of the square and with four water
molecules attached to each Zr atom.
The mineral malachite is a typical example of a basic carbonate. The
formula, Cu2CO3(OH)2 shows that it is halfway 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. 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. 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
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. Thus, the high-temperature forms of KOH and NaOH have
the sodium chloride structure, 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. 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. 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
The amphoteric hydroxide Al(OH)3 has four major crystalline forms:
gibbsite (most stable), bayerite, nordstrandite and doyleite.[note 7]
All these polymorphs are built up of double layers of hydroxide
ions – the aluminium atoms on two-thirds of the octahedral
holes between the two layers – and differ only in the stacking
sequence of the layers. The structures are similar to the brucite
structure. However, whereas the brucite structure can be described as
a close-packed structure in gibbsite the OH groups on the underside of
one layer rest on the groups of the layer below. This arrangement led
to the suggestion that there are directional bonds between OH groups
in adjacent layers. This is an unusual form of hydrogen bonding
since the two hydroxide ion involved would be expected to point away
from each other. The hydrogen atoms have been located by neutron
diffraction experiments on α-AlO(OH) (diaspore). The O–H–O
distance is very short, at 265 pm; the hydrogen is not
equidistant between the oxygen atoms and the short OH bond makes an
angle of 12° with the O–O line. A similar type of hydrogen bond
has been proposed for other amphoteric hydroxides, including Be(OH)2,
Zn(OH)2 and Fe(OH)3
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. Layered double
hydroxides may be represented by the formula [Mz+
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
Potassium hydroxide and sodium hydroxide are two well-known reagents
in organic chemistry.
The hydroxide ion may act as a base catalyst. 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.
Dissociation can occur in the presence of a suitable base.
RC(O)CH2R′ + B ⇌ RC(O)CH−R′ + BH+
The base should have a pKa value not less than about 4 log units
smaller or the equilibrium will lie almost completely to the left.
The hydroxide ion by itself is not a strong enough base, but it can be
converted in one by adding sodium hydroxide to ethanol
OH− + EtOH ⇌ 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 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 can also act as a Lewis-base catalyst.
As a nucleophilic reagent
Nucleophilic acyl substitution
Nucleophilic acyl substitution with nucleophile (Nu) and leaving group
The hydroxide ion is intermediate in nucleophilicity between the
fluoride ion F−, and the amide ion NH−
2. The hydrolysis of an ester
R1C(O)OR2 + H2O ⇌ R1C(O)OH + HOR2
also known as saponification is an example of a nucleophilic acyl
substitution with the hydroxide ion acting as a nucleophile. In this
case the leaving group is an alkoxide ion, which immediately removes a
proton from a water molecule to form an alcohol. In the manufacture of
soap, sodium chloride is added to salt out the sodium salt of the
carboxylic acid; this is an example of the application of the
Other cases where hydroxide can act as a nucleophilic reagent are
amide hydrolysis, the Cannizzaro reaction, nucleophilic aliphatic
substitution, nucleophilic aromatic substitution and in elimination
reactions. The reaction medium for KOH and NaOH is usually water but
with a phase-transfer catalyst the hydroxide anion can be shuttled
into an organic solvent as well, for example in the generation of
^ [H+] denotes the concentration of hydrogen cations and [OH−] the
concentration of hydroxide ions
^ Strictly speaking pH is the cologarithm of the hydrogen cation
^ pOH signifies the minus the logarithm to base 10 of OH− ,
alternatively the logarithm of 1/ OH−
^ In this context proton is the term used for a solvated hydrogen
^ In aqueous solution the ligands L are water molecules, but they may
be replaced by other ligands
^ The name is not derived from "period", but from "iodine": per-iodic
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