The hydroxide ion is a diatomic anion with chemical formula OH⁻. It consists of an oxygen and a hydrogen atom held together by a covalent bond, and carrying a negative electric charge. It is a natural constituent of water. It acts as a base towards an acid, or as a ligand towards a metal, donating one, two or three electron pairs to metal centers. The hydroxide ion is present in the hydroxides of group 1 and group 2 elements, in the solid state. Those hydroxides dissociate in aqueous solution, liberating solvated hydroxide ions. Sodium hydroxide is a multi-million-ton per annum commodity chemical. A hydroxide group attached to a strongly electropositive center may itself dissociate, liberating a hydrogen cation (H⁺), making the parent compound an acid.

In organic chemistry, the hydroxide ion can act as a catalyst or as a nucleophilic reagent. An OH group, known as an hydroxyl group, is present in alcohols, phenols, carboxylic acids and related organic compounds.

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 :K_w = [H⁺][OH⁻] 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) which 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 [Al(OH)₄]^–. It is also often found in mixed-ligand complexes of the type [ML_x(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 [PtMe₃OH]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 [HCrO₄]^–, dissociates according to :[O₃CrO-H]^– [CrO₄]^2– + H⁺ with a pK_a of about 5.9.

Vibrational spectra

The infrared spectra of compounds containing the OH group have strong absorption bands in the region centered around 3500 cm⁻¹. The high frequency 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 an (OH)Zn(OH) bending vibration at 300 cm⁻¹.

Applications

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 :CO₃^2– + H₂O HCO₃^– + OH^–; (pK_a2 = 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 hydroxides.

Inorganic hydroxides

Alkali metals

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 weak base, with pK_b of 0.2. Lithium hydroxide is used in breathing gas purification systems for spacecraft, submarines, and rebreathers to remove carbon dioxide from exhaled gas. :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 strong bases.

Alkaline earth metals

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 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 lime water and can be used to test for the weak acid carbon dioxide. The reaction Ca(OH)₂ + CO₂ 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

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 solution. :Al(OH)₃ (solid) + OH^- (aq) [Al(OH)₄]^- (aq) 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 is unstable in aqueous solution: :C(OH)₄ → HCO₃^– + H₃O⁺ :HCO₃^– + 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 [SiO_x(OH)_4-2x]_n. ''Orthosilicic acid'' have been identified in very dilute aqueous solution. It is a weak acid with pK_a1 = 9.84, pK_a2 = 13.2 at 25 °C. It is usually written as H₄SiO₄ but the formula SiO₂(OH)₂ is generally accepted . 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 acid 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 hydroxide ion.

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

{|class="wikitable" style="text-align:center" | | | | | |- |Phosphorous acid |Phosphoric acid |Sulfuric acid |Telluric acid |''ortho''periodic 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 which can dissociate.

Telluric acid is often written with the formula H₂TeO₄·2H₂O but is better described structurally as Te(OH)₆.

''Ortho''-periodic acid can loose 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, particularly in salts such as the ''meso''periodate 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 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.

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. Hydroxides of metals in the +1 oxidation state are also poorly defned or unstable. For example, silver hydroxide, Ag(OH), decomposes spontaneously to the oxide (Ag₂O). Copper(I) and gold(I) hydroxides are similarly unstable, though stable adducts of CuOH and AuOH are known. The polymeric compounds M(OH)₂ and M(OH)₃ are generally prepared by increasing the pH of an aqueous solutions of the corresponding metal cations until the hydroxide precipitates out of solution. Conversely, the hydroxides dissolve in acidic solution. Zinc hydroxide, Zn(OH)₂, is amphoteric, forming the zincate ion, Zn(OH)₄^2– in strongly alkaline solution.

Numerous mixed ligand complexes of these metals with the hydroxide ion exist, in fact these are generally better defined than the simpler derivatives. Many can be made by causing dissociation of a co-ordinated water molecule. :L_nM(OH₂) + B L_nM(OH) + BH⁺ (L = ligand, B = base)

Vanadic acid, H₃VO₄, shows similarities with phosphoric acid, H₃PO₄, though it has a much more complex 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 conceptually 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. 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 poisoning.

Structural chemistry

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 can not 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 ions.

The amphoteric hydroxide Al(OH)₃ has four major crystalline forms: gibbsite (most stable), bayerite, nordstrandite and doyleite. 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 αAlO(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 A₃M^III(OH)₆, A₂M^IV(OH)₆ and AM^V(OH)₆. As the formula suggests these substances contain M(OH)₆ octahedral structural units. Layered double hydroxides may be represented by the formula [M^z+_1–xM³⁺_x(OH)₂]^q+(X^n–)_q/n·''y''H₂O. Most commonly, z = 2, and M^{2+ = 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.

Base catalysis

The hydroxide ion may act as a base catalyst. The base abstracts a proton from a weak acid to give an intermediate which goes on to react with another reagent. Common substrates for proton abstraction are alcohols, phenols, amines and carbon acids. The pK_a value for dissociation of a C-H bond is extremely high, but the pK_a alpha hydrogens of a carbonyl compound are about 3 log units lower. Typical pK_a 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 pK_a 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 which can be catalyzed by the presence of hydroxide. Hydroxide can also act as a Lewis-base catalyst.

As a nucleophilic reagent

The hydroxide ion is intermediate in nucleophilicity between the fluoride ion, F⁻, and the amide ion, NH₂⁻. 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 common ion effect.

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 dichlorocarbene.

Hydroxyl groups in organic compounds

Organic compounds such as alcohols, phenols and carboxylic acids contain hydroxyl groups. Each class of compound undergoes reactions specific to that class.

See also

Hydroxyl radical

Notes

References

Bibliography

Category:Bases Category:Oxoanions Category:Water chemistry

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