Manyinorganic substances which bear the wordhydroxide in their names are notionic compounds of the hydroxide ion, but covalent compounds which containhydroxy groups.
has a value close to 10−14 at 25 °C, so theconcentration of hydroxide ions in pure water is close to 10−7 mol∙dm−3, to satisfy the equal charge constraint. ThepH of a solution is equal to the decimalcologarithm of thehydrogen 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 ofpOH, which is close to (14 − pH),[note 3] so the 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 (decrease pH, increase 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 decreases the hydrogen cation concentration, which increases the hydroxide ion concentration. pOH can be kept at a nearly constant value with variousbuffer solutions.
Schematic representation of the bihydroxide ion[3]
In anaqueous solution[4] the hydroxide ion is abase in theBrønsted–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 aLewis base by donating a pair of electrons to a Lewis acid. In aqueous solution both hydrogen ions and hydroxide ions are stronglysolvated, withhydrogen bonds between oxygen and hydrogen atoms. Indeed, the bihydroxide ionH 3O− 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 thebifluoride ionHF− 2 (114 pm).[3] 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 highviscosity due to the formation of an extended network of hydrogen bonds as inhydrogen fluoride solutions.
In solution, exposed to air, the hydroxide ion reacts rapidly with atmosphericcarbon dioxide, which acts as a lewis acid, to form, initially, thebicarbonate ion.
OH− + CO2 ⇌HCO− 3
Theequilibrium constant for this reaction can be specified either as a reaction with dissolved carbon dioxide or as a reaction with carbon dioxide gas (seeCarbonic acid for values and details). At neutral or acid pH, the reaction is slow, but is catalyzed by theenzymecarbonic anhydrase, which effectively creates hydroxide ions at the active site.
Solutions containing the hydroxide ion attackglass. In this case, thesilicates in glass are acting as acids. Basic hydroxides, whether solids or in solution, are stored inairtight plastic containers.
The hydroxide ion can function as a typical electron-pair donorligand, forming such complexes as tetrahydroxoaluminate/tetrahydroxidoaluminate [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 abridging 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.[5]
When bound to a strongly electron-withdrawing metal centre, hydroxide ligands tend toionise into oxide ligands. For example, the bichromate ion [HCrO4]− dissociates according to
Theinfrared spectra of compounds containing the OHfunctional group have strongabsorption bands in the region centered around 3500 cm−1.[7] The high frequency ofmolecular 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 byinfrared spectroscopy relatively easy. A band due to an OH group tends to be sharp. However, theband 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 acoordination 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).[8] M−OH stretching vibrations occur below about 600 cm−1. For example, thetetrahedral 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.[9]
Solutions containing the hydroxide ion are generated when a salt of aweak acid is dissolved in water.Sodium carbonate is used as an alkali, for example, by virtue of thehydrolysis reaction
CO2− 3 + H2O ⇌HCO− 3 + OH− (pKa2 = 10.33 at 25 °C and zeroionic strength)
An example of the use of sodium carbonate as an alkali is whenwashing soda (another name for sodium carbonate) acts on insolubleesters, such astriglycerides, commonly known as fats, to hydrolyze them and make them soluble.
Bauxite, a basic hydroxide ofaluminium, is the principal ore from which the metal is manufactured.[11] Similarly,goethite (α-FeO(OH)) andlepidocrocite (γ-FeO(OH)), basic hydroxides ofiron, are among the principal ores used for the manufacture of metallic iron.[12]
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 isamphoteric.[15] The hydroxide itself isinsoluble in water, with asolubility product log K*sp of −11.7. Addition of acid gives solublehydrolysis 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.[16] At very low pH theaqua ion [Be(H2O)4]2+ is formed. Addition of hydroxide to Be(OH)2 gives the soluble tetrahydroxoberyllate or tetrahydroxidoberyllate anion, [Be(OH)4]2−.
The solubility in water of the other hydroxides in this group increases with increasingatomic number.[17]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, andbarium hydroxide. A solution or suspension of calcium hydroxide is known aslimewater and can be used to test for theweak 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.
Aluminium hydrolysis as a function of pH. Water molecules attached to Al are omitted
The simplest hydroxide of boron B(OH)3, known asboric 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.
In theBayer process[19] for the production of pure aluminium oxide frombauxite minerals this equilibrium is manipulated by careful control of temperature and alkali concentration. In the first phase, aluminium dissolves in hot alkaline solution asAl(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-calledred 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/hydroxido 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 formsboehmite ordiaspore, depending on crystal structure.Gallium hydroxide,[15]indium hydroxide, andthallium(III) hydroxide are also amphoteric.Thallium(I) hydroxide is a strong base.[20]
Carbon dioxide is also known as carbonic anhydride, meaning that it forms by dehydration ofcarbonic 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 can be written as H4SiO4 or Si(OH)4.[6] 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 ofgermanium have been characterized.Tin(II) hydroxide Sn(OH)2 was prepared in anhydrous media. Whentin(II) oxide is treated with alkali the pyramidal hydroxo complexSn(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 whichstannates, with a formula [Sn(OH)6]2−, are derived by reaction with the (Lewis) basic hydroxide ion.[26]
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 solubleplumbate ions are formed, including [Pb(OH)6]2−.[27]
In the higher oxidation states of thepnictogens,chalcogens,halogens, andnoble gases there are oxoacids in which the central atom is attached to oxide ions and hydroxide ions. Examples includephosphoric acid H3PO4, andsulfuric acid H2SO4. In these compounds one or more hydroxide groups candissociate with the liberation of hydrogen cations as in a standardBrønsted–Lowry acid. Many oxoacids of sulfur are known and all feature OH groups that can dissociate.[28]
Telluric acid is often written with the formula H2TeO4·2H2O but is better described structurally as Te(OH)6.[29]
Orthoperiodic 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 theisoelectronic 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 ofchlorine,bromine, andiodine have the formula On−1/2A(OH), wheren is theoxidation number: +1, +3, +5, or +7, and A = Cl, Br, or I. The only oxoacid offluorine is F(OH),hypofluorous acid. When these acids are neutralized the hydrogen atom is removed from the hydroxide group.[33]
The hydroxides of thetransition metals andpost-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 calledolation. 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 solution of the corresponding metal cation until the hydroxideprecipitates out of solution. On the converse, the hydroxides dissolve in acidic solution.Zinc hydroxide Zn(OH)2 is amphoteric, forming the tetrahydroxidozincate ionZn(OH)2− 4 in strongly alkaline solution.[15]
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 correspondingmetal aquo complex.
LnM(OH2) + B ⇌ LnM(OH)– + BH+ (L = ligand, B = base)
Vanadic acid H3VO4shows similarities with phosphoric acid H3PO4 though it has a much more complexvanadate oxoanion chemistry.Chromic acid H2CrO4, has similarities with sulfuric acid H2SO4; for example, both formacid 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 ofolation, formingpolyoxometalates.[35]
In some cases, the products of partial hydrolysis of metal ion, described above, can be found in crystalline compounds. A striking example is found withzirconium(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 atetrameric 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.[36]
The mineralmalachite is a typical example of a basic carbonate. The formula, Cu2CO3(OH)2 shows that it is halfway betweencopper carbonate andcopper hydroxide. Indeed, in the past the formula was written as CuCO3·Cu(OH)2. Thecrystal structure is made up of copper, carbonate and hydroxide ions.[36] The mineralatacamite 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 hydroxyphosphate (libethenite), arsenate (olivenite), sulfate (brochantite), and nitrate compounds.White lead is a basiclead carbonate, (PbCO3)2·Pb(OH)2, which has been used as a whitepigment because of its opaque quality, though its use is now restricted because it can be a source forlead poisoning.[36]
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 effectiveionic radius of about 153 pm.[38] Thus, the high-temperature forms of KOH and NaOH have thesodium chloride structure,[39] which gradually freezes in a monoclinically 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 byX-ray diffraction.[40] The room-temperature form of NaOH has thethallium 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 thecadmium iodide layer structure, with a kind of close-packing of magnesium and hydroxide ions.[38][41]
Theamphoteric hydroxide Al(OH)3 has four major crystalline forms:gibbsite (most stable),bayerite,nordstrandite, anddoyleite.[note 7]All thesepolymorphs 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.[42] 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.[43] This is an unusual form ofhydrogen bonding since the two hydroxide ions involved would be expected to point away from each other. The hydrogen atoms have been located byneutron 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.[44] A similar type of hydrogen bond has been proposed for other amphoteric hydroxides, including Be(OH)2, Zn(OH)2, and Fe(OH)3.[38]
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·yH 2O. Most commonly,z = 2, and M2+ = Ca2+, Mg2+, Mn2+, Fe2+, Co2+, Ni2+, Cu2+, or Zn2+; henceq = x.
The hydroxide ion may act as abase 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 arealcohols,phenols,amines, andcarbon acids. ThepKa value for dissociation of a C–H bond is extremely high, but the pKaalpha hydrogens of a carbonyl compound are about 3 log units lower. Typical pKa values are 16.7 foracetaldehyde and 19 foracetone.[47] 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 to one by adding sodium hydroxide toethanol
OH− + EtOH ⇌ EtO− + H2O
to produce theethoxide ion. ThepKa for self-dissociation of ethanol is about 16, so the alkoxide ion is a strong enough base.[48] The addition of an alcohol to analdehyde to form ahemiacetal 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]
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