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Hydroxide
Lewis structure of the hydroxide ion showing three lone pairs on the oxygen atom
Space-filling representation of the hydroxide ion
Ball-and-stick model of the hydroxide ion
Names
IUPAC name
Hydroxide
Systematic IUPAC name
Oxidanide (not recommended)
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
UNII
  • InChI=1S/H2O/h1H2/p-1
  • [OH-]
Properties
OH−
Molar mass 17.007 g·mol−1
Basicity (pKb) 0.0 [1]
Conjugate acid Water
Conjugate base Oxide anion
Related compounds
Related compounds
 O2H+
OH•
O22−
H2O
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Hydroxide is a diatomic anion with chemical formula OH−. It consists of an oxygen and hydrogen atom held together by a single 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. The corresponding electrically neutral compound HO• is the hydroxyl radical. The corresponding covalently bound group −OH of atoms is the hydroxy group. Both the hydroxide ion and hydroxy group are nucleophiles and can act as catalysts 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 hydroxy groups.

Hydroxide ion

[edit]

The hydroxide ion is naturally produced from water by the self-ionization reaction:[2]

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, 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 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 various buffer solutions.

Schematic representation of the bihydroxide ion[3]

In an aqueous solution[4] 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 ions and hydroxide ions are strongly solvated, with hydrogen bonds between oxygen and hydrogen atoms. Indeed, the bihydroxide ion H
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 the bifluoride ion HF−
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 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, which acts as a lewis acid, to form, initially, the bicarbonate ion.

OH− + CO2 ⇌ HCO−
3

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 airtight plastic containers.

The hydroxide ion can function as a typical electron-pair donor ligand, 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 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.[5]

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.[6]

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.[7] 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).[8] 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.[9]

Applications

[edit]

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.[10] The principal method of manufacture is the chloralkali 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

CO2−
3 + H2O ⇌ HCO−
3 + OH−       (pKa2 = 10.33 at 25 °C and zero ionic strength)

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.[11] 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.[12]

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 (LiOH) is used in breathing gas purification systems for spacecraft, submarines, and rebreathers to remove carbon dioxide from exhaled gas.[13]

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.[14]

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.[15] 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.[16] At very low pH the aqua 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 increasing atomic 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, and barium hydroxide. A solution or 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

[edit]
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.[18]

Tetrahydroxo-
aluminate(III) ion

Aluminium hydroxide Al(OH)3 is amphoteric and dissolves in alkaline solution.[15]

Al(OH)3 (solid) + OH− (aq) ⇌ Al(OH)−
4 (aq)

In the Bayer process[19] 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/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 forms boehmite or diaspore, depending on crystal structure. Gallium hydroxide,[15] indium hydroxide, and thallium(III) hydroxide are also amphoteric. Thallium(I) hydroxide is a strong base.[20]

Carbon group elements

[edit]

Carbon forms no simple hydroxides. The hypothetical compound C(OH)4 (orthocarbonic acid or methanetetrol) is unstable in aqueous solution:[21]

C(OH)4 → HCO−
3 + H3O+
HCO−
3 + H+ ⇌ 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 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 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]

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−.[27]

Other main-group elements

[edit]
Phosphorous acid Phosphoric acid Sulfuric acid Telluric acid Orthoperiodic acid Xenic acid

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 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 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 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.[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 solution of the corresponding metal cation 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 tetrahydroxidozincate ion Zn(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 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.[35]

Basic salts containing hydroxide

[edit]

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.[36]

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.[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 hydroxyphosphate (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 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 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]

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.[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 of hydrogen bonding since the two hydroxide ions 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.[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+; hence q = x.

Organic reactions

[edit]

Potassium hydroxide and sodium hydroxide are two well-known reagents in organic chemistry.

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 ⇌ 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 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.[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]
Nucleophilic acyl substitution with an anionic nucleophile (Nu−) and leaving group (L−)

The hydroxide ion is intermediate in nucleophilicity between the fluoride ion F−, and the amide ion NH−
2.[50] Ester hydrolysis under alkaline conditions (also known as base hydrolysis)

R1C(O)OR2 + OH− ⇌ R1CO(O)H + −OR2 ⇌ R1CO2− + HOR2

is an example of a hydroxide ion serving as a nucleophile.[51]

Early methods for manufacturing soap treated triglycerides from animal fat (the ester) with lye.

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 the reactive intermediate dichlorocarbene.

Notes

[edit]

References

[edit]

Bibliography

[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Hydroxide

View on Grokipedia
from Grokipedia
The hydroxide ion, denoted as OH⁻, is a diatomic inorganic anion consisting of a single oxygen atom covalently bonded to a , with the negative charge primarily localized on the oxygen. It serves as the conjugate base of (H₂O) and is a fundamental species in aqueous chemistry, produced when bases dissociate in solution according to the Arrhenius definition, where its presence elevates the above 7. With the molecular formula OH⁻ and a of 17.007 g/mol, the hydroxide ion exhibits strong basicity and nucleophilicity due to the high on oxygen, enabling it to accept protons or participate in substitution reactions. Its preferred IUPAC name is (systematic name: oxidanide), and it forms stable ionic compounds known as hydroxides (e.g., , NaOH) that are widely used in industry for adjustment, cleaning, and synthesis. In biological systems, trace amounts function as metabolites, while in environmental contexts, it influences in natural s. The 's anomalous diffusion in water, faster than other small s, arises from a proton-hopping mechanism involving hydrogen-bonded networks.

Hydroxide Ion

Properties

The hydroxide (OH⁻) is a monovalent diatomic anion composed of an oxygen atom covalently bonded to a , carrying a negative charge primarily on the oxygen. It possesses an effective of approximately 133 pm when in a coordination environment of six, consistent with its role in ionic compounds and solvation shells. As the conjugate base of , OH⁻ demonstrates strong basicity in aqueous environments, with a pK_b \approx -1.7, reflecting its high affinity for protons and ability to deprotonate weak acids effectively. In , OH⁻ accepts protons to form H₂O, driving acid-base equilibria toward neutralization in solutions where [OH⁻] exceeds 10⁻⁷ M, resulting in > 7. This basic strength positions OH⁻ as a key in alkaline conditions, though its reactivity is moderated by effects. In aqueous solutions, the hydroxide ion undergoes hydration, forming solvated clusters denoted as [OH(H₂O)_n]⁻, where n ≈ 3–5 constitutes the primary hydration shell through hydrogen bonding from molecules donating protons to the oxygen of OH⁻. These clusters stabilize the and influence its mobility and reactivity, contributing to the elevated in basic media. The equilibrium concentration of OH⁻ is linked to H⁺ via the ion product of , defined as Kw=[\ceH+][\ceOH−]=1.0×10−14K_w = [\ce{H+}] [\ce{OH-}] = 1.0 \times 10^{-14} at 25°C under standard conditions. This constant arises from the autoionization of (2H₂O ⇌ H₃O⁺ + OH⁻) and shows temperature dependence, increasing to approximately 5.5 × 10^{-14} at 50°C due to the endothermic of the process, thereby shifting neutral below 7 at higher s. The properties of OH⁻ are evident in its production during the cathodic reduction of oxygen in alkaline solutions, characterized by the \ceO2+2H2O+4e−−>4OH−\ce{O2 + 2H2O + 4e- -> 4OH-} with a E∘=0.401E^\circ = 0.401 V versus the at 25°C and 14. This potential underscores OH⁻'s role in electrochemical processes like and fuel cells, where it serves as a product in oxygen reduction reactions.

Spectroscopic Characteristics

The hydroxide ion (OH⁻) exhibits distinct vibrational signatures that facilitate its detection across different phases. In the gas phase, the free OH⁻ ion displays a sharp O-H stretching vibration in the (IR) spectrum at approximately 3555 cm⁻¹, corresponding to the fundamental mode of the ionic bond. In aqueous or hydrogen-bonded environments, this band broadens and shifts to lower wavenumbers, typically centering around 3000–3600 cm⁻¹ with a continuum extending down to 800 cm⁻¹ due to delocalized proton transfer and effects. complements IR by revealing symmetric stretching modes; for instance, in concentrated aqueous NaOH solutions, a broad OH⁻ stretching band appears near 3500 cm⁻¹, while in solid hydroxides like NaOH, it sharpens to 3633 cm⁻¹. Electronic of OH⁻ reveals (UV) absorption primarily below 200 nm. A characteristic band around 190 nm arises from an n→σ* transition involving the oxygen to the O-H antibonding orbital, observable in both gas-phase and aqueous measurements of hydroxide solutions. This transition is intense and shifts slightly in solvated systems due to environmental polarization, but remains a key identifier for the isolated ion. Nuclear magnetic resonance (NMR) spectroscopy provides insights into the electronic environment of OH⁻ via ¹⁷O enrichment. In aqueous solutions, such as NaOH, the ¹⁷O for OH⁻ is approximately 20 ppm relative to , reflecting its high and minimal deshielding. The one-bond scalar ¹J(¹⁷O,¹H) is around 82 Hz, arising from the strong O-H bond and observable in high-resolution spectra despite quadrupolar broadening. In mass spectrometry, particularly electrospray ionization (ESI-MS), the OH⁻ ion is directly detected as the base peak at m/z 17 in negative-ion mode from hydroxide-containing samples. Fragmentation patterns of larger solvated or complexed species often yield this m/z 17 ion via neutral loss of water or other ligands, confirming the presence of intact OH⁻ without extensive dissociation due to ESI's soft ionization. Isotopic substitution highlights mass-dependent effects in OH⁻ spectra. Replacing ¹⁶O with ¹⁸O reduces the O-H stretching frequency by about 5.8%, from ~3555 cm⁻¹ to ~3350 cm⁻¹ in the gas phase, due to the increased reduced mass altering the vibrational potential. Similar downshifts occur in IR and Raman bands for ¹⁸OH⁻ in clusters or solutions, enabling precise quantification of isotopic ratios and aiding structural assignments.

Structural Chemistry

Bonding and Geometry

The hydroxide (OH⁻) is a diatomic anion characterized by a linear and an O–H of 0.964 , as determined from high-resolution spectroscopic measurements. In , the valence electrons occupy a σ orbital formed primarily from the oxygen 2s and 2p orbitals interacting with the 1s orbital, resulting in a of 1 and a closed-shell electronic configuration with eight valence electrons. This contrasts with the neutral OH radical, which has seven valence electrons, placing an in the σ orbital and yielding a slightly longer of 0.971 while maintaining a linear . and (DFT) calculations confirm the electron density distribution, with significant charge accumulation on the oxygen atom, contributing to the ion's high basicity. In coordination chemistry, the hydroxide ion serves as the hydroxo (OH⁻) ligand, typically binding through the oxygen atom in a monodentate fashion or as a bridging μ-OH group between metal centers. Monodentate coordination is exemplified in octahedral hexahydroxo complexes such as [Al(OH)₆]³⁻, where six OH⁻ ligands surround the central aluminum ion, forming regular M–O bonds with M–O–H angles approaching 110° due to the sp³-like hybridization at oxygen. Bridging μ-OH ligands are common in dinuclear and polynuclear complexes, such as [(H₂O)₄Cr(μ-OH)₂Cr(OH₂)₄]⁴⁺, where the OH⁻ group donates to two metal atoms simultaneously, often stabilizing higher oxidation states through interactions. These coordination modes influence the overall , with terminal hydroxo ligands leading to bent M–O–H arrangements and bridging forms adopting linear or asymmetric configurations depending on the metal–metal distance. The hydroxide ion plays a central role in hydrogen bonding networks, acting both as a hydrogen bond donor via its O–H group and as a strong acceptor through the negatively charged oxygen lone pairs. In aqueous solutions, OH⁻ forms extensive three-dimensional networks with molecules, typically engaging in three to four acceptor interactions and one donor interaction per ion, which contributes to its anomalously high mobility despite strong . The energy of individual O–H···O hydrogen bonds involving OH⁻ ranges from 20 to 40 kJ/mol, reflecting moderate to strong interactions that stabilize ionic structures and influence properties like in alkaline solutions. Quantum chemical calculations, including DFT and coupled-cluster methods, provide insights into the electronic structure of OH⁻. These computations also map the , showing a pronounced with partial negative charge on oxygen (~ -1.2 e) and positive on hydrogen (~ +0.2 e), consistent with the ion's role in proton transfer processes. In acidic conditions, coordinated OH⁻ ligands exhibit amphoteric behavior, readily undergoing to form aqua ligands (M–OH + H⁺ → M–OH₂⁺), which initiates reactions in metal complexes.

Crystal Structures

Hydroxide-containing materials often exhibit layered crystal structures, where metal cations are coordinated by hydroxide ions in octahedral arrangements, forming sheets that are stacked via hydrogen bonding. A prototypical example is , Mg(OH)₂, which adopts a trigonal structure in the P-3m1, consisting of brucite-type layers of edge-sharing MgO₆ octahedra. These layers are held together by weak interlayer hydrogen bonds between the hydroxyl groups of adjacent sheets, resulting in a highly anisotropic structure with cleavage parallel to the layers. Polymorphism is prevalent among hydroxide structures, particularly in aluminum hydroxides, where different stacking arrangements and coordination environments lead to distinct phases. , the α-form of Al(OH)₃, features a monoclinic with double layers of Al(OH)₆ octahedra linked by bonds, forming a dense packing that makes it the most thermodynamically stable polymorph under ambient conditions. In contrast, bayerite (β-Al(OH)₃) has a similar layered motif but with a different interlayer -bonding configuration, leading to a hexagonal , while (γ-AlOOH) represents an oxyhydroxide polymorph with orthorhombic and chains of edge-sharing AlO₆ octahedra connected via bonds. These polymorphs arise from variations in synthesis conditions, such as and , influencing their relative stability and applications in . Ionic hydroxides display diverse lattice types depending on the cation size and charge. , NaOH, forms an orthorhombic in the Cmcm at , with Na⁺ ions surrounded by OH⁻ ions in a distorted octahedral coordination, transitioning to a cubic phase at elevated temperatures around 575 K where the OH⁻ ions exhibit rotational disorder. In comparison, , Ca(OH)₂, known as , adopts a hexagonal in the P-3m1, isostructural with , featuring layers of CaO₆ octahedra linked by hydrogen bonds between hydroxyl groups pointing toward interlayer spaces. These structural differences reflect the larger of Ca²⁺ compared to Na⁺, promoting layered over cubic packing. Hydrogen bonding networks are crucial for stabilizing hydroxide crystal architectures, often forming infinite chains within layers or three-dimensional frameworks across the structure. In layered hydroxides like and , these networks involve O-H···O bonds with typical donor-acceptor O···O distances ranging from 2.7 to 3.0 , which dictate interlayer cohesion and facilitate properties such as swelling or . Such bonds create a balance between covalent intralayer interactions and weaker intermolecular forces, enabling polymorphism and influencing vibrational spectra. Defects and doping in hydroxide layers significantly alter their electronic and transport properties, particularly in (LDHs) derived from brucite-like structures. Vacancies, such as cation or anion defects, introduce local charge imbalances that enhance ionic conductivity by creating pathways for proton or hydroxide migration, as seen in NiFe-LDHs where oxygen vacancies improve rates. Doping with aliovalent ions, like partial substitution of Mg²⁺ with Al³⁺ in LDHs, generates positive layer charges balanced by interlayer anions, further tuning defect densities and boosting applications in electrocatalysis. These modifications, often introduced via synthesis or post-treatment, exemplify how structural imperfections can optimize material performance without disrupting the overall lattice.

Inorganic Hydroxides

Alkali and Alkaline Earth Metal Hydroxides

Alkali metal hydroxides, such as those of , sodium, and , are typically synthesized by the direct reaction of the alkali metal with , producing the hydroxide and gas; for example, sodium reacts as 2Na + 2H₂O → 2NaOH + H₂, a highly that generates heat and often ignites the . Alkaline earth metal hydroxides are commonly prepared by the hydration of the corresponding oxides, known as slaking; , for instance, reacts with to form : CaO + H₂O → Ca(OH)₂. Solubility in increases down , with exhibiting lower (12.8 g/100 mL at 20°C) compared to (109 g/100 mL at 20°C) and (112 g/100 mL at 20°C), primarily due to the higher of LiOH arising from the small size of the Li⁺ ion, which makes dissociation less favorable despite similar hydration energies. In , increases with increasing cation size, but hydroxides like Ca(OH)₂ remain sparingly soluble overall, with a solubility product constant (Ksp) of 5.5 × 10−6 at 25°C, reflecting its limited dissociation into Ca²⁺ and OH⁻ ions. These hydroxides display varying thermal stability, with NaOH remaining stable up to its around 323°C before decomposing into Na₂O and H₂O at higher temperatures under certain conditions. In contrast, Mg(OH)₂ decomposes at approximately 350°C to form MgO and : Mg(OH)₂ → MgO + H₂O. Sodium and hydroxides are highly hygroscopic and deliquescent, readily absorbing atmospheric moisture to form hydrates such as NaOH·H₂O, which crystallizes from aqueous solutions in specific temperature ranges. A notable application involves in limewater, a saturated used for detecting ; the reaction Ca(OH)₂ + CO₂ → CaCO₃ + H₂O produces an insoluble white precipitate of , turning the clear solution milky. This property underscores the strong basicity and reactivity of these s-block hydroxides in aqueous environments.

Transition and Post-Transition Metal Hydroxides

Transition and hydroxides exhibit diverse structures, ranging from amorphous gels to crystalline layered materials, and display variable reactivities influenced by the metal's d-electron configuration. These compounds are typically insoluble in , forming gelatinous precipitates that play key roles in qualitative inorganic and . Unlike the highly soluble s-block hydroxides, d-block and post-d-block variants often show amphoteric character, allowing dissolution in either acidic or basic conditions, and can adopt multiple oxidation states leading to mixed hydroxide phases. Precipitation reactions are central to identifying these metals in solution, particularly through the addition of hydroxide ions to form characteristic insoluble hydroxides. For instance, in qualitative analysis schemes, Fe³⁺ ions as red-brown Fe(OH)₃ upon reaction with OH⁻, as seen in group III cation separations where the hydroxide's color and insolubility distinguish iron from other metals. This reaction, Fe³⁺ + 3OH⁻ → Fe(OH)₃ (red-brown ppt), exemplifies the low solubility products (Ksp ≈ 10⁻³⁸) typical of trivalent hydroxides, enabling selective isolation in analytical procedures. Similar precipitations occur for other ions, such as Cr³⁺ forming green Cr(OH)₃, aiding in systematic metal identification. Amphoteric behavior is prominent in hydroxides of post-transition metals like aluminum and transition metals like zinc, where the precipitates redissolve in excess base to form soluble hydroxy complexes. Aluminum hydroxide, Al(OH)₃, initially forms a white gelatinous precipitate but dissolves in strong base via Al(OH)₃ + OH⁻ → [Al(OH)₄]⁻, demonstrating its ability to act as a Lewis acid by accepting additional hydroxide ligands. This property arises from the borderline acidic character of Al³⁺, with the tetrahydroxoaluminate ion stable in alkaline media (pH > 13). Zinc hydroxide exhibits analogous amphoterism, precipitating as white Zn(OH)₂ before dissolving in excess OH⁻ to yield [Zn(OH)₄]²⁻, a process driven by the coordination chemistry of Zn²⁺ in tetrahedral geometry. These reactions highlight the pH-dependent solubility of such hydroxides, contrasting with purely basic behavior in other metals. Variable oxidation states in transition metals lead to a range of hydroxide phases with distinct colors and stabilities. Manganese provides a classic example: , a pale pink precipitate from Mn²⁺ + 2OH⁻, represents the divalent state and is prone to aerial oxidation, while embodies the trivalent form, occurring as a black in hydrothermal deposits with a structure of edge-sharing Mn(III)O₆ octahedra distorted by Jahn-Teller effects. These phases underscore the versatility of , influencing formation and catalytic applications. Magnetic properties of these hydroxides often stem from unpaired d-electrons in the metal centers, conferring . Fe(OH)₃, with Fe(III) in a high-spin d⁵ configuration, exhibits due to five unpaired electrons, as confirmed in synthetic studies where the material shows susceptibility consistent with isolated Fe³⁺ sites before any antiferromagnetic ordering upon aging or . Similarly, Cr(OH)₃ displays from its Cr(III) d³ electrons (three unpaired), with magnetic moments around 3.8 μB, reflecting weak exchange interactions in the polymeric structure. These properties are exploited in nanomaterial design for . Coprecipitation and aging processes enable the formation of mixed hydroxides, notably (LDHs), which consist of brucite-like sheets of divalent and trivalent metals (e.g., Mg²⁺/Al³⁺ or Zn²⁺/Fe³⁺) with intercalated anions. Synthesized by of metal salts at constant pH (typically 8–10), LDHs undergo aging to crystallize into structures, where positive layer charge is balanced by anions like CO₃²⁻ or Cl⁻. The weak electrostatic binding allows facile anion exchange, enabling applications in remediation by swapping interlayer species for pollutants, with exchange capacities up to 3–4 meq/g depending on layer composition. This versatility arises from the tunable metal ratios and high surface area (50–200 m²/g) post-aging.

Hydroxides of p-Block Elements

The hydroxides of p-block elements, spanning Groups 13 through 16, exhibit predominantly covalent bonding due to the increasing across the block, leading to molecular or polymeric structures rather than ionic lattices. These compounds are often unstable in isolation, prone to , , or conversion to oxyanions, reflecting the tendency of p-block elements to form stable multiple bonds with oxygen. Unlike the more ionic metal hydroxides, p-block variants frequently display acidic character and volatility, with applications limited by their reactivity and in some cases. In Group 13, , B(OH)3, is a prototypical covalent hydroxide, existing as a trigonal planar where acts as a Lewis acid by accepting a to form the tetrahedral [B(OH)4]- in basic solution. It behaves as a very weak monoprotic acid with an Ka = 5.8 × 10−10 at 25°C, owing to the electron-deficient center that weakly polarizes an O–H bond in the hydrated form. Aluminum hydroxide, Al(OH)3, marks a transition toward more metallic behavior; it is amphoteric, dissolving in both acids and bases, though detailed and aspects overlap with hydroxides. Group 14 hydroxides highlight the distinction between true inorganic species and organic analogs. For carbon, compounds like methanol (CH3OH) incorporate an –OH group but function as alcohols with covalent C–O bonds, lacking the ionic OH- character of metal hydroxides and instead undergoing nucleophilic substitution or dehydration. Orthosilicic acid, Si(OH)4, is a tetrahedral monomer stable only in dilute aqueous solutions below approximately 100 mg/L SiO2; it readily undergoes condensation polymerization, releasing water to form siloxane (Si–O–Si) chains and networks that constitute silicates and silica gels. Tin hydroxides, such as Sn(OH)2 and Sn(OH)4, adopt polymeric or cluster structures for stability; for instance, Sn(II) species form cations like [Sn3(OH)4]2+ with bridging hydroxo groups, while Sn(IV) variants exhibit octahedral coordination in extended lattices, contributing to their limited solubility and tendency to precipitate as hydrous oxides. In Group 15, , H3PO3, features a structure with two ionizable –OH groups attached to phosphorus and a direct P–H bond, rendering it diprotic with pKa values of 2.00 and 6.59, distinct from the triprotic H3PO4. Arsenic trihydroxide, As(OH)3 (also known as arsenious acid, H3AsO3), is highly toxic, with acute oral exposure causing severe gastrointestinal distress, , and multiorgan failure at doses as low as 70–180 mg; chronic environmental contamination, particularly in , poses global health risks including carcinogenicity and skin lesions, affecting millions through in food chains. Group 16 elements form oxoacids rather than simple hydroxides; , H2SO4, is structured as (HO)2SO2 with two –OH groups but is classified as a strong diprotic oxoacid, not a hydroxide, serving as a key industrial source of ions via ionization. For , the selenate dianion SeO42- from selenic acid H2SeO4 can be conceptualized in hydrated basic media, though discrete Se(OH)62- species are not commonly isolated, unlike in analogous chemistry.

Reactions and Applications

Industrial and Laboratory Uses

Hydroxide compounds, particularly (NaOH), play a central role in various industrial processes. In soap production, NaOH facilitates , the reaction of fats or oils with alkali to form and , serving as a key raw material for both traditional and synthetic detergents. In the for paper manufacturing, NaOH, combined with , treats wood chips under high pressure to dissolve and separate fibers, enabling pulp production that accounts for the majority of global paper output. Additionally, NaOH is essential in the for alumina extraction, where it dissolves aluminum oxide from ore at elevated temperatures and pressures, forming from which pure alumina is subsequently precipitated. Calcium hydroxide, Ca(OH)₂, is widely employed in for adjustment, softening by precipitating calcium and magnesium ions as carbonates, and neutralizing acidic to prevent environmental harm. In battery technology, (KOH) acts as the in alkaline batteries, enabling the electrochemical reaction between and ; the overall process can be represented as: Zn+2MnO2+H2O→ZnO+2MnOOH\mathrm{Zn + 2MnO_2 + H_2O \rightarrow ZnO + 2MnOOH} This reaction provides high energy density and long shelf life, making KOH-based alkaline batteries a staple in consumer electronics. In laboratory settings, NaOH is a standard reagent for acid-base titrations, where it is used to standardize solutions or determine acid concentrations by reaching equivalence points with indicators like phenolphthalein. It is also integral to aqueous workups in organic extractions, where basic NaOH solutions deprotonate acidic compounds to partition them into the aqueous phase, facilitating purification of reaction mixtures. Magnesium hydroxide, Mg(OH)₂, functions as a non-halogenated in polymers and composites, undergoing endothermic above 300°C to release , which dilutes combustible gases and absorbs from the substrate. This mechanism enhances in applications such as electrical cables and construction materials without generating toxic byproducts.

Role in Acid-Base Chemistry

In acid-base chemistry, the hydroxide ion (OH⁻) plays a central role in defining basicity through its concentration in aqueous solutions. The pOH scale measures this concentration logarithmically as pOH = -log[OH⁻], providing a complementary metric to pH, which quantifies hydronium ion (H₃O⁺) activity. At 25°C, the relationship pH + pOH = 14 arises from the ion product of water (K_w = 1.0 × 10⁻¹⁴), establishing neutrality at pH 7 where [OH⁻] = [H₃O⁺] = 1.0 × 10⁻⁷ M. This framework allows chemists to assess solution basicity; for instance, a solution with [OH⁻] = 0.01 M has pOH = 2 and pH = 12, indicating strong basicity. Hydroxide ions are essential in acid-base titrations, particularly when a strong base like NaOH neutralizes a strong acid such as HCl. The reaction proceeds as NaOH + HCl → NaCl + H₂O, with the equivalence point occurring at 7 due to complete neutralization and formation of water under neutral conditions. In titration curves for strong acid-strong base systems, the rises gradually before the equivalence point, then sharply increases beyond it as excess OH⁻ dominates, enabling precise determination of analyte concentration through stoichiometric ratios. In buffering systems, OH⁻ interacts with weak acids to maintain stable . For example, in an acetate buffer comprising acetic acid (CH₃COOH) and its conjugate base (CH₃COO⁻ from ), added OH⁻ reacts with CH₃COOH to form CH₃COO⁻ + H₂O, shifting the equilibrium without significant pH change according to the Henderson-Hasselbalch . This resistance to pH variation is crucial for applications requiring constant acidity, such as biological assays, where the buffer's capacity depends on the weak acid's pK_a near the desired pH. Acid-base indicators like rely on OH⁻-induced for visual endpoint detection in titrations. undergoes a color change from colorless (protonated form, HIn) to pink (deprotonated form, In²⁻) over the range 8.2–10.0, as OH⁻ shifts the equilibrium HIn ⇌ H⁺ + In²⁻ toward the colored anion in basic media. This transition aligns well with equivalence points in weak acid-strong base titrations, providing a sharp visual cue for completion. For polyprotic acids, OH⁻ facilitates stepwise neutralization, allowing sequential of multiple acidic protons. (H₃PO₄), a triprotic acid, reacts progressively: H₃PO₄ + OH⁻ → H₂PO₄⁻ + H₂O (pK_{a1} ≈ 2.1), H₂PO₄⁻ + OH⁻ → HPO₄²⁻ + H₂O (pK_{a2} ≈ 7.2), and HPO₄²⁻ + OH⁻ → PO₄³⁻ + H₂O (pK_{a3} ≈ 12.7), with full neutralization requiring three equivalents of base to yield PO₄³⁻ + 3H₂O. curves exhibit distinct inflection points corresponding to each pK_a, enabling selective quantification of acid forms in mixtures like fertilizers or biological fluids.

Organic Chemistry of Hydroxide

Base-Catalyzed Reactions

In base-catalyzed reactions, the hydroxide ion (OH⁻) functions primarily as a Brønsted base by abstracting a proton from an organic substrate, thereby facilitating transformations such as eliminations, condensations, and isomerizations. These processes are fundamental in , where the strong basicity of OH⁻ in protic solvents like or enables the generation of reactive intermediates like carbanions or enolates. The mechanisms typically proceed via concerted or stepwise pathways, influenced by and substrate structure, and are widely studied for their role in both laboratory and industrial applications. A prominent example is the E2 elimination, a concerted bimolecular process where OH⁻ abstracts a β-proton from an while the departs, forming an . In , this is efficient for primary alkyl bromides, such as the conversion of ethyl bromide to ethylene: CH3CH2Br+OH−→CH2=CH2+Br−+H2O\text{CH}_3\text{CH}_2\text{Br} + \text{OH}^- \rightarrow \text{CH}_2=\text{CH}_2 + \text{Br}^- + \text{H}_2\text{O} The depends on composition, with ethanol-water mixtures enhancing the elimination over substitution due to reduced of OH⁻, leading to higher basicity. This mechanism requires anti-periplanar geometry for optimal orbital overlap and is favored under kinetic control with strong bases like OH⁻. In , OH⁻ catalyzes the self-addition of carbonyl compounds by deprotonating the α-carbon to form an intermediate, which then attacks another . For , the process begins with enolate formation: CH3CHO+OH−⇌−CH2CHO+H2O\text{CH}_3\text{CHO} + \text{OH}^- \rightleftharpoons ^-\text{CH}_2\text{CHO} + \text{H}_2\text{O} followed by to yield the β-hydroxy , and subsequent under basic conditions to the α,β-unsaturated carbonyl. The rate-limiting step often involves the loss of OH⁻ during dehydration, particularly in aqueous media, highlighting the role of OH⁻ in both initiation and termination. This reaction is versatile for C-C bond formation and exemplifies base-promoted chemistry. The of esters, known as , proceeds via a base-catalyzed mechanism where OH⁻ attacks the carbonyl carbon, forming a tetrahedral intermediate that expels the , yielding a and alcohol: RCOOR’+OH−→RCOO−+R’OH\text{RCOOR'} + \text{OH}^- \rightarrow \text{RCOO}^- + \text{R'OH} This addition-elimination pathway is second-order overall, first-order in both ester and OH⁻ concentrations, with the rate-determining step being the formation of the tetrahedral intermediate in aqueous solutions. Theoretical studies confirm multiple pathways, but the BAC2 mechanism dominates for simple alkyl esters, driven by the basicity of OH⁻. is industrially significant for production from fats. Hofmann elimination involves the thermal decomposition of quaternary ammonium hydroxides, where OH⁻ abstracts a β-proton in an E2-like manner, leading to an , tertiary , and . The general reaction is: R4N++OH−→alkene+R3N+H2O\text{R}_4\text{N}^+ + \text{OH}^- \rightarrow \text{alkene} + \text{R}_3\text{N} + \text{H}_2\text{O} This process favors the least substituted due to steric factors in the and may involve an intermediate for certain substrates, enhancing selectivity for terminal s. It is commonly used for exhaustive followed by elimination to determine structures. Base-induced isomerization of alkenes occurs through reversible proton abstraction by OH⁻, allowing double bond migration toward more stable conjugated or internal positions. For 1-butene, OH⁻ in alcoholic media catalyzes the shift to 2-butene via allylic carbanion intermediates, with solvent polarity influencing the equilibrium. This proton transfer mechanism is equilibrium-controlled and is applied in refining processes to optimize alkene stability.

Nucleophilic Reactions

Hydroxide ion (OH⁻) acts as a in by attacking electron-deficient centers, such as carbon atoms bearing good leaving groups or electrophilic carbonyl carbons, leading to substitution or products. This reactivity is particularly prominent in aqueous or alcoholic media, where OH⁻ is generated from bases like NaOH or KOH. Unlike its role as a base in reactions, here the focus is on direct bond formation via nucleophilic attack. In SN2 reactions, OH⁻ displaces leaving groups from primary or methyl alkyl halides through a concerted backside attack, resulting in inversion of at the carbon center. For example, the reaction of methyl iodide with hydroxide yields and iodide ion: CH3I+OH−→CH3OH+I−\text{CH}_3\text{I} + \text{OH}^- \rightarrow \text{CH}_3\text{OH} + \text{I}^- This mechanism is favored under basic conditions due to the strong nucleophilicity of OH⁻ and the low steric hindrance at primary carbons. The stereochemical inversion is a hallmark of the SN2 pathway, distinguishing it from SN1 processes that involve intermediates. Epoxide ring opening by OH⁻ proceeds via nucleophilic attack at the less substituted carbon under basic conditions, driven by the strain relief in the three-membered ring and steric accessibility. This contrasts with acid-catalyzed openings, where attack occurs at the more substituted site. A classic example is the of (oxirane) to : \ce(CH2)2O+OH−−>HOCH2CH2OH\ce{(CH2)2O + OH^- -> HOCH2CH2OH} The reaction is typically carried out in aqueous NaOH and is industrially important for glycol production. Nucleophilic acyl substitution with OH⁻ is a key hydrolysis pathway for activated carboxylic acid derivatives, where the hydroxide adds to the carbonyl carbon, forming a tetrahedral intermediate that expels the . Acid chlorides react rapidly with OH⁻ to form salts, which upon acidification yield s: RCOCl+OH−→RCOO−+Cl−\text{RCOCl} + \text{OH}^- \rightarrow \text{RCOO}^- + \text{Cl}^- This process is significantly faster for acid chlorides than for less reactive , where base hydrolysis requires harsher conditions like heating in concentrated NaOH to cleave the amide bond and produce the and . The reactivity order—acid chlorides > anhydrides > esters > —reflects the quality of the and the stability of the tetrahedral intermediate. OH⁻ also adds directly to carbonyl groups in certain aldehydes, particularly , forming gem-diols or their conjugate bases in non-catalyzed or base-promoted equilibria. For , the hydration involves OH⁻ addition to yield the hydroxymethoxide ion: HCHO+OH−⇌H2C(OH)O−\text{HCHO} + \text{OH}^- \rightleftharpoons \text{H}_2\text{C(OH)O}^- This equilibrium lies far toward the due to the lack of steric hindrance and electron-withdrawing effects in , unlike higher aldehydes where hydration is less favorable. Such additions are relevant in aqueous environments and prebiotic chemistry simulations. To extend OH⁻ nucleophilicity to non-aqueous media, phase-transfer catalysis employs lipophilic quaternary salts (e.g., ) to transport the anion from an aqueous base layer into an organic solvent. This enables efficient SN2 displacements, openings, or acyl substitutions in immiscible systems, enhancing reaction rates by increasing local OH⁻ concentration. The catalysts form pairs with OH⁻, solubilizing it without altering its nucleophilic character.

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