Editing Hydroxide

{{Short description|Chemical compound}}
{{Chembox
| ImageFile1 = Hydroxide lone pairs-2D.svg
| ImageSize1 = 170px
| ImageClass1 = skin-invert-image
| ImageAlt1 = [[Lewis structure]] of the hydroxide ion showing three [[lone pair]]s on the oxygen atom
| ImageFileL1 = Hydroxide-3D-vdW.png
| ImageClassL1 = bg-transparent
| ImageAltL1 = Space-filling representation of the hydroxide ion
| ImageFileR1 = Hydroxide-3D-balls.png
| ImageClassR1 = bg-transparent
| ImageAltR1 = [[Ball-and-stick model]] of the hydroxide ion
|IUPACName = Hydroxide
| SystematicName = Oxidanide ''(not recommended)''
| OtherNames =  
| Section1 = {{Chembox Identifiers
| CASNo = 14280-30-9
| PubChem = 961
| ChEBI =16234
| SMILES = [OH-]
| UNII = 9159UV381P
| InChI = 1S/H2O/h1H2/p-1
| ChemSpiderID = 936
}}
| Section2 = {{Chembox Properties
| Formula = {{chem|OH|−}}
| O=1|H=1
| Appearance =
| Solubility =
| pKb = 0.0 <ref>{{cite journal|doi=10.1002/hlca.202400103|title=The pKa of Water and the Fundamental Laws Describing Solution Equilibria: An Appeal for a Consistent Thermodynamic Pedagogy|author=Neils, T.L.; Schaertel, S. and Silverstein, T.P.|journal=Helv. Chim. Acta|year=2024|volume=107|issue=11|doi-access=free}}</ref>
| ConjugateAcid = [[Properties of water|Water]]
| ConjugateBase = [[Oxide anion]]}}
| Section3 = {{Chembox Hazards
| MainHazards =
| FlashPt =
| AutoignitionPt =}}
|Section8={{Chembox Related
|OtherCompounds = [[Dioxidanylium|O<sub>2</sub>H<sup>+</sup>]]<br/>[[Hydroxyl radical|OH<sup>•</sup>]]<br/>[[Peroxide|O<sub>2</sub><sup>2−</sup>]]<br/>[[water|H<sub>2</sub>O]]
 }}
}}

'''Hydroxide''' is a [[polyatomic ion|diatomic anion]] with [[chemical formula]] OH<sup>−</sup>. 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 [[Self-ionization of water|minor constituent of water]]. It functions as a [[base (chemistry)|base]], a [[ligand]], a [[nucleophile]], and a [[catalyst]]. The hydroxide ion forms [[salt (chemistry)|salts]], some of which [[dissociation (chemistry)|dissociate]] in aqueous solution, liberating solvated hydroxide ions. [[Sodium hydroxide]] is a multi-million-ton per annum [[commodity chemicals|commodity chemical]].
The corresponding [[electrically neutral]] compound&nbsp;HO<sup>•</sup> is the [[hydroxyl radical]]<!-- explain further? -->. The corresponding [[covalent bond|covalently]] bound [[functional group|group]]&nbsp;{{chem2|\sOH}} of atoms is the [[hydroxy group]].
Both the hydroxide ion and hydroxy group are [[nucleophile]]s and can act as catalysts in [[organic chemistry]].

Many [[inorganic]] substances which bear the word ''hydroxide'' in their names are not [[ionic compound]]s of the hydroxide ion, but covalent compounds which contain [[hydroxy group]]s.

==Hydroxide ion==
The hydroxide ion is naturally produced from [[water]] by the [[self-ionization of water|self-ionization]] reaction:<ref>{{cite journal|author1= Geissler, P. L.|author2-link=Christoph Dellago|author2= Dellago, C.|author3= Chandler, D.|author4= Hutter, J.|author5= Parrinello, M.|year= 2001|title= Autoionization in liquid water|journal= [[Science (journal)|Science]]|volume= 291|pages= 2121–2124|doi= 10.1126/science.1056991|pmid= 11251111|issue= 5511|bibcode= 2001Sci...291.2121G|url= http://gold.cchem.berkeley.edu:8080/Pubs/DC174.pdf|citeseerx= 10.1.1.6.4964|s2cid= 1081091|access-date= 2017-10-25|archive-url= https://web.archive.org/web/20070625233942/http://gold.cchem.berkeley.edu:8080/Pubs/DC174.pdf|archive-date= 2007-06-25|url-status= dead}}</ref>
:[[Hydronium|H<sub>3</sub>O<sup>+</sup>]] + OH<sup>−</sup> {{eqm}} 2H<sub>2</sub>O
The [[equilibrium constant]] for this reaction, defined as
:''K''<sub>w</sub> = [H<sup>+</sup>][OH<sup>−</sup>]<ref group=note>[H<sup>+</sup>] denotes the concentration of [[hydron (chemistry)|hydrogen cations]] and [OH<sup>−</sup>] the concentration of hydroxide ions</ref>
has a value close to 10<sup>−14</sup> at 25&nbsp;°C, so the [[concentration]] of hydroxide ions in pure water is close to 10<sup>−7</sup>&nbsp;mol∙dm<sup>−3</sup>, to satisfy the equal charge constraint. The [[pH]] of a solution is equal to the decimal [[cologarithm]] of the [[hydron (chemistry)|hydrogen cation]] concentration;<ref group=note>Strictly speaking pH is the cologarithm of the hydrogen cation [[activity (chemistry)|activity]]</ref> the pH of pure water is close to 7 at ambient temperatures. The concentration of hydroxide ions can be expressed in terms of [[pH#pOH|pOH]], which is close to (14&nbsp;−&nbsp;pH),<ref group=note>pOH signifies the negative logarithm to base 10 of [OH<sup>−</sup>], alternatively the logarithm of {{sfrac|1|[OH<sup>−</sup>]<nowiki/>}}</ref> 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 NH<sub>3</sub> + H<sup>+</sup> {{eqm}} {{chem|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 solution]]s.

[[File:Bihydoxide.png|thumb|150px|Schematic representation of the bihydroxide ion<ref name=ARF/>]]

In an [[aqueous solution]]<ref>{{cite journal|last=Marx|first=D.|author2=Chandra, A |author3=Tuckerman, M.E. |year=2010|title=Aqueous Basic Solutions: Hydroxide Solvation, Structural Diffusion, and Comparison to the Hydrated Proton|journal=Chem. Rev.|volume=110|issue=4|pages=2174–2216|doi=10.1021/cr900233f|pmid=20170203}}</ref> the hydroxide ion is a [[base (chemistry)|base]] in the [[Brønsted–Lowry acid–base theory|Brønsted–Lowry]] sense as it can accept a proton<ref group=note>In this context proton is the term used for a solvated hydrogen cation</ref> 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 bond]]s between oxygen and hydrogen atoms. Indeed, the bihydroxide ion {{chem|H|3|O|2|−}} has been characterized in the solid state. This compound is centrosymmetric and has a very short hydrogen bond (114.5&nbsp;[[picometre|pm]]) that is similar to the length in the [[bifluoride]] ion {{chem|HF|2|−}} (114&nbsp;pm).<ref name=ARF>{{cite journal|title= The bihydroxide ({{chem|H|3|O|2|−}}) anion. A very short, symmetric hydrogen bond|author1=Kamal Abu-Dari |author2=Kenneth N. Raymond |author3=Derek P. Freyberg |journal= [[J. Am. Chem. Soc.]]|year= 1979|volume= 101|pages= 3688–3689|doi= 10.1021/ja00507a059|issue= 13}}</ref> 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<sup>−</sup> + CO<sub>2</sub> {{eqm}} {{chem|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 [[silicate]]s 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/tetrahydroxido[[aluminate]] [Al(OH)<sub>4</sub>]<sup>−</sup>. It is also often found in mixed-ligand complexes of the type [ML<sub>''x''</sub>(OH)<sub>''y''</sub>]<sup>''z''+</sup>, 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 [Pb<sub>2</sub>(OH)]<sup>3+</sup>, metal hydroxides are often written in a simplified format. It can even act as a 3-electron-pair donor, as in the tetramer [PtMe<sub>3</sub>(OH)]<sub>4</sub>.<ref>Greenwood, p.&nbsp;1168</ref>

When bound to a strongly electron-withdrawing metal centre, hydroxide ligands tend to [[dissociation (chemistry)|ionise]] into oxide ligands. For example, the bichromate ion [HCrO<sub>4</sub>]<sup>−</sup> dissociates according to
:[O<sub>3</sub>CrO–H]<sup>−</sup> {{eqm}} [CrO<sub>4</sub>]<sup>2−</sup> + H<sup>+</sup>
with a p''K''<sub>a</sub> of about 5.9.<ref name=scdb>[http://www.acadsoft.co.uk/scdbase/scdbase.htm IUPAC SC-Database] {{Webarchive|url=https://web.archive.org/web/20170619235720/http://www.acadsoft.co.uk/scdbase/scdbase.htm |date=2017-06-19 }} A comprehensive database of published data on equilibrium constants of metal complexes and ligands</ref>

===Vibrational spectra===
The [[infrared spectrum|infrared spectra]] of compounds containing the OH [[functional group]] have strong [[spectral line|absorption bands]] in the region centered around 3500&nbsp;cm<sup>−1</sup>.<ref name=nakamoto>{{cite book|last=Nakamoto|first=K.|title=Infrared and Raman spectra of Inorganic and Coordination compounds|edition=5th|series=Part A|year=1997|publisher=Wiley|isbn=978-0-471-16394-7}}</ref> 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 [[spectral linewidth|band width]] increases when the OH group is involved in hydrogen bonding. A water molecule has an HOH bending mode at about 1600&nbsp;cm<sup>−1</sup>, 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<!-- WP:MOSCHEM#Skeletal_formulas --> bending mode can be observed. For example, in [Sn(OH)<sub>6</sub>]<sup>2−</sup> it occurs at 1065&nbsp;cm<sup>−1</sup>. The bending mode for a bridging hydroxide tends to be at a lower frequency as in [([[bipyridine]])Cu(OH)<sub>2</sub>Cu([[bipyridine]])]<sup>2+</sup> (955&nbsp;cm<sup>−1</sup>).<ref>Nakamoto, Part B, p.&nbsp;57</ref> M−OH stretching vibrations occur below about 600&nbsp;cm<sup>−1</sup>. For example, the [[tetrahedron|tetrahedral]] ion [Zn(OH)<sub>4</sub>]<sup>2−</sup> has bands at 470&nbsp;cm<sup>−1</sup> ([[Raman spectroscopy|Raman]]-active, polarized) and 420&nbsp;cm<sup>−1</sup> (infrared). The same ion has a (HO)–Zn–(OH) bending vibration at 300&nbsp;cm<sup>−1</sup>.<ref>{{cite book|last=Adams|first=D.M.|title=Metal–Ligand and Related Vibrations|year=1967|publisher=Edward Arnold|location=London}} Chapter 5.</ref>

==Applications==
[[Sodium hydroxide]] solutions, also known as [[lye]] and caustic soda, are used in the manufacture of [[wood pulp|pulp]] and [[paper]], [[textile]]s, [[drinking water]], [[soap]]s and [[detergent]]s, and as a [[drain cleaner]]. Worldwide production in 2004 was approximately 60&nbsp;million [[tonne]]s.<ref name= Ullmann>{{Ullmann|doi= 10.1002/14356007.a24_345.pub2|title= Sodium Hydroxide|author= Cetin Kurt, Jürgen Bittner |}}</ref> 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
:{{chem|CO|3|2−}} + H<sub>2</sub>O {{eqm}} {{chem|HCO|3|−}} + OH<sup>−</sup> {{spaces|5}} ([[acid dissociation constant|p''K''<sub>a2</sub>]] = 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 [[ester]]s, such as [[triglyceride]]s, 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.<ref>{{cite book|title= Nature's Building Blocks: An A–Z Guide to the Elements|last= Emsley|first= John|publisher= Oxford University Press|year= 2001|location= Oxford, UK|isbn= 978-0-19-850340-8|chapter= Aluminium|page= [https://archive.org/details/naturesbuildingb0000emsl/page/24 24]|chapter-url= https://books.google.com/books?id=j-Xu07p3cKwC&pg=PA24|url= https://archive.org/details/naturesbuildingb0000emsl/page/24}}</ref> 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.<ref>{{cite book|title= Nature's Building Blocks: An A–Z Guide to the Elements|last= Emsley|first= John|publisher= Oxford University Press|year= 2001|location= Oxford, UK|isbn= 978-0-19-850340-8|chapter= Aluminium|page= [https://archive.org/details/naturesbuildingb0000emsl/page/209 209]|chapter-url= https://books.google.com/books?id=j-Xu07p3cKwC&pg=PA209|url= https://archive.org/details/naturesbuildingb0000emsl/page/209}}</ref>
<!--==Notes==
*[http://pubs3.acs.org/acs/journals/doilookup?in_doi=10.1021/j100016a003 Solvation and Transport of H<sub>3</sub>O<sup>+</sup> and OH<sup>−</sup> Ions in Water (JCP 99, 5749 (1995)]
The original article was a direct copy of a Russian encyclopedia article!-->

==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]] (LiOH) is used in [[breathing gas]] purification systems for [[spacecraft]], [[submarine]]s, and [[rebreather]]s to remove [[carbon dioxide]] from exhaled gas.<ref>{{cite journal |last=Jaunsen |first=JR |title=The Behavior and Capabilities of Lithium Hydroxide Carbon Dioxide Scrubbers in a Deep Sea Environment |journal=US Naval Academy Technical Report |volume=USNA-TSPR-157 |year=1989 |url=http://archive.rubicon-foundation.org/4998 |access-date=2008-06-17 |archive-url=https://web.archive.org/web/20090824104846/http://archive.rubicon-foundation.org/4998 |archive-date=2009-08-24 |url-status=usurped }}</ref>
:2 LiOH + CO<sub>2</sub> → Li<sub>2</sub>CO<sub>3</sub> + H<sub>2</sub>O
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 metal]]s are also [[strong base]]s.<ref>Holleman, p.&nbsp;1108</ref>

===Alkaline earth metals===
[[File:Beryllium trimer.svg|thumb|left|130px|Trimeric hydrolysis product of beryllium dication<ref group=note>In aqueous solution the ligands L are water molecules, but they may be replaced by other ligands</ref>]]
[[File:BeHydrolysis.png|thumb|Beryllium hydrolysis as a function of pH. Water molecules attached to Be are omitted.]]
[[Beryllium hydroxide]] Be(OH)<sub>2</sub> is [[amphoteric]].<ref name=amph>Thomas R. Dulski [https://books.google.com/books?id=ViOMjoLKB1gC&pg=PA100 A manual for the chemical analysis of metals], ASTM International, 1996, {{ISBN|0-8031-2066-4}} p.&nbsp;100</ref> The hydroxide itself is [[insoluble]] in water, with a [[solubility product]] log&nbsp;''K''*<sub>sp</sub> of −11.7. Addition of acid gives soluble [[hydrolysis]] products, including the trimeric ion [Be<sub>3</sub>(OH)<sub>3</sub>(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup>, which has OH groups bridging between pairs of beryllium ions making a 6-membered ring.<ref>{{cite journal|last=Alderighi|first=L|author2=Dominguez, S. |author3=Gans, P. |author4=Midollini, S. |author5=Sabatini, A. |author6= Vacca, A. |year=2009|title=Beryllium binding to adenosine 5'-phosphates in aqueous solution at 25°C|journal=J. Coord. Chem.|volume=62|issue=1|pages=14–22|doi=10.1080/00958970802474862|s2cid=93623985}}</ref> At very low pH the [[Metal ions in aqueous solution|aqua ion]] [Be(H<sub>2</sub>O)<sub>4</sub>]<sup>2+</sup> is formed. Addition of hydroxide to Be(OH)<sub>2</sub> gives the soluble tetrahydroxoberyllate or tetrahydroxido[[beryllate]] anion, [Be(OH)<sub>4</sub>]<sup>2−</sup>.

The solubility in water of the other hydroxides in this group increases with increasing [[atomic number]].<ref>Housecroft, p.&nbsp;241</ref> [[Magnesium hydroxide]] Mg(OH)<sub>2</sub> 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)<sub>2</sub> + CO<sub>2</sub> {{eqm}} Ca<sup>2+</sup> + {{chem|HCO|3|−}} + OH<sup>−</sup> illustrates the basicity of calcium hydroxide. [[Soda lime]], which is a mixture of the strong bases NaOH and KOH with Ca(OH)<sub>2</sub>, is used as a CO<sub>2</sub> absorbent.

===Boron group elements===
[[File:AlHydrolysis.png|thumb|Aluminium hydrolysis as a function of pH. Water molecules attached to Al are omitted]]
The simplest hydroxide of boron B(OH)<sub>3</sub>, 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)<sub>3</sub> + H<sub>2</sub>O {{eqm}} [[tetrahydroxyborate|{{chem|B(OH)|4|−}}]] + H<sup>+</sup>
A variety of [[oxyanion]]s of boron are known, which, in the protonated form, contain hydroxide groups.<ref>Housectroft, p.&nbsp;263</ref>
[[File:Tetrahydroxoaluminate ion.svg|thumb|100px|left|Tetrahydroxo-<br>aluminate(III) ion]]
[[Aluminium hydroxide]] Al(OH)<sub>3</sub> is amphoteric and dissolves in alkaline solution.<ref name=amph/>
:Al(OH)<sub>3</sub> (solid) + OH<sup>−</sup>&nbsp;(aq) {{eqm}} [[aluminate|{{chem|Al(OH)|4|−}}]]&nbsp;(aq)
In the [[Bayer process]]<ref>[http://www.world-aluminium.org/?pg=85 Bayer process chemistry]</ref> 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 {{chem|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 [Al<sub>13</sub>(OH)<sub>32</sub>]<sup>7+</sup> 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]],<ref name=amph/> [[indium hydroxide]], and [[thallium(III) hydroxide]] are also amphoteric. [[Thallium(I) hydroxide]] is a strong base.<ref>James E. House [https://books.google.com/books?id=ocKWuxOur-kC&pg=PA764 Inorganic chemistry], Academic Press, 2008, {{ISBN|0-12-356786-6}}, p.&nbsp;764</ref>

===Carbon group elements===
Carbon forms no simple hydroxides. The [[hypothetical compound]] C(OH)<sub>4</sub> ([[orthocarbonic acid]] or methanetetrol) is unstable in aqueous solution:<ref>{{Cite journal|last1=Böhm|first1=Stanislav|last2=Antipova|first2=Diana|last3=Kuthan|first3=Josef|date=1997|title=A study of methanetetraol dehydration to carbonic acid|journal=International Journal of Quantum Chemistry|language=en|volume=62|issue=3|pages=315–322|doi=10.1002/(SICI)1097-461X(1997)62:3<315::AID-QUA10>3.0.CO;2-8|issn=1097-461X}}</ref>

:C(OH)<sub>4</sub> → {{chem|HCO|3|−}} + H<sub>3</sub>O<sup>+</sup>
:{{chem|HCO|3|−}} + H<sup>+</sup> {{eqm}} H<sub>2</sub>CO<sub>3</sub>
[[Carbon dioxide]] is also known as carbonic anhydride, meaning that it forms by dehydration of [[carbonic acid]] H<sub>2</sub>CO<sub>3</sub> (OC(OH)<sub>2</sub>).<ref>Greenwood, p.&nbsp;310</ref>

[[Silicic acid]] is the name given to a variety of compounds with a generic formula [SiO<sub>''x''</sub>(OH)<sub>4−2''x''</sub>]<sub>''n''</sub>.<ref>Greenwood, p.&nbsp;346</ref><ref>R. K. Iler, ''The Chemistry of Silica'', Wiley, New York, 1979 {{ISBN|0-471-02404-X}}</ref> Orthosilicic acid has been identified in very dilute aqueous solution. It is a weak acid with p''K''<sub>a1</sub>&nbsp;=&nbsp;9.84, p''K''<sub>a2</sub>&nbsp;=&nbsp;13.2 at 25&nbsp;°C. It is usually written as H<sub>4</sub>SiO<sub>4</sub>, but the formula Si(OH)<sub>4</sub> is generally accepted.<ref name=scdb/>{{dubious|discuss|date=November 2014}} Other silicic acids such as metasilicic acid (H<sub>2</sub>SiO<sub>3</sub>), disilicic acid (H<sub>2</sub>Si<sub>2</sub>O<sub>5</sub>), and pyrosilicic acid (H<sub>6</sub>Si<sub>2</sub>O<sub>7</sub>) have been characterized. These acids also have hydroxide groups attached to the silicon; the formulas suggest that these acids are protonated forms of poly[[oxyanion]]s.

Few hydroxo complexes of [[germanium]] have been characterized. [[Tin(II) hydroxide]] Sn(OH)<sub>2</sub> was prepared in anhydrous media. When [[tin(II) oxide]] is treated with alkali the pyramidal hydroxo complex {{chem|Sn(OH)|3|−}} is formed. When solutions containing this ion are acidified, the ion [Sn<sub>3</sub>(OH)<sub>4</sub>]<sup>2+</sup> is formed together with some basic hydroxo complexes. The structure of [Sn<sub>3</sub>(OH)<sub>4</sub>]<sup>2+</sup> has a triangle of tin atoms connected by bridging hydroxide groups.<ref>Greenwood, p.&nbsp;384</ref> Tin(IV) hydroxide is unknown but can be regarded as the hypothetical acid from which [[stannate]]s, with a formula [Sn(OH)<sub>6</sub>]<sup>2−</sup>, are derived by reaction with the (Lewis) basic hydroxide ion.<ref>Greenwood, pp.&nbsp;383–384</ref>

Hydrolysis of Pb<sup>2+</sup> in aqueous solution is accompanied by the formation of various hydroxo-containing complexes, some of which are insoluble. The basic hydroxo complex [Pb<sub>6</sub>O(OH)<sub>6</sub>]<sup>4+</sup> 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 Pb<sub>4</sub> tetrahedra. In strongly alkaline solutions soluble [[plumbate]] ions are formed, including [Pb(OH)<sub>6</sub>]<sup>2−</sup>.<ref>Greenwood, p.&nbsp;395</ref>

===Other main-group elements===
{|class="wikitable" style="text-align:center"
|[[File:Phosphonic-acid-2D-dimensions-vector.svg|center|150px]]
|[[File:Phosphoric-acid-2D-dimensions.svg|center|180px]]
|[[File:Sulfuric-acid-2D-dimensions.svg|center|180px]]
|[[File:Telluric acid.svg|center|150px]]
|[[File:Ortho-Periodsäure.svg|center|150px]]
|[[File:Xenic acid.png|center|150px]]
|-
|[[Phosphorous acid]]
|[[Phosphoric acid]]
|[[Sulfuric acid]]
|[[Telluric acid]]
|[[Orthoperiodic acid]]
|[[Xenic acid]]
|}

In the higher oxidation states of the [[pnictogen]]s, [[chalcogen]]s, [[halogen]]s, and [[noble gas]]es there are oxoacids in which the central atom is attached to oxide ions and hydroxide ions. Examples include [[phosphoric acid]] H<sub>3</sub>PO<sub>4</sub>, and [[sulfuric acid]] H<sub>2</sub>SO<sub>4</sub>. In these compounds one or more hydroxide groups can [[dissociation (chemistry)|dissociate]] with the liberation of hydrogen cations as in a standard [[Brønsted–Lowry acid–base theory|Brønsted–Lowry]] acid. Many oxoacids of sulfur are known and all feature OH groups that can dissociate.<ref>Greenwood, p.&nbsp;705</ref>

[[Telluric acid]] is often written with the formula H<sub>2</sub>TeO<sub>4</sub>·2H<sub>2</sub>O but is better described structurally as Te(OH)<sub>6</sub>.<ref>Greenwood, p.&nbsp;781</ref>

Orthoperiodic acid<ref group=note>The name is '''not''' derived from "period", but from "iodine": periodic acid (compare [[iodic acid]], [[perchloric acid]]), and it is thus pronounced per-iodic {{IPAc-en|ˌ|p|ɜːr|aɪ|ˈ|ɒ|d|ᵻ|k}} {{respell|PUR|eye|OD|ik}}, and not as {{IPAc-en|ˌ|p|ɪər|ɪ|-}} {{respell|PEER|ee-}}.</ref> can lose all its protons, eventually forming the periodate ion [IO<sub>4</sub>]<sup>−</sup>. It can also be protonated in strongly acidic conditions to give the octahedral ion [I(OH)<sub>6</sub>]<sup>+</sup>, completing the [[isoelectronic]] series, [E(OH)<sub>6</sub>]<sup>''z''</sup>, 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 K<sub>4</sub>[I<sub>2</sub>O<sub>8</sub>(OH)<sub>2</sub>]·8H<sub>2</sub>O.<ref>Greenwood, pp.&nbsp;873–874</ref>

As is common outside of the alkali metals, hydroxides of the elements in lower oxidation states are complicated. For example, [[phosphorous acid]] H<sub>3</sub>PO<sub>3</sub> predominantly has the structure OP(H)(OH)<sub>2</sub>, in equilibrium with a small amount of P(OH)<sub>3</sub>.<ref>{{cite journal|title= Stabilization of tautomeric forms P(OH)<sub>3</sub> and HP(OH)<sub>2</sub> and their derivatives by coordination to palladium and nickel atoms in heterometallic clusters with the {{chem|Mo|3|MQ|4|4+}} core (M = Ni, Pd; Q = S, Se) |author=M. N. Sokolov |author2=E. V. Chubarova |author3=K. A. Kovalenko |author4=I. V. Mironov |author5=A. V. Virovets |author6=E. Peresypkina |author7=V. P. Fedin |doi= 10.1007/s11172-005-0296-1|year= 2005|journal= Russian Chemical Bulletin|volume= 54|pages= 615|issue= 3|s2cid=93718865 }}</ref><ref>Holleman, pp.&nbsp;711–718</ref>

The oxoacids of [[chlorine]], [[bromine]], and [[iodine]] have the formula O<sub>{{sfrac|''n''−1|2}}</sub>A(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.<ref>Greenwood, p.&nbsp;853</ref>

===Transition and post-transition metals===
The hydroxides of the [[transition metal]]s and [[post-transition metal]]s 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 (Ag<sub>2</sub>O). Copper(I) and gold(I) hydroxides are also unstable, although stable adducts of CuOH and AuOH are known.<ref>{{cite journal|last=Fortman|first=George C. |author2=Slawin, Alexandra M. Z. |author3=Nolan, Steven P. |year=2010|title=A Versatile Cuprous Synthon: [Cu(IPr)(OH)] (IPr = 1,3 bis(diisopropylphenyl)imidazol-2-ylidene)|journal=Organometallics|volume=29|issue=17|pages=3966–3972|doi=10.1021/om100733n}}</ref> The polymeric compounds M(OH)<sub>2</sub> and M(OH)<sub>3</sub> are in general prepared by increasing the pH of an aqueous solution of the corresponding metal cation until the hydroxide [[precipitate]]s out of solution. On the converse, the hydroxides dissolve in acidic solution. [[Zinc hydroxide]] Zn(OH)<sub>2</sub> is amphoteric, forming the tetrahydroxido[[zincate]] ion {{chem|Zn(OH)|4|2−}} in strongly alkaline solution.<ref name=amph/>

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]].
:L<sub>''n''</sub>M(OH<sub>2</sub>) + B {{eqm}} L<sub>''n''</sub>M(OH)<sup>–</sup> + BH<sup>+</sup> (L = ligand, B = base)

[[Vanadic acid]] H<sub>3</sub>VO<sub>4</sub> [[acid dissociation constant#Polyprotic acids|shows similarities]] with phosphoric acid H<sub>3</sub>PO<sub>4</sub> though it has a much more complex [[vanadate]] oxoanion chemistry. [[Chromic acid]] H<sub>2</sub>CrO<sub>4</sub>, has similarities with sulfuric acid H<sub>2</sub>SO<sub>4</sub>; for example, both form [[acid salt]]s A<sup>+</sup>[HMO<sub>4</sub>]<sup>−</sup>. 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 [[polyoxometalate]]s.<ref>Juan J. Borrás-Almenar, Eugenio Coronado, Achim Müller [https://books.google.com/books?id=RJwQO1Ip0SQC&pg=PA4 Polyoxometalate Molecular Science], Springer, 2003, {{ISBN|1-4020-1242-X}}, p.&nbsp;4</ref>

==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 Zr<sup>4+</sup> are extensively hydrolyzed in water even at low pH. The compound originally formulated as ZrOCl<sub>2</sub>·8H<sub>2</sub>O was found to be the chloride salt of a [[tetrameric]] cation [Zr<sub>4</sub>(OH)<sub>8</sub>(H<sub>2</sub>O)<sub>16</sub>]<sup>8+</sup> in which there is a square of Zr<sup>4+</sup> 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.<ref name=wells561>Wells, p.&nbsp;561</ref>

The mineral [[malachite]] is a typical example of a basic carbonate. The formula, Cu<sub>2</sub>CO<sub>3</sub>(OH)<sub>2</sub> shows that it is halfway between [[basic copper carbonate|copper carbonate]] and [[copper hydroxide]]. Indeed, in the past the formula was written as CuCO<sub>3</sub>·Cu(OH)<sub>2</sub>. The [[crystal structure]] is made up of copper, carbonate and hydroxide ions.<ref name=wells561/> The mineral [[atacamite]] is an example of a basic chloride. It has the formula Cu<sub>2</sub>Cl(OH)<sub>3</sub>. In this case the composition is nearer to that of the hydroxide than that of the chloride: CuCl<sub>2</sub>·3Cu(OH)<sub>2</sub>.<ref>Wells, p.&nbsp;393</ref> Copper forms hydroxyphosphate ([[libethenite]]), arsenate ([[olivenite]]), sulfate ([[brochantite]]), and nitrate compounds. [[White lead]] is a basic [[lead]] carbonate, (PbCO<sub>3</sub>)<sub>2</sub>·Pb(OH)<sub>2</sub>, 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]].<ref name=wells561/>

==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&nbsp;pm.<ref name=wells548/> Thus, the high-temperature forms of KOH and NaOH have the [[sodium chloride#Crystal structure|sodium chloride]] structure,<ref>Victoria M. Nield, David A. Keen [https://books.google.com/books?id=QDT6VQ4GRL8C&pg=PA276 Diffuse neutron scattering from crystalline materials], Oxford University Press, 2001 {{ISBN|0-19-851790-4}}, p.&nbsp;276</ref> which gradually freezes in a monoclinically distorted sodium chloride structure at temperatures below about 300&nbsp;°C. The OH groups still rotate even at room temperature around their symmetry axes and, therefore, cannot be detected by [[X-ray diffraction]].<ref>{{cite journal|last1=Jacobs|first1=H.|last2=Kockelkorn|first2=J.|last3=Tacke|first3=Th.|title=Hydroxide des Natriums, Kaliums und Rubidiums: Einkristallzüchtung und röntgenographische Strukturbestimmung an der bei Raumtemperatur stabilen Modifikation|journal=Zeitschrift für Anorganische und Allgemeine Chemie|volume=531|pages=119|year=1985|doi=10.1002/zaac.19855311217|issue=12}}</ref> The room-temperature form of NaOH has the [[thallium(I) iodide|thallium iodide]] structure. LiOH, however, has a layered structure, made up of tetrahedral Li(OH)<sub>4</sub> and (OH)Li<sub>4</sub> units.<ref name=wells548>Wells, p.&nbsp;548</ref> 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)<sub>2</sub> ([[brucite]]) crystallizes with the [[cadmium iodide]] layer structure, with a kind of close-packing of magnesium and hydroxide ions.<ref name=wells548/><ref>{{cite journal|last1=Enoki|first1=Toshiaki|last2=Tsujikawa|first2=Ikuji|title=Magnetic Behaviours of a Random Magnet, Ni<sub>''p''</sub>Mg<sub>1−''p''</sub>(OH)<sub>2</sub>|journal=Journal of the Physical Society of Japan|volume=39|pages=317|year=1975|doi=10.1143/JPSJ.39.317|issue=2|bibcode= 1975JPSJ...39..317E}}</ref>

The [[amphoterism|amphoteric]] hydroxide Al(OH)<sub>3</sub> has four major crystalline forms: [[gibbsite]] (most stable), [[bayerite]], [[nordstrandite]], and [[doyleite]].<ref group=note>Crystal structures are illustrated at Web mineral: [http://webmineral.com/data/Gibbsite.shtml Gibbsite], [http://webmineral.com/data/Bayerite.shtml Bayerite], [http://webmineral.com/data/Nordstrandite.shtml Norstrandite] and [http://webmineral.com/data/Doyleite.shtml Doyleite]</ref>
All these [[Polymorphism (materials science)|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.<ref>Athanasios K. Karamalidis, David A. Dzombak [https://books.google.com/books?id=XULsOFSipsgC&pg=PA15 Surface Complexation Modeling: Gibbsite], John Wiley and Sons, 2010 {{ISBN|0-470-58768-7}} pp. 15 ff</ref> 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.<ref>{{cite journal|last=Bernal|first=J.D.|author2=Megaw, H.D.|year=1935|title=The Function of Hydrogen in Intermolecular Forces|journal=Proc. R. Soc. A|volume=151|pages=384–420|doi=10.1098/rspa.1935.0157|issue=873|bibcode= 1935RSPSA.151..384B|doi-access=free}}</ref> This is an unusual form of [[hydrogen bond]]ing 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&nbsp;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.<ref>Wells, p.&nbsp;557</ref> A similar type of hydrogen bond has been proposed for other amphoteric hydroxides, including Be(OH)<sub>2</sub>, Zn(OH)<sub>2</sub>, and Fe(OH)<sub>3</sub>.<ref name=wells548/>

A number of mixed hydroxides are known with stoichiometry A<sub>3</sub>M<sup>III</sup>(OH)<sub>6</sub>, A<sub>2</sub>M<sup>IV</sup>(OH)<sub>6</sub>, and AM<sup>V</sup>(OH)<sub>6</sub>. As the formula suggests these substances contain M(OH)<sub>6</sub> octahedral structural units.<ref>Wells, p.&nbsp;555</ref> [[Layered double hydroxides]] may be represented by the formula {{chem|[M|1−''x''|''z''+|M|''x''|3+|(OH)|2|]<sup>''q''+</sup>(X<sup>''n''−</sup>)|{{frac|''q''|''n''}}|·''y''H|2|O}}. Most commonly, ''z''&nbsp;=&nbsp;2, and M<sup>2+</sup> = Ca<sup>2+</sup>, Mg<sup>2+</sup>, Mn<sup>2+</sup>, Fe<sup>2+</sup>, Co<sup>2+</sup>, Ni<sup>2+</sup>, Cu<sup>2+</sup>, or Zn<sup>2+</sup>; hence ''q''&nbsp;=&nbsp;''x''.

==Organic reactions==
[[Potassium hydroxide]] and [[sodium hydroxide]] are two well-known [[reagent]]s in [[organic chemistry]].

===Base catalysis===
The hydroxide ion may act as a [[base catalyst]].<ref>{{cite book|editor-last1= Hattori |editor-first1=H.|editor-last2=Misono|editor-first2=M.|editor-last3=Ono|editor-first3=Y.|title=Acid–Base catalysis II|year=1994|publisher=Elsevier|isbn= 978-0-444-98655-9}}</ref> 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 [[Alcohol (chemistry)|alcohol]]s, [[phenol]]s, [[amine]]s, and [[carbon acid]]s. The [[acid dissociation constant|p''K''<sub>a</sub>]] value for dissociation of a C–H bond is extremely high, but the pK<sub>a</sub> [[alpha hydrogen]]s of a carbonyl compound are about 3 log units lower. Typical p''K''<sub>a</sub> values are 16.7 for [[acetaldehyde]] and 19 for [[acetone]].<ref>Ouellette, R.J. and Rawn, J.D. "Organic Chemistry" 1st Ed. Prentice-Hall, Inc., 1996: New Jersey. {{ISBN|0-02-390171-3}}.</ref> Dissociation can occur in the presence of a suitable base.
:RC(O)CH<sub>2</sub>R' + B {{eqm}} RC(O)CH<sup>−</sup>R' + BH<sup>+</sup>
The base should have a p''K''<sub>a</sub> 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<sup>−</sup> + EtOH {{eqm}} EtO<sup>−</sup> + H<sub>2</sub>O
to produce the [[ethoxide]] ion. The [[acid dissociation constant|pK<sub>a</sub>]] for self-dissociation of ethanol is about 16, so the alkoxide ion is a strong enough base.<ref>{{cite book |title=Organic chemistry |url=https://archive.org/details/organicchemistry04pine |url-access=limited |last=Pine |first=S.H.
|author2=Hendrickson, J.B. |author3=Cram, D.J. |author4= Hammond, G.S. |year=1980 |publisher=McGraw–Hill |isbn=978-0-07-050115-7
|page=[https://archive.org/details/organicchemistry04pine/page/206 206]}}</ref> 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.<ref>{{cite journal|journal=Angewandte Chemie International Edition |last=Denmark |first=S.E.|author2=Beutne, G.L. |year=2008 |title=Lewis Base Catalysis in Organic Synthesis |volume=47 |issue=9 |pages=1560–1638|doi=10.1002/anie.200604943|pmid=18236505 }}</ref>

===As a nucleophilic reagent===
[[File:AcylSubstitution.svg|thumb|300px|Nucleophilic acyl substitution with an anionic [[nucleophile]] (Nu<sup>&minus;</sup>) and [[leaving group]] (L<sup>&minus;</sup>)]]
The hydroxide ion is intermediate in [[nucleophilicity]] between the [[fluoride]] ion F<sup>−</sup>, and the [[azanide|amide]] ion {{chem|NH|2|−}}.<ref>{{cite journal|last=Mullins|first=J. J.|year=2008|title=Six Pillars of Organic Chemistry|journal=[[J. Chem. Educ.]]|volume=85|issue=1|page=83|doi=10.1021/ed085p83|bibcode= 2008JChEd..85...83M}}[http://bastoslab.com/2010/wp-content/uploads/2010/09/ed085p83.pdf pdf] {{Webarchive|url=https://web.archive.org/web/20110707213623/http://bastoslab.com/2010/wp-content/uploads/2010/09/ed085p83.pdf |date=2011-07-07 }}</ref> [[Ester hydrolysis]] under alkaline conditions (also known as [[base hydrolysis]])
:R<sup>1</sup>C(O)OR<sup>2</sup> + OH<sup>&minus;</sup> {{eqm}} R<sup>1</sup>CO(O)H + <sup>&minus;</sup>OR<sup>2</sup> {{eqm}} R<sup>1</sup>CO<sub>2</sub><sup>&minus;</sup> + HOR<sup>2</sup>
is an example of a hydroxide ion serving as a nucleophile.<ref name=Hardinger>{{cite web|url = http://www.chem.ucla.edu/~harding/IGOC/S/saponification.html|title = Illustrated Glossary of Organic Chemistry: Saponification|year = 2017|first = Steven A.|last = Hardinger|publisher = [[UCLA College of Letters and Science|Department of Chemistry & Biochemistry, UCLA]]|access-date = April 10, 2023}}</ref> 

Early methods for [[saponification|manufacturing soap]] treated [[triglyceride]]s 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 reaction]]s. 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==
{{Reflist|group=note}}

==References==
{{Reflist|30em}}

==Bibliography==
*{{cite book|last1=Holleman|first1=A.F.|last2=Wiberg|first2=E.|last3=Wiberg|first3=N.|title=Inorganic Chemistry|year=2001|publisher=Academic press|isbn=978-0-12-352651-9}}
*{{Housecroft3rd}}
*{{Greenwood&Earnshaw2nd}}
*{{cite book|last1=Shriver|first1=D.F|last2=Atkins|first2=P.W|title=Inorganic Chemistry|edition=3rd|year=1999|publisher=Oxford University Press|location=Oxford|isbn=978-0-19-850330-9}}
*{{cite book|last=Wells|first=A.F|title=Structural Inorganic Chemistry|edition=3rd.|year=1962|publisher=Clarendon Press|location=Oxford|isbn=978-0-19-855125-6}}

{{Hydroxides}}

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[[Category:Hydroxides| ]]
[[Category:Oxyanions]]
[[Category:Water chemistry]]