Editing Hydroxide (section)

==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>