Editing Hydroxide (section)

==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>