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[[File:Autoprotolyse eau.svg|thumb|upright=1.5]]
{{Acids_and_Bases}}
The '''self-ionization of water''' (also '''autoionization of water''', and '''autodissociation of water''') is the chemical reaction in which two [[water molecule]]s react to produce a [[hydronium]] ion (H<sub>3</sub>O<sup>+</sup>) and a [[hydroxide]] ion (OH<sup>−</sup>):

: 2 H<sub>2</sub>O (l) {{eqm}} H<sub>3</sub>O<sup>+</sup> (aq) + OH<sup>−</sup> (aq)

It is an example of [[autoprotolysis]], and relies on the [[amphoteric]] nature of water.

Water, however pure, is not a simple collection of H<sub>2</sub>O molecules.  Even in "pure" water, sensitive equipment can detect a very slight electrical [[Conductivity (electrolytic)|conductivity]] of 0.055  [[Siemens (unit)|&micro;S]]·cm<sup>−1</sup>. According to the theories of [[Svante Arrhenius]], this must be due to the presence of ions.

==Concentration and frequency==
The preceding reaction has a chemical [[equilibrium constant]] of ''K''<sub>eq</sub> = ([H<sub>3</sub>O<sup>+</sup>] [OH<sup>−</sup>]) / [H<sub>2</sub>O]<sup>2</sup> = 3.23 &times; 10<sup>−18</sup>. So the [[acidity constant]] which is K<sub>a</sub> = K<sub>eq</sub> &times; [H<sub>2</sub>O] = ([H<sub>3</sub>O<sup>+</sup>] [OH<sup>−</sup>]) / [H<sub>2</sub>O] = 1.8 &times; 10<sup>−16</sup>.<ref>McMurry, John. (2004) Organic Chemistry, pg 44</ref> For reactions in water (or diluted aqueous solutions), the [[molarity]] (a unit of concentration) of water, [H<sub>2</sub>O], is practically constant and is omitted from the acidity constant expression by convention.  The resulting equilibrium constant is called the '''ionization constant''', '''dissociation constant''', or '''self-ionization constant''', or ion product '''of water''' and is symbolized by K<sub>w</sub>.

:''K''<sub>w</sub> = ''K''<sub>a</sub> [H<sub>2</sub>O] = ''K''<sub>eq</sub> [H<sub>2</sub>O]<sup>2</sup> = [H<sub>3</sub>O<sup>+</sup>] [OH<sup>−</sup>]

:where
::[H<sub>3</sub>O<sup>+</sup>] = molarity of hydrogen or hydronium ion, and
::[OH<sup>−</sup>] = molarity of hydroxide ion.

At [[Standard Ambient Temperature and Pressure]] (SATP), about 25 °C (298 K), K<sub>w</sub> = [H<sub>3</sub>O<sup>+</sup>][OH<sup>−</sup>] = 1.0&times;10<sup>−14</sup>. Pure water ionizes or dissociates into equal amounts of H<sub>3</sub>O<sup>+</sup> and OH<sup>−</sup>, so their molarities are equal:

:[H<sub>3</sub>O<sup>+</sup>] = [OH<sup>−</sup>].

At SATP, the concentrations of hydroxide and hydronium are both very low at 1.0 &times; 10<sup>−7</sup>&nbsp;mol/L and the ions are rarely produced: a randomly selected water molecule will dissociate within approximately 10 hours.<ref>{{cite journal | author = Eigen, M.; de Maeyer, L. | year = 1955 | title = Untersuchungen über die Kinetik der Neutralisation I | journal = [[Z. Elektrochem.]] | volume = 59 | pages = 986}}</ref> Since the concentration of water molecules in water is largely unaffected by dissociation and [H<sub>2</sub>O] equals approximately 56&nbsp;mol/l, it follows that for every 5.6&times;10<sup>8</sup> water molecules, one pair will exist as ions.  Any solution in which the H<sub>3</sub>O<sup>+</sup> and OH<sup>−</sup> concentrations equal each other is considered a '''neutral''' solution.  Absolutely pure water is neutral, although even trace amounts of impurities could affect these ion concentrations and the water may no longer be neutral.  ''K''<sub>w</sub> is sensitive to both pressure and temperature; it increases when either increases.

[[Deionized water]] (also called DI water) is water that has had most ''impurity'' ions common in tap water or natural water sources (such as Na<sup>+</sup> and Cl<sup>−</sup>) removed by means of [[distillation]] or some other [[water purification]] method. Removal of ''all'' ions from water is next to impossible, since water self-ionizes quickly to reach equilibrium.

==Dependence on temperature and pressure==
[[Image:Temperature dependence water ionization.svg|thumb|300px|Temperature dependence of the water ionization constant at 25 MPa]]
[[Image:Pressure dependence water ionization pKw on P.svg|thumb|300px|Pressure dependence of the water ionization constant at 25 °C]]By definition, p''K''<sub>w</sub> = −log<sub>10</sub> ''K''<sub>w</sub>.  At SATP, p''K''<sub>w</sub> = −log<sub>10</sub> (1.0&times;10<sup>−14</sup>) = 14.0. The value of p''K''<sub>w</sub> varies with temperature.  As temperature increases, p''K''<sub>w</sub> decreases; and as temperature decreases, p''K''<sub>w</sub> increases (for temperatures up to about 250 °C). This means that ionization of water typically increases with temperature.

There is also a (usually small) dependence on pressure (ionization increases with increasing pressure). The dependence of the water ionization on temperature and pressure has been well investigated and a standard formulation exists<ref>[http://www.iapws.org/ International Association for the Properties of Water and Steam (IAPWS)<!-- Bot generated title -->]</ref>.

==Acidity==
[[pH]] is a [[logarithm]]ic measure of the acidity (or alkalinity) of an aqueous solution.  By definition, pH = −log<sub>10</sub> [H<sub>3</sub>O<sup>+</sup>].  Since [H<sub>3</sub>O<sup>+</sup>] = [OH<sup>−</sup>] in a neutral solution, by mathematics, for a neutral aqueous solution pH = 7 at SATP.

Self-ionization is the process that determines the [[pH]] of water. Since the concentration of hydronium at SATP (25 °C) is approximately 1.0&times;10<sup>−7</sup>mol/l, the pH of pure liquid water at this temperature is 7.  Since ''K''<sub>w</sub> increases as temperature increases, hot water has a higher concentration of hydronium than cold water, but this does not mean it is more acidic, as the hydroxide concentration is also higher by the same amount.

==Mechanism==
Geissler et al. have determined that electric field fluctuations in liquid water cause molecular dissociation.<ref>{{cite journal | author = Geissler, P. L.; Dellago, C.; Chandler, D.; Hutter, J.; 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}}</ref> They propose the following sequence of events that takes place in about 150 [[femtosecond|fs]]: the system begins in a neutral state; random fluctuations in molecular motions occasionally (about once every 10 hours per water molecule) produce an electric field strong enough to break an oxygen-hydrogen bond, resulting in a hydroxide (OH<sup>−</sup>) and hydronium ion (H<sub>3</sub>O<sup>+</sup>); the proton of the hydronium ion travels along water molecules by the [[Grotthuss mechanism]]; and a change in the hydrogen bond network in the solvent isolates the two ions, which are stabilized by solvation.

Within 1 [[picosecond]], however, a second reorganization of the hydrogen bond network allows rapid proton transfer down the electric potential difference and subsequent recombination of the ions. This timescale is consistent with the time it takes for hydrogen bonds to reorient themselves in water.<ref>{{cite journal | author = Stillinger, F. H. | year = 1975 | journal = [[Adv. Chem. Phys.]] | volume = 31 | pages = 1 | doi = 10.1002/9780470143834.ch1 | title = Theory and Molecular Models for Water}}</ref><ref>{{cite journal | author = Rapaport, D. C. | year = 1983 | journal = [[Mol. Phys.]] | volume = 50 | pages = 1151 | doi = 10.1080/00268978300102931 | title = Hydrogen bonds in water}}</ref><ref>{{cite journal | author = Chen, S.-H. &amp; Teixeira, J. | year = 1986 | journal = [[Adv. Chem. Phys]] | volume = 64 | pages = 1 | doi = 10.1002/9780470142882.ch1 | title = Structure and Dynamics of Low-Temperature Water as Studied by Scattering Techniques}}</ref>

==Isotope effects==
[[Heavy water]], D<sub>2</sub>O, self-ionizes less than normal water, H<sub>2</sub>O; oxygen forms a slightly stronger bond to [[deuterium]] because the larger mass of deuterium difference results in a lower [[zero-point energy]], a quantum mechanical effect. The following table compares the values of p''K''<sub>w</sub> for H<sub>2</sub>O and D<sub>2</sub>O.<ref name=crc>{{cite book | author=Lide, D. R. (Ed.) | title=CRC Handbook of Chemistry and Physics (70th Edn.) | publisher=Boca Raton (FL):CRC Press | year=1990}}</ref>

:{| class="wikitable" style="text-align: center;"
! ''T'' (°C) || H<sub>2</sub>O || D<sub>2</sub>O
|-
| 10 || 14.5346 || 15.439
|-
| 20 || 14.1669 || 15.049
|-
| 25 || 13.9965 || 14.869
|-
| 30 || 13.8330 || 14.699
|-
| 40 || 13.5348 || 14.385
|-
| 50 || 13.2617 || 14.103
|}

==See also==
{{Portal|Water}}
*[[Chemical equilibrium]]
*[[Acid]]
*[[base (chemistry)|Base]]
*[[Acid-base reaction theories]]

==References==
<references/>

==External links==
* [http://www.vias.org/genchem/acidbase_equ_12591_04.html General Chemistry]—Autoionization of Water

{{DEFAULTSORT:Self-Ionization Of Water}}
[[Category:Water chemistry]]
[[Category:Equilibrium chemistry]]

[[ca:Producte iònic de l'aigua]]
[[cs:Autoionizace vody]]
[[da:Vands autoprotolyse]]
[[el:Αυτοϊοντισμός νερού]]
[[es:Autoionización del agua]]
[[fr:Autoprotolyse]]
[[it:Autoionizzazione dell'acqua]]
[[mk:Автојонизација на водата]]
[[ja:自己解離]]
[[pl:Autodysocjacja wody]]
[[pt:Autoionização da água]]
[[fi:Autoprotolyysi]]
[[sv:Vattnets autoprotolys]]
[[zh:水的电离]]
[[Category:2Fix]]