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{{Short description|Negatively charged polyatomic ion containing oxygen}}An '''oxyanion''', or '''oxoanion''', is an [[ion]] with the generic formula {{chem|A|''x''|O|''y''|''z''−}} (where A represents a [[chemical element]] and O represents an [[oxygen]] atom). Oxyanions are formed by a large majority of the [[chemical element]]s.<ref>{{Greenwood&Earnshaw}}</ref> The formulae of simple oxyanions are determined by the [[octet rule]].  The corresponding [[oxyacid]] of an oxyanion is the compound {{chem|H|''z''|A|''x''|O|''y''}}.  The structures of condensed oxyanions can be rationalized in terms of AO<sub>''n''</sub> polyhedral units with sharing of corners or edges between polyhedra. The oxyanions (specifically, phosphate and polyphosphate esters) adenosine monophosphate ([[adenosine monophosphate|AMP]]), adenosine diphosphate ([[adenosine diphosphate|ADP]]) and [[adenosine triphosphate]] (ATP) are important in biology.

==Monomeric oxyanions==
The formula of [[monomeric]] oxyanions, {{chem|AO|''n''|''m''−}}, is dictated by the [[oxidation state]] of the element A and its position in the [[periodic table]]. Elements of the first row are limited to a maximum coordination number of 4. However, none of the first row elements has a monomeric oxyanion with that coordination number. Instead, [[carbonate]] ({{chem|CO|3|2−}}) and [[nitrate]] ({{chem|NO|3|−}}) have a [[trigonal planar]] structure with [[π bond]]ing between the central atom and the oxygen atoms. This π bonding is favoured by the similarity in size of the central atom and oxygen.

The oxyanions of second-row elements in the [[group (periodic table)|group]] oxidation state are [[tetrahedral molecular geometry|tetrahedral]]. Tetrahedral {{chem2|SiO4}} units are found in [[olivine]] minerals, {{chem2|(Mg,Fe)2SiO4}}, but the anion does not have a separate existence as the oxygen atoms are surrounded tetrahedrally by cations in the solid state. [[Phosphate]] ({{chem|PO|4|3−}}), [[sulfate]] ({{chem|SO|4|2−}}), and [[perchlorate]] ({{chem|ClO|4|−}}) ions can be found as such in various salts. Many oxyanions of elements in lower oxidation state obey the [[octet rule]] and this can be used to rationalize the formulae adopted. For example, chlorine(V) has two valence electrons so it can accommodate three electron pairs from bonds with oxide ions. The charge on the ion is +5&nbsp;−&nbsp;3&nbsp;×&nbsp;2 =&nbsp;−1, and so the formula is {{chem|ClO|3|−}}. The structure of the ion is predicted by [[VSEPR]] theory to be pyramidal, with three bonding electron pairs and one lone pair. In a similar way,
The oxyanion of chlorine(III) has the formula {{chem|ClO|2|−}}, and is bent with two lone pairs and two bonding pairs.

{| class="wikitable"
!Oxidation state
!Name
!Formula
!Image
|-
|align=center|+1
|The [[hypochlorite]] ion
|ClO<sup>−</sup>
|[[Image:Hypochlorite-ion-3D-vdW.png|70px]]
|-
|align=center|+3
|The [[chlorite]] ion
|{{chem|ClO|2|−}}
|[[Image:Chlorite-ion-3D-vdW.png|70px]]
|-
|align=center|+5
|The [[chlorate]] ion
|{{chem|ClO|3|−}}
|[[Image:Chlorate-ion-3D-vdW.png|70px]]
|-
|align=center|+7
|The [[perchlorate]] ion
|{{chem|ClO|4|−}}
|[[Image:Perchlorate-ion-3D-vdW.png|70px]]
|}

In the third and subsequent rows of the periodic table, 6-coordination is possible, but isolated octahedral oxyanions are not known because they would carry too high an electrical charge. Thus molybdenum(VI) does not form {{chem|MoO|6|6−}}, but forms the tetrahedral [[molybdate]] anion, {{chem|MoO|4|2−}}. MoO<sub>6</sub> units are found in condensed molybdates. Fully protonated oxyanions with an octahedral structure are found in such species as {{chem|Sn(OH)|6|2−}} and {{chem|Sb(OH)|6|−}}. In addition, [[orthoperiodate]] can be only partially deprotonated,<ref group="Note"> the high value of the fourth p''K''<sub>a</sub> makes it very unlikely the fifth and sixth deprotonation will occur in water solution.</ref> with 
:<chem>H3IO6^{2-} \ _{\longrightarrow}^{\longleftarrow} \ H2IO6^{3-} \ + \ H^{+}</chem> having p''K''<sub>a</sub>=11.60.<ref>{{cite book|last=Aylett|first=founded by A.F. Holleman; continued by Egon Wiberg; translated by Mary Eagleson, William Brewer; revised by Bernhard J.|title=Inorganic chemistry|year=2001|publisher=Academic Press, W. de Gruyter.|location=San Diego, Calif.: Berlin|isbn=0123526515|page=454|edition=1st English ed., [edited] by Nils Wiberg.}}</ref><ref>{{cite book|last=Burgot|first=Jean-Louis|title=Ionic equilibria in analytical chemistry|publisher=Springer|location=New York|isbn=978-1441983824|page=358|date=2012-03-30}}</ref>

===Naming===
The naming of monomeric oxyanions follows the following rules.

Here the [[Halogen|halogen group]] (group '''7'''A, 17) is referred to as group VII and the [[Noble gas|noble gases group]] (group '''8'''A) is referred to as group VIII.

; If central atom is not in Group VII or VIII
{| class=wikitable
! Central atom oxidation number !! Naming scheme !! Examples
|-
| = Group number || '''*-ate''' || [[Borate]] ({{chem|BO|3|3−}}), [[Carbonate]] ({{chem|CO|3|2−}}), [[Nitrate]] ({{chem|NO|3|−}}), [[Phosphate]] ({{chem|PO|4|3−}}), [[Sulfate]] ({{chem|SO|4|2−}}), [[Chromate ion|Chromate]] ({{chem|CrO|4|2−}}), [[Arsenate]] ({{chem|AsO|4|3−}}), [[Ferrate]] ({{chem|FeO|4|2−}})
|-
| = Group number −&nbsp;2 || '''*-ite''' || [[Nitrite]] ({{chem|NO|2|−}}), [[Phosphite anion|Phosphite]] ({{chem|PO|3|3−}}), [[Sulfite]] ({{chem|SO|3|2−}}), [[Arsenite]] ({{chem|AsO|3|3−}})
|-
| = Group number −&nbsp;4 || '''hypo-*-ite''' || [[Hypophosphite]] ({{chem|PO|2|3−}}), [[Hyposulfite]] ({{chem|SO|2|2−}})
|}
; If central atom is in Group VII or VIII
{| class=wikitable
! Central atom oxidation number !! Naming scheme !! Examples
|-
| = Group number || '''per-*-ate''' || [[Perchlorate]] ({{chem|ClO|4|−}}), [[Perbromate]] ({{chem|BrO|4|−}}), [[Periodate]] ({{chem|IO|4|−}}), [[Permanganate]] ({{chem|MnO|4|−}}), [[Perxenate]] ({{chem|XeO|6|4-}})
|-
| = Group number −&nbsp;2 || '''*-ate''' || [[Chlorate]] ({{chem|ClO|3|−}}), [[Bromate]] ({{chem|BrO|3|−}}), [[Iodate]] ({{chem|IO|3|−}})
|-
| = Group number −&nbsp;4 || '''*-ite''' || [[Chlorite]] ({{chem|ClO|2|−}}), [[Bromite]] ({{chem|BrO|2|−}})
|-
| = Group number −&nbsp;6 || '''hypo-*-ite''' || [[Hypochlorite]] (ClO<sup>−</sup>), [[Hypobromite]] (BrO<sup>−</sup>)
|}

== Condensation reactions ==
[[Image:Dichromate-3D-balls.png|thumb|150 px|The dichromate ion; two tetrahedra share one corner]]
In aqueous solution, oxyanions with high charge can undergo condensation reactions, such as in the formation of the [[dichromate]] ion, {{chem2|Cr2O7(2-)}}:

:<chem>2 CrO4^2- + 2 H+ <=> Cr2O7^2- + H2O</chem>

The driving force for this reaction is the reduction of electrical charge density on the anion and the elimination of the [[hydronium]] ({{chem2|H+}}) ion. The amount of order in the solution is decreased, releasing a certain amount of [[Entropy (order and disorder)|entropy]] which makes the [[Gibbs free energy]] more negative and favors the forward reaction. It is an example of an [[acid–base reaction]] with the monomeric oxyanion acting as a base and the condensed oxyanion acting as its [[conjugate acid]]. The reverse reaction is a [[hydrolysis]] reaction, as a [[water molecule]], acting as a base, is split. Further condensation may occur, particularly with anions of higher charge, as occurs with adenosine phosphates.
{|
|[[Image:AMP structure.svg|200px]]
|[[Image:Adenosindiphosphat protoniert.svg|x150px]]
|[[Image:ATP structure.svg|x150px]]
|-
|align="center"|AMP
|align="center"|ADP
|align="center"|ATP
|}
The conversion of ATP to ADP is a hydrolysis reaction and is an important source of energy in biological systems.

The formation of most [[silicate]] minerals can be viewed as the result of a de-condensation reaction in which [[silica]] reacts with a basic oxide, an acid–base reaction in the [[acid–base reaction#Lux–Flood definition|Lux–Flood]] sense.
:<chem>\overset{base}{CaO} + \overset{acid}{SiO2} -> CaSiO3</chem>

== Structures and formulae of polyoxyanions ==
{{See also|Polyoxometalate}}
[[Image:Ammonium-metavanadate-chains-3D.png|thumb|right|200px|Metavanadate chains in ammonium metavanadate]]
<!-- [[image:Trimetaphosphoric-acid-3D-vdW.png|thumb| 150px|Cyclotriphosphoric acid]] -->A '''polyoxyanion''' is a [[polymer]]ic oxyanion in which multiple oxyanion monomers, usually regarded as {{chem2|MO_{''n''} }} polyhedra, are joined by sharing corners or edges.<ref>{{cite book|last=Mueller|first=U.|title=Inorganic Structural Chemistry |publisher=Wiley|year=1993|isbn=0-471-93717-7 }}</ref> When two corners of a polyhedron are shared the resulting structure may be a chain or a ring. Short chains occur, for example, in [[polyphosphate]]s. Inosilicates, such as [[pyroxene]]s, have a long chain of {{chem2|SiO4}} tetrahedra each sharing two corners. The same structure occurs in so-called meta-vanadates, such as [[ammonium metavanadate]], {{chem2|NH4VO3}}.

The formula of the oxyanion {{chem2|SiO3(2-)}} is obtained as follows: each nominal silicon ion ({{chem2|Si(4+)}}) is attached to two nominal oxide ions ({{chem2|O(2-)}}) and has a half share in two others. Thus the stoichiometry and charge are given by:

:<math chem>\text{Stoichiometry:  } \ce{Si{} + 2O{} + (2 \times 1/2)O} = \ce{SiO3}</math>
:<math>\text{Charge:  } +4 + (2 \times -2) + (2 \times (\tfrac{1}{2} \times -2)) = -2</math>

A ring can be viewed as a chain in which the two ends have been joined. Cyclic [[Phosphoric acids and phosphates|triphosphate]], {{chem2|P3O9(3-)}} is an example.

When three corners are shared the structure extends into two dimensions. In [[amphibole]]s, (of which  [[asbestos]] is an example) two chains are linked together by sharing of a third corner on alternate places along the chain. This results in an ideal formula {{chem2|Si4O11(6-)}} and a linear chain structure which explains the fibrous nature of these minerals. Sharing of all three corners can result in a sheet structure, as in [[mica]], {{chem2|Si2O5(2-)}}, in which each silicon has one oxygen to itself and a half-share in three others. Crystalline mica can be cleaved into very thin sheets.

The sharing of all four corners of the tetrahedra results in a 3-dimensional structure, such as in [[quartz]]. [[Aluminosilicate]]s are minerals in which some silicon is replaced by aluminium. However, the oxidation state of aluminium is one less than that of silicon, so the replacement must be accompanied by the addition of another cation. The number of possible combinations of such a structure is very large, which is, in part, the reason why there are so many aluminosilicates.
[[Image:Decavanadate polyhedra.png|thumb|Decavanadate ion, {{chem2|V10O28(6-)}}]]

Octahedral {{chem2|MO6}} units are common in oxyanions of the larger transition metals. Some compounds, such as salts of the chain-polymeric ion, {{chem2|Mo2O7(2-)}} even contain both tetrahedral and octahedral units.<ref>{{cite journal|last=Lindqvist|first=I.|year=1950|title=Crystal Structure Studies on Anhydrous Sodium Molybdates and Tungstates|journal=Acta Chem. Scand.|volume=4|pages=1066–1074|doi= 10.3891/acta.chem.scand.04-1066|last2=Hassel|first2=O.|last3=Webb|first3=M.|last4=Rottenberg|first4=Max|doi-access=free}}</ref><ref name=Wells>{{cite book|last=Wells|first=A.F.|title=Structural Inorganic Chemistry|publisher=Clarendon Press|location=Oxford|year=1962|edition=3rd}} p446</ref> Edge-sharing is common in ions containing octahedral building blocks and the octahedra are usually distorted to reduce the strain at the bridging oxygen atoms. This results in 3-dimensional structures called [[polyoxometalate]]s. Typical examples occur in the [[Keggin structure]] of the [[phosphomolybdic acid|phosphomolybdate]] ion. Edge sharing is an effective means of reducing electrical charge density, as can be seen with the hypothetical condensation reaction involving two octahedra:

:<chem>2 MO6^{\mathit{n}-}{} + 4H+ -> Mo2O10^{(\mathit{n}-4) - }{} + 2 H2O</chem>

Here, the average charge on each M atom is reduced by 2. The efficacy of edge-sharing is demonstrated by the following reaction, which occurs when an alkaline aqueous solution of molybdate is acidified.

:<chem>7 MoO4^2- + 8 H+ <=> Mo7O24^6- + 4 H2O</chem>

The tetrahedral molybdate ion is converted into a cluster of 7 edge-linked octahedra<ref name=Wells/><ref>{{cite journal|last=Lindqvist|first=I.|year=1950|journal=Arkiv för Kemi|volume=2|pages=325|title=Arkiv för Kemi}}</ref> giving an average charge on each molybdenum of {{frac|6|7}}. The heptamolybdate cluster is so stable that clusters with between 2 and 6 molybdate units have not been detected even though they must be formed as intermediates.

==Heuristic for acidity==
The pKa of the related acids can be guessed from the number of double bonds to oxygen. Thus perchloric acid is a very strong acid while hypochlorous acid is very weak. A simple rule usually works to within about 1&nbsp;pH unit.

== Acid–base properties ==
Most oxyanions are weak [[base (chemistry)|base]]s and can be protonated to give acids or acid salts. For example, the phosphate ion can be successively protonated to form phosphoric acid.

:<chem>PO4^3- + H+ <=> HPO4^2-</chem>
:<chem>HPO4^2- + H+ <=> H2PO4-</chem>
:<chem>H2PO4- + H+ <=> H3PO4</chem>

[[Image:Phosphite-ion-from-xtal-3D-balls.png|thumb|left|100px|{{chem2|HPO3(2-)}} (phosphite ion) structure]]
[[Image:Sulfuric-acid-3D-vdW.png|thumb|100px|right|Sulfuric acid molecule]]
The extent of protonation in aqueous solution will depend on the [[acid dissociation constant]]s and [[pH]]. For example, AMP (adenosine monophosphate) has a p''K''<sub>a</sub> value of 6.21,<ref>{{cite journal|last=da Costa|first=C.P.|author2=Sigel, H.|year=2000|title=Lead(II)-Binding Properties of the 5′-Monophosphates of Adenosine (AMP<sup>2−</sup>), Inosine (IMP<sup>2−</sup>), and Guanosine (GMP<sup>2−</sup>) in Aqueous Solution. Evidence for Nucleobase−Lead(II) Interactions|journal=Inorg. Chem.|volume=39|issue=26|pages=5985–5993|doi=10.1021/ic0007207|pmid=11151499}}</ref> so at pH&nbsp;7 it will be about 10% protonated. Charge neutralization is an important factor in these protonation reactions. By contrast, the univalent anions [[perchlorate]] and [[permanganate]] ions are very difficult to protonate and so the corresponding acids are [[strong acids]].

Although acids such as phosphoric acid are written as {{chem2|H3PO4}}, the protons are attached to oxygen atoms forming hydroxyl groups, so the formula can also be written as {{chem2|OP(OH)3}} to better reflect the structure. Sulfuric acid may be written as {{chem2|O2S(OH)2}}; this is the molecule observed in the gas phase.

The [[phosphite]] ion, {{chem2|PO3(3-)}}, is a [[strong base]], and so always carries at least one proton. In this case the proton is attached directly to the phosphorus atom with the structure {{chem2|HPO3(2-)}}. In forming this ion, the phosphite ion is behaving as a [[Lewis base]] and donating a pair of electrons to the Lewis acid, {{chem2|H+}}.

[[Image:Predominance diagram Cr.png|thumb|250px|Predominance diagram for chromate]]
As mentioned above, a condensation reaction is also an acid–base reaction. In many systems, both protonation and condensation reactions can occur. The case of the chromate ion provides a relatively simple example. In the [[predominance diagram]] for chromate, shown at the right, pCr stands for the negative [[logarithm]] of the chromium concentration and [[pH]] stands for the negative logarithm of {{chem2|H+}}&nbsp;ion concentration. There are two independent equilibria. [[Equilibrium constants]] are defined as follows.<ref>{{cite journal|last=Brito|first=F. |author2=Ascanioa, J. |author3=Mateoa, S. |author4=Hernándeza, C. |author5=Araujoa, L. |author6=Gili, P. |author7=Martín-Zarzab, P. |author8=Domínguez, S. |author9=Mederos, A.|year=1997|title=Equilibria of chromate(VI) species in acid medium and ab initio studies of these species |journal=Polyhedron|volume=16|issue=21|pages=3835–3846 |doi=10.1016/S0277-5387(97)00128-9   }}</ref>

:{|
|<chem>CrO4^2- + H+ <=>  HCrO4-</chem>
|{{spaces|8}}<math>K_1=\frac{[\mathrm{HCrO_4^-}]}{[\mathrm{CrO_4^{2-}}][\mathrm{H^+}]}</math>
|{{spaces|8}}<math>\log K_1 = 5.89</math>
|-
|<chem>2 HCrO4- <=> Cr2O7^2- + H2O</chem>
|{{spaces|8}}<math>K_2=\frac{[\mathrm{Cr_2O_7^{2-}}]}{[\mathrm{HCrO_4^-}]^2}</math>
|{{spaces|8}}<math>\log K_2 = 2.05</math>
|}

The predominance diagram is interpreted as follows. 
*The chromate ion, {{chem2|CrO4(2-)}}, is the predominant species at high pH.  As pH rises the chromate ion becomes ever more predominant, until it is the only species in solutions with pH&nbsp;>&nbsp;6.75.
*At pH&nbsp;<&nbsp;p''K''<sub>1</sub> the hydrogen chromate ion, {{chem2|HCrO4-}} is predominant in dilute solution.
*The dichromate ion, {{chem2|Cr2O7(2-)}}, is predominant in more concentrated solutions, except at high pH.

The species {{chem2|H2CrO4}} and  {{chem2|HCr2O7-}} are not shown as they are formed only at very low pH. 

Predominance diagrams can become very complicated when many polymeric species can be formed,<ref>{{cite book|last=Pope|first=M.T.|title=Heteropoly and Isopoly Oxometalates |publisher=Springer |year=1983 |isbn=0-387-11889-6 }}</ref> such as in [[vanadate]]s, [[molybdate]]s, and [[tungstate]]s. Another complication is that many of the higher polymers are formed extremely slowly, such that equilibrium may not be attained even in months, leading to possible errors in the equilibrium constants and the predominance diagram.

== See also ==
* [[Oxycation]]
* [[Fluoroanion]]
* [[Carbanion]]

==References and notes==
=== Notes ===
<references group="Note" />
=== References ===
{{Reflist}}

==External links==
*{{Commonscatinline|Oxoanions}}

{{chlorates}}
{{sulfates}}
{{nitrates}}
{{phosphates}}
{{perchlorates}}
{{bromates}}
{{carbonates}}

[[Category:Oxyanions| ]]
[[Category:Acid–base chemistry]]
[[Category:Equilibrium chemistry]]