View source for Oxide

{{Short description|Chemical compound}}{{For|negatively-charged polyatomic ion containing oxygen|Oxyanions}}
[[Image:Rutile-unit-cell-3D-balls.png|thumb|right|The [[Crystal structure#Unit cell|unit cell]] of [[rutile]], an important oxide of titanium. Ti(IV) centers are grey; oxygen centers are red. Notice that oxygen forms three bonds to titanium and titanium forms six bonds to oxygen.]]
An '''oxide''' ({{IPAc-en|ˈ|ɒ|k|s|aɪ|d}}) is a [[chemical compound]] containing at least one [[oxygen]] [[atom]] and one other [[chemical element|element]]<ref>{{cite book|title = Foundations of College Chemistry|edition = 12th |last1= Hein|first1= Morris|last2 = Arena|first2= Susan|date = 2006|publisher = Wiley |isbn = 978-0-471-74153-4}}</ref> in its [[chemical formula]]. "Oxide" itself is the [[dianion]] (anion bearing a net charge of −2) of oxygen, an O<sup>2−</sup> ion with oxygen in the [[oxidation state]] of −2. Most of the [[Earth's crust]] consists of oxides. Even materials considered pure elements often develop an oxide coating. For example, [[aluminium foil]] develops a thin skin of {{chem2|link=alumina|Al2O3}} (called a [[Passivation (chemistry)|passivation layer]]) that protects the foil from further [[oxidation]].<ref name=Greenwood>Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. {{ISBN|0-7506-3365-4}}.</ref>

==Stoichiometry==
Oxides are extraordinarily diverse in terms of [[stoichiometries]] (the measurable relationship between reactants and chemical equations of an equation or reaction) and in terms of the structures of each stoichiometry. Most elements form oxides of more than one stoichiometry. A well known example is [[carbon monoxide]] and [[carbon dioxide]].<ref name=Greenwood>Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. {{ISBN|0-7506-3365-4}}.</ref> This applies to ''binary'' oxides, that is, compounds containing only oxide and another element.  Far more common than binary oxides are oxides of more complex stoichiometries.  Such complexity can arise by the introduction of other cations (a positively charged ion, i.e. one that would be attracted to the cathode in electrolysis) or other anions (a negatively charged ion). [[Iron silicate]], Fe<sub>2</sub>SiO<sub>4</sub>, the mineral [[fayalite]], is one of many examples of a ternary oxide. For many metal oxides, the possibilities of polymorphism and [[Non-stoichiometric_compound|nonstoichiometry]] exist as well.<ref>{{cite book|title=Transition Metal Oxides|author=C. N. R. Rao, B. Raveau| publisher=VCH|location=New York|year=1995|isbn=1-56081-647-3}}</ref>  The commercially important dioxides of titanium exists in three distinct structures, for example.  Many metal oxides exist in various nonstoichiometric states.  Many molecular oxides exist with diverse ligands as well.<ref>{{cite journal |doi=10.1021/cr020376q|title=Organometallic Oxides of Main Group and Transition Elements Downsizing Inorganic Solids to Small Molecular Fragments|first1=Herbert W. |last1=Roesky|first2=Ionel |last2=Haiduc|first3=Narayan S. |last3=Hosmane
|journal=Chem. Rev.|year=2003|volume=103|issue=7 |pages=2579–2596|pmid=12848580 }}</ref>

For simplicity's sake, most of this article focuses on binary oxides.

==Formation==
Oxides are associated with all elements except a few noble gases. The pathways for the formation of this diverse family of compounds are correspondingly numerous.

===Metal oxides===
Many metal oxides arise by decomposition of other metal compounds, e.g. [[carbonate]]s, [[hydroxide]]s, and [[nitrate]]s.  In the making of [[calcium oxide]], [[calcium carbonate]] ([[limestone]]) breaks down upon heating, releasing carbon dioxide:<ref name=Greenwood/><!--p120-->
:{{chem2|CaCO3 -> CaO + CO2}}
The reaction of elements with oxygen in air is a key step in [[corrosion]] relevant to the commercial use of iron especially. Almost all elements form oxides upon heating with oxygen atmosphere. For example, zinc powder will burn in air to give [[zinc oxide]]:<ref>{{cite book |doi=10.1002/14356007.a28_509|chapter=Zinc |title=Ullmann's Encyclopedia of Industrial Chemistry |year=2000 |last1=Graf |first1=Günter G. |isbn=3-527-30673-0 }}</ref>
:{{chem2|2 Zn + O2 -> 2 ZnO}}
The production of metals from ores often involves the production of oxides by roasting (heating) metal sulfide minerals in air. In this way, {{chem2|MoS2}} ([[molybdenite]]) is converted to [[molybdenum trioxide]], the precursor to virtually all molybdenum compounds:<ref>{{Ullmann|author=Roger F. Sebenik|display-authors=etal|title=Molybdenum and Molybdenum Compounds|year=2005|doi=10.1002/14356007.a16_655|isbn=978-3527306732}}</ref>
:{{chem2|2 MoS2 + 7 O2 -> 2 MoO3 + 4 SO2}}

[[Noble metal]]s (such as [[gold]] and [[platinum]]) are prized because they resist direct chemical combination with oxygen.<ref name=Greenwood/>

===Non-metal oxides===
Important and prevalent nonmetal oxides are [[carbon dioxide]] and [[carbon monoxide]].  These species form upon full or partial oxidation of carbon or hydrocarbons.  With a deficiency of oxygen, the monoxide is produced:<ref name=Greenwood/>
:{{chem2|2 CH4 + 3 O2 -> 2 CO + 4 H2O}}
:{{chem2|2 C + O2 -> 2 CO}}
With excess oxygen, the dioxide is the product, the pathway proceeds by the intermediacy of carbon monoxide:
:{{chem2|CH4 + 2 O2 -> CO2 + 2 H2O}}
:{{chem2|C + O2 -> CO2}}

Elemental nitrogen ({{chem2|N2}}) is difficult to convert to oxides, but the combustion of ammonia gives [[nitric oxide]], which further reacts with oxygen:
:{{chem2|4 NH3 + 5 O2 -> 4 NO + 6 H2O}}
:{{chem2|2 NO + O2 -> 2 NO2}}
These reactions are practiced in the production of [[nitric acid]], a commodity chemical.<ref name=Ullmann>{{Ullmann|last1=Thiemann |first1=Michael |last2=Scheibler |first2=Erich |last3=Wiegand |first3=Karl Wilhelm |date=2000 |title=Nitric Acid, Nitrous Acid, and Nitrogen Oxides |doi=10.1002/14356007.a17_293|isbn=978-3527306732 }}</ref>

The chemical produced on the largest scale industrially is [[sulfuric acid]].  It is produced by the oxidation of sulfur to [[sulfur dioxide]], which is separately oxidized to [[sulfur trioxide]]:<ref>{{Ullmann|doi=10.1002/14356007.a25_635|title=Sulfuric Acid and Sulfur Trioxide |year=2000 |last1=Müller |first1=Hermann |isbn=3527306730 }}</ref> 
:{{chem2|S + O2 -> SO2}}
:{{chem2|2 SO2 + O2 -> 2 SO3}}
Finally the trioxide is converted to sulfuric acid by a [[hydration reaction]]:
:{{chem2|SO3 + H2O -> H2SO4}}

==Structure==
Oxides have a range of structures, from individual molecules to [[polymer]]ic and [[crystalline]] structures.  At standard conditions, oxides may range from solids to gases. Solid oxides of metals usually have polymeric structures at ambient conditions.<ref>{{cite book|title=Transition Metal Oxides. An Introduction to Their Electronic Structure and Properties|author=P.A. Cox |publisher=Oxford University Press|year=2010|isbn=978-0-19-958894-7}}</ref> 

===Molecular oxides===
<gallery widths="140" heights="100" caption="Some important gaseous oxides">
File:Carbon-dioxide-2D-dimensions.svg|[[Carbon dioxide]] is the main product of fossil fuel combustion.
File:Carbon monoxide 2D.svg|[[Carbon monoxide]] is the product of the incomplete combustion of carbon-based fuels and a precursor to many useful chemicals.
File:Nitrogen-dioxide-2D-dimensions-vector.svg|[[Nitrogen dioxide]] is a problematic pollutant from internal combustion engines.
File:Sulfur-dioxide-2D.svg|[[Sulfur dioxide]], the principal oxide of [[sulfur]], is emitted from volcanoes.
File:Nitrous-oxide-2D-dimensions.png|[[Nitrous oxide]] ("laughing gas") is a potent greenhouse gas produced by soil bacteria. 
</gallery>

Although most metal oxides are crystalline solids, many non-metal oxides are molecules.  Examples of molecular oxides are [[carbon dioxide]] and [[carbon monoxide]]. All simple oxides of nitrogen are molecular, e.g., [[Nitric oxide|NO]], [[Nitrous oxide|N<sub>2</sub>O]], [[Nitrogen dioxide|NO<sub>2</sub>]] and [[Dinitrogen tetroxide|N<sub>2</sub>O<sub>4</sub>]]. [[Phosphorus pentoxide]] is a more complex molecular oxide with a deceptive name, the real formula being P<sub>4</sub>O<sub>10</sub>. Tetroxides are rare, with a few more common examples being  [[ruthenium tetroxide]], [[osmium tetroxide]], and [[xenon tetroxide]].<ref name=Greenwood/>

==Reactions==
===Reduction===
{{see also|Carbothermic reduction}}
Reduction of metal oxide to the metal is practiced on a large scale in the production of some metals.  Many metal oxides convert to metals simply by heating ([[thermal decomposition]]). For example, [[silver oxide]] decomposes at 200&nbsp;°C:<ref>{{Cite web|url=http://chemister.ru/Database/properties-en.php?dbid=1&id=4098|title = Silver oxide}}</ref>
:{{chem2|2 Ag2O -> 4 Ag + O2}}
Most often, however, metal oxides are reduced by a chemical reagent.  A common and cheap reducing agent is carbon in the form of [[coke (fuel)|coke]].  The most prominent example is that of [[Iron#Industrial production|iron ore smelting]]. Many reactions are involved, but the simplified equation is usually shown as:<ref name=Greenwood/>
:{{chem2|2 Fe2O3 + 3 C -> 4 Fe + 3 CO2}}

Some [[metal oxide]]s dissolve in the presence of reducing agents, which can include organic compounds.  Reductive dissolution of [[Ferric Oxide|ferric oxides]] is integral to [[geochemical]] phenomena such as the [[iron cycle]].<ref name="CornellSchwertmann2003-323">{{cite book|title=The Iron Oxides: Structure, Properties, Reactions, Occurrences and Uses, Second Edition|last1=Cornell|first1=R. M.|last2=Schwertmann|first2=U.|year=2003|page=323|doi=10.1002/3527602097|isbn=978-3-527-30274-1}}</ref>

===Hydrolysis and dissolution===
Because the M–O bonds are typically strong, metal oxides tend to be insoluble in solvents, though they may be attacked by aqueous acids and bases.<ref name=Greenwood/>

Dissolution of oxides often gives [[Oxyanion|oxyanions]].  Adding aqueous base to {{chem2|P4O10}} gives various [[phosphate]]s.  Adding aqueous base to {{chem2|MoO3}} gives [[polyoxometalate]]s.  Oxycations are rarer, some examples being [[nitrosonium]] ({{chem2|NO+}}), [[vanadyl]] ({{chem2|VO(2+)}}), and [[uranyl]] ({{chem2|UO2(2+)}}). Many compounds are known with both oxides and other groups.  In [[organic chemistry]], these include [[ketone]]s and many related [[carbonyl]] compounds. For the transition metals, many [[Oxo ligand|oxo complex]]es are known as well as [[oxyhalide]]s.<ref name=Greenwood/>

==Nomenclature and formulas==
The [[chemical formula]]s of the oxides of the [[chemical element]]s in their highest [[oxidation state]] are predictable and are derived from the number of [[valence electron]]s for that element. Even the chemical formula of O<sub>4</sub>, [[tetraoxygen]], is predictable as a [[group 16 element]]. One exception is [[copper]], for which the highest oxidation state oxide is [[copper(II) oxide]] and not [[copper(I) oxide]]. Another exception is [[fluoride]], which does not exist as one might expect—as F<sub>2</sub>O<sub>7</sub>—but as [[oxygen difluoride|OF<sub>2</sub>]].<ref>{{cite journal|doi=10.1021/ed082p1649 |title=Fully Exploiting the Potential of the Periodic Table through Pattern Recognition|author=Schultz, Emeric|journal= J. Chem. Educ. |year=2005|volume=82 |issue=11|page=1649|bibcode = 2005JChEd..82.1649S }}</ref>

== See also ==
{{Wiktionary}}
* Other oxygen ions: [[ozonide]] ({{chem2|O3-}}), [[superoxide]] ({{chem2|O2-}}), [[peroxide]] ({{chem2|O2(2-)}}) and [[dioxygenyl]] ({{chem2|O2+}}).
* [[Suboxide]]
* [[Oxohalide]]
* [[Oxyanion]]
* [[Complex oxide]]
* See [[:Template:Oxides]] for a list of oxides.
* [[Salt (chemistry)|Salt]]
* [[Wet electrons]]

== References ==
{{reflist}}

{{Oxides}}
{{Monatomic anion compounds}}

{{Authority control}}

[[Category:Oxides| ]]
[[Category:Anions]]