Oxide: Difference between revisions
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{{Short description|Chemical compound | {{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.]] | [[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> | 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== | ==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 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 | 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> | |journal=Chem. Rev.|year=2003|volume=103|issue=7 |pages=2579–2596|pmid=12848580 }}</ref> | ||
For simplicity sake, most of this article focuses on binary oxides. | For simplicity's sake, most of this article focuses on binary oxides. | ||
==Formation== | ==Formation== | ||
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===Metal oxides=== | ===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--> | 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> | 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> | 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/> | [[Noble metal]]s (such as [[gold]] and [[platinum]]) are prized because they resist direct chemical combination with oxygen.<ref name=Greenwood/> | ||
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===Non-metal oxides=== | ===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/> | 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: | 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: | 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> | 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> | 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]]: | Finally the trioxide is converted to sulfuric acid by a [[hydration reaction]]: | ||
: | :{{chem2|SO3 + H2O -> H2SO4}} | ||
==Structure== | ==Structure== | ||
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{{see also|Carbothermic 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 °C:<ref>{{Cite web|url=http://chemister.ru/Database/properties-en.php?dbid=1&id=4098|title = Silver oxide}}</ref> | 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 °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/> | 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> | 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=== | ===Hydrolysis and dissolution=== | ||
Because the | 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/> | 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/> | ||
Latest revision as of 18:13, 5 November 2025
Template:Short descriptionScript error: No such module "For".
An oxide (Template:IPAc-en) is a chemical compound containing at least one oxygen atom and one other element[1] in its chemical formula. "Oxide" itself is the dianion (anion bearing a net charge of −2) of oxygen, an O2− 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 Template:Chem2 (called a passivation layer) that protects the foil from further oxidation.[2]
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.[2] 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, Fe2SiO4, the mineral fayalite, is one of many examples of a ternary oxide. For many metal oxides, the possibilities of polymorphism and nonstoichiometry exist as well.[3] 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.[4]
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. carbonates, hydroxides, and nitrates. In the making of calcium oxide, calcium carbonate (limestone) breaks down upon heating, releasing carbon dioxide:[2]
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:[5]
The production of metals from ores often involves the production of oxides by roasting (heating) metal sulfide minerals in air. In this way, Template:Chem2 (molybdenite) is converted to molybdenum trioxide, the precursor to virtually all molybdenum compounds:[6]
Noble metals (such as gold and platinum) are prized because they resist direct chemical combination with oxygen.[2]
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:[2]
With excess oxygen, the dioxide is the product, the pathway proceeds by the intermediacy of carbon monoxide:
Elemental nitrogen (Template:Chem2) is difficult to convert to oxides, but the combustion of ammonia gives nitric oxide, which further reacts with oxygen:
These reactions are practiced in the production of nitric acid, a commodity chemical.[7]
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:[8]
Finally the trioxide is converted to sulfuric acid by a hydration reaction:
Structure
Oxides have a range of structures, from individual molecules to polymeric and crystalline structures. At standard conditions, oxides may range from solids to gases. Solid oxides of metals usually have polymeric structures at ambient conditions.[9]
Molecular oxides
- Some important gaseous oxides
-
Carbon dioxide is the main product of fossil fuel combustion.
-
Carbon monoxide is the product of the incomplete combustion of carbon-based fuels and a precursor to many useful chemicals.
-
Nitrogen dioxide is a problematic pollutant from internal combustion engines.
-
Sulfur dioxide, the principal oxide of sulfur, is emitted from volcanoes.
-
Nitrous oxide ("laughing gas") is a potent greenhouse gas produced by soil bacteria.
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., NO, N2O, NO2 and N2O4. Phosphorus pentoxide is a more complex molecular oxide with a deceptive name, the real formula being P4O10. Tetroxides are rare, with a few more common examples being ruthenium tetroxide, osmium tetroxide, and xenon tetroxide.[2]
Reactions
Reduction
Script error: No such module "Labelled list hatnote". 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 °C:[10]
Most often, however, metal oxides are reduced by a chemical reagent. A common and cheap reducing agent is carbon in the form of coke. The most prominent example is that of iron ore smelting. Many reactions are involved, but the simplified equation is usually shown as:[2]
Some metal oxides dissolve in the presence of reducing agents, which can include organic compounds. Reductive dissolution of ferric oxides is integral to geochemical phenomena such as the iron cycle.[11]
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.[2]
Dissolution of oxides often gives oxyanions. Adding aqueous base to Template:Chem2 gives various phosphates. Adding aqueous base to Template:Chem2 gives polyoxometalates. Oxycations are rarer, some examples being nitrosonium (Template:Chem2), vanadyl (Template:Chem2), and uranyl (Template:Chem2). Many compounds are known with both oxides and other groups. In organic chemistry, these include ketones and many related carbonyl compounds. For the transition metals, many oxo complexes are known as well as oxyhalides.[2]
Nomenclature and formulas
The chemical formulas of the oxides of the chemical elements in their highest oxidation state are predictable and are derived from the number of valence electrons for that element. Even the chemical formula of O4, 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 F2O7—but as OF2.[12]
See also
- Other oxygen ions: ozonide (Template:Chem2), superoxide (Template:Chem2), peroxide (Template:Chem2) and dioxygenyl (Template:Chem2).
- Suboxide
- Oxohalide
- Oxyanion
- Complex oxide
- See Template:Oxides for a list of oxides.
- Salt
- Wet electrons
References
Template:Oxides Template:Monatomic anion compounds
- ↑ Script error: No such module "citation/CS1".
- ↑ a b c d e f g h i Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. Template:ISBN.
- ↑ Script error: No such module "citation/CS1".
- ↑ Script error: No such module "Citation/CS1".
- ↑ Script error: No such module "citation/CS1".
- ↑ Template:Ullmann
- ↑ Template:Ullmann
- ↑ Template:Ullmann
- ↑ Script error: No such module "citation/CS1".
- ↑ Script error: No such module "citation/CS1".
- ↑ Script error: No such module "citation/CS1".
- ↑ Script error: No such module "Citation/CS1".