Oxide

 ( Anions ) 

An oxide () is a chemical compound containing at least one oxygen atom and one other chemical element

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 (called a passivation layer) that protects the foil from further oxidation.Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. .

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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.Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. . 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.

C. N. R. Rao, B. Raveau (1995). 9781560816478, VCH. ISBN 9781560816478

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.

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

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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.

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Metal oxides Many metal oxides arise by decomposition of other metal compounds, e.g. , , and . In the making of calcium oxide, calcium carbonate (limestone) breaks down upon heating, releasing carbon dioxide:

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:

(2025). 9783527306732 ISBN 9783527306732

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, (molybdenite) is converted to molybdenum trioxide, the precursor to virtually all molybdenum compounds:

2 MoS2 + 7 O2 -> 2MoO3 + 4 SO2

(such as gold and platinum) are prized because they resist direct chemical combination with oxygen.

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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 CH4 + 3 O2 -> 2 CO + 4 H2O

2 C + O2 -> 2 CO

With excess oxygen, the dioxide is the product, the pathway proceeds by the intermediacy of carbon monoxide:

CH4 + 2 O2 -> CO2 + 2 H2O

C + O2 -> CO2

Elemental nitrogen () is difficult to convert to oxides, but the combustion of ammonia gives nitric oxide, which further reacts with oxygen:

4 NH3 + 5 O2 -> 4 NO + 6 H2O

2 NO + O2 -> 2 NO2

These reactions are practiced in the production of nitric acid, a commodity chemical.

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:

S + O2 -> SO2

2 SO2 + O2 -> 2 SO3

Finally the trioxide is converted to sulfuric acid by a hydration reaction:

SO3 + H2O -> H2SO4

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Structure

Oxides have a range of structures, from individual molecules to and crystalline structures. At standard conditions, oxides may range from solids to gases. Solid oxides of metals usually have polymeric structures at ambient conditions.

P.A. Cox (2025). 9780199588947, Oxford University Press. ISBN 9780199588947

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Molecular 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.

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, Nitrous oxide, Nitrogen dioxide 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.

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Reactions

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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:

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. The most prominent example is that of iron ore smelting. Many reactions are involved, but the simplified equation is usually shown as:

2 Fe2O3 + 3 C -> 4 Fe + 3 CO2

Some dissolve in the presence of reducing agents, which can include organic compounds. Reductive dissolution of Ferric Oxide is integral to geochemical phenomena such as the iron cycle.

(2025). 9783527302741 ISBN 9783527302741

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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.

Dissolution of oxides often gives Oxyanion. Adding aqueous base to gives various . Adding aqueous base to gives . Oxycations are rarer, some examples being nitrosonium (), vanadyl (), and uranyl (). Many compounds are known with both oxides and other groups. In organic chemistry, these include and many related carbonyl compounds. For the transition metals, many Oxo ligand are known as well as .

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Nomenclature and formulas

The of the oxides of the in their highest oxidation state are predictable and are derived from the number of 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.

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See also

• Other oxygen ions: ozonide (), superoxide (), peroxide () and dioxygenyl ().
• Suboxide
• Oxohalide
• Oxyanion
• Complex oxide
• See Oxides for a list of oxides.
• Salt
• Wet electrons

Categories: Anions

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