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The unit cell of rutile. Ti(IV) centers are grey; oxygen centers are red. Notice that oxygen forms three bonds to titanium and titanium forms six bonds to oxygen. Rutile-unit-cell-3D-balls.png
The unit cell of rutile. 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 ( /ˈɒksd/ ) is a chemical compound that contains at least one oxygen atom and one other element [1] in its chemical formula. "Oxide" itself is the dianion of oxygen, an O2– (molecular) ion. Metal oxides thus typically contain an anion of oxygen in the oxidation state of −2. Most of the Earth's crust consists of solid oxides, the result of elements being oxidized by the oxygen in air or in water. Even materials considered pure elements often develop an oxide coating. For example, aluminium foil develops a thin skin of Al2O3 (called a passivation layer) that protects the foil from further corrosion. [2] Certain elements can form multiple oxides, differing in the amounts of the element combining with the oxygen. Examples are carbon, iron, nitrogen (see nitrogen oxide), silicon, titanium, and aluminium. In such cases the oxides are distinguished by specifying the numbers of atoms involved, as in carbon monoxide and carbon dioxide, or by specifying the element's oxidation number, as in iron(II) oxide and iron(III) oxide.



Due to its electronegativity, oxygen forms stable chemical bonds with almost all elements to give the corresponding oxides. Noble metals (such as gold or platinum) are prized because they resist direct chemical combination with oxygen, and substances like gold(III) oxide must be generated by indirect routes.

Two independent pathways for corrosion of elements are hydrolysis and oxidation by oxygen. The combination of water and oxygen is even more corrosive. Virtually all elements burn in an atmosphere of oxygen or an oxygen-rich environment. In the presence of water and oxygen (or simply air), some elements— sodium—react rapidly, to give the hydroxides. In part, for this reason, alkali and alkaline earth metals are not found in nature in their metallic, i.e., native, form. Cesium is so reactive with oxygen that it is used as a getter in vacuum tubes, and solutions of potassium and sodium, so-called NaK are used to deoxygenate and dehydrate some organic solvents. The surface of most metals consists of oxides and hydroxides in the presence of air. A well-known example is aluminium foil, which is coated with a thin film of aluminium oxide that passivates the metal, slowing further corrosion. The aluminum oxide layer can be built to greater thickness by the process of electrolytic anodizing. Though solid magnesium and aluminum react slowly with oxygen at STP—they, like most metals, burn in air, generating very high temperatures. Finely grained powders of most metals can be dangerously explosive in air. Consequently, they are often used in solid-fuel rockets.

Oxides, such as iron(III) oxide or rust, which consists of hydrated iron(III) oxides Fe2O3*nH2O and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3), form when oxygen combines with other elements Rust screw.jpg
Oxides, such as iron(III) oxide or rust, which consists of hydrated iron(III) oxides Fe2O3·nH2O and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3), form when oxygen combines with other elements

In dry oxygen, iron readily forms iron(II) oxide, but the formation of the hydrated ferric oxides, Fe2O3−x(OH)2x, that mainly comprise rust, typically requires oxygen and water. Free oxygen production by photosynthetic bacteria some 3.5 billion years ago precipitated iron out of solution in the oceans as Fe2O3 in the economically important iron ore hematite.


Oxides have a range of different structures, from individual molecules to polymeric and crystalline structures. At standard conditions, oxides may range from solids to gases.

Oxides of metals

Oxides of most metals adopt polymeric structures. [3] The oxide typically links three metal atoms (e.g., rutile structure) or six metal atoms (carborundum or rock salt structures). Because the M-O bonds are typically strong and these compounds are crosslinked polymers, the solids tend to be insoluble in solvents, though they are attacked by acids and bases. The formulas are often deceptively simple where many are nonstoichiometric compounds. [2]

Molecular oxides

Although most metal oxides are polymeric, some 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. Some polymeric oxides depolymerize when heated to give molecules, examples being selenium dioxide and sulfur trioxide. Tetroxides are rare. The more common examples: ruthenium tetroxide, osmium tetroxide, and xenon tetroxide.

Many oxyanions are known, such as polyphosphates and polyoxometalates. Oxycations are rarer, some examples being nitrosonium (NO+), vanadyl (VO2+), and uranyl (UO2+
). Of course 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.


Conversion of a metal oxide to the metal is called reduction. The reduction can be induced by many reagents. Many metal oxides convert to metals simply by heating.

Reduction by carbon

Metals are "won" from their oxides by chemical reduction, i.e. by the addition of 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]

2 Fe2O3 + 3 C → 4 Fe + 3 CO2

Metal oxides can be reduced by organic compounds. This redox process is the basis for many important transformations in chemistry, such as the detoxification of drugs by the P450 enzymes and the production of ethylene oxide, which is converted to antifreeze. In such systems, the metal center transfers an oxide ligand to the organic compound followed by regeneration of the metal oxide, often by oxygen in the air.

Reduction by heating

Metals that are lower in the reactivity series can be reduced by heating alone. For example, silver oxide decomposes at 200 °C: [4]

2 Ag2O → 4 Ag + O2

Reduction by displacement

Metals that are more reactive displace the oxide of the metals that are less reactive. For example, zinc is more reactive than copper, so it displaces copper (II) oxide to form zinc oxide:

Zn + CuO → ZnO + Cu

Reduction by hydrogen

Apart from metals, hydrogen can also displace metal oxides to form hydrogen oxide, also known as water:

H2 + CuO → Cu + H2O

Reduction by electrolysis

Since metals that are reactive form oxides that are stable, some metal oxides must be electrolyzed to be reduced. This includes sodium oxide, potassium oxide, calcium oxide, magnesium oxide, and aluminium oxide. The oxides must be molten before immersing graphite electrodes in them:

2Al2O3 → 4Al + 3O2

Hydrolysis and dissolution

Oxides typically react with acids or bases, sometimes both. Those reacting only with acids are labeled basic oxides. Those reacting only by bases are called "acidic oxides". Oxides that react with both are amphoteric. Metals tend to form basic oxides, non-metals tend to form acidic oxides, and amphoteric oxides are formed by elements near the boundary between metals and non-metals (metalloids). This reactivity is the basis of many practical processes, such as the extraction of some metals from their ores in the process called hydrometallurgy.

Oxides of more electropositive elements tend to be basic. They are called basic anhydrides. Exposed to water, they may form basic hydroxides. For example, sodium oxide is basic—when hydrated, it forms sodium hydroxide. Oxides of more electronegative elements tend to be acidic. They are called "acid anhydrides"; adding water, they form oxoacids. For example, dichlorine heptoxide is an acid anhydride; perchloric acid is its fully hydrated form. Some oxides can act as both acid and base. They are amphoteric. An example is aluminium oxide. Some oxides do not show behavior as either acid or base.

The oxide ion has the formula O2−. It is the conjugate base of the hydroxide ion, OH and is encountered in ionic solids such as calcium oxide. O2− is unstable in aqueous solution − its affinity for H+ is so great (pKb ~ −38) that it abstracts a proton from a solvent H2O molecule:

O2− + H2O → 2 OH

The equilibrium constant of aforesaid reactions is pKeq ~ −22

In the 18th century, oxides were named calxes or calces after the calcination process used to produce oxides. Calx was later replaced by oxyd.

Reductive dissolution

The reductive dissolution of a transition metal oxide occurs when dissolution is coupled to a redox event. [5] For example, ferric oxides dissolve in the presence of reductants, which can include organic compounds. [6] or bacteria [7] Reductive dissolution is integral to geochemical phenomena such as the iron cycle. [8]

Reductive dissolution does not necessarily occur at the site where the reductant adsorbs. Instead, the added electron travel through the particle, causing reductive dissolution elsewhere on the particle. [9] [10]

Nomenclature and formulas

Sometimes, metal-oxygen ratios are used to name oxides. Thus, NbO would be called niobium monoxide and TiO2 is titanium dioxide. This naming follows the Greek numerical prefixes. In the older literature and continuing in industry, oxides are named by adding the suffix -a to the element's name. Hence alumina, magnesia and chromia, are, respectively, Al2O3, MgO and Cr2O3.

Special types of oxides are peroxide, O2−
, and superoxide, O
. In such species, oxygen is assigned higher (less negative) oxidation states than oxide.

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. [11]

Since fluorine is more electronegative than oxygen, oxygen difluoride (OF2) does not represent an oxide of fluorine, but instead represents a fluoride of oxygen.

Examples of oxides

The following table gives examples of commonly encountered oxides. Only a few representatives are given, as the number of polyatomic ions encountered in practice is very large.

Water (hydrogen oxide) H
Common solvent, required by carbon-based life
Nitrous oxide N
Laughing gas, anesthetic (used in a combination with diatomic oxygen (O2) to make nitrous oxide and oxygen anesthesia), produced by nitrogen-fixing bacteria, nitrous, oxidizing agent in rocketry, aerosol propellant, recreational drug, greenhouse gas. Other nitrogen oxides such as NO
(nitrogen dioxide), NO (nitrogen oxide), N
(dinitrogen trioxide) and N
(dinitrogen tetroxide) exist, particularly in areas with notable air pollution. They are also strong oxidisers, can add nitric acid to acid rain, and are harmful to health.
Silicon dioxide SiO
Sand, quartz
Iron(II,III) oxide Fe
Iron ore, rust, along with iron(III) oxide (Fe
Aluminium oxide Al
Aluminium ore, alumina, corundum, ruby (corundum with impurities of chromium).
Zinc oxide ZnORequired for vulcanization of rubber, additive to concrete, sunscreen, skin care lotions, antibacterial and antifungal properties, food additive, white pigment.
Carbon dioxide CO
Constituent of the atmosphere of Earth, the most abundant and important greenhouse gas, used by plants in photosynthesis to make sugars, product of biological processes such as respiration and chemical reactions such as combustion and chemical decomposition of carbonates. CO or Carbon monoxide exists as a product of incomplete combustion and is a highly toxic gas.
Calcium oxide CaO Quicklime (used in construction to make mortar and concrete), used in self-heating cans due to exothermic reaction with water to produce calcium hydroxide, possible ingredient in Greek fire and produces limelight when heated over 2,400 °Celsius.

See also

Related Research Articles

Chlorine Chemical element, symbol Cl and atomic number 17

Chlorine is a chemical element with the symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity on the Pauling scale, behind only oxygen and fluorine. Nitrogen is also more electronegative in many other scales such as Mulliken-Jaffe, Allred-Rochow, Noorizadeh-Shakerzadeh, Nagle, Martynov-Batsanov and Allen electronegativity scales.

Carboxylic acid Organic compound

A carboxylic acid is an organic acid that contains a carboxyl group (C(=O)OH) attached to an R-group. The general formula of a carboxylic acid is R−COOH or R−CO2H, with R referring to the alkyl, alkenyl, aryl, or other group. Carboxylic acids occur widely. Important examples include the amino acids and fatty acids. Deprotonation of a carboxylic acid gives a carboxylate anion.

Hydroxide Chemical compound

Hydroxide is a diatomic anion with chemical formula OH. It consists of an oxygen and hydrogen atom held together by a single covalent bond, and carries a negative electric charge. It is an important but usually minor constituent of water. It functions as a base, a ligand, a nucleophile, and a catalyst. The hydroxide ion forms salts, some of which dissociate in aqueous solution, liberating solvated hydroxide ions. Sodium hydroxide is a multi-million-ton per annum commodity chemical. A hydroxide attached to a strongly electropositive center may itself ionize, liberating a hydrogen cation (H+), making the parent compound an acid.

Hydrolysis is any chemical reaction in which a molecule of water breaks one or more chemical bonds. The term is used broadly for substitution, elimination, and solvation reactions in which water is the nucleophile.

Nitrogen Chemical element, symbol N and atomic number 7

Nitrogen is the chemical element with the symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. The name nitrogène was suggested by French chemist Jean-Antoine-Claude Chaptal in 1790 when it was found that nitrogen was present in nitric acid and nitrates. Antoine Lavoisier suggested instead the name azote, from the Ancient Greek: ἀζωτικός "no life", as it is an asphyxiant gas; this name is used instead in many languages, such as French, Italian, Russian, Romanian, Portuguese and Turkish, and appears in the English names of some nitrogen compounds such as hydrazine, azides and azo compounds.

Rust Type of iron oxide

Rust is an iron oxide, a usually reddish-brown oxide formed by the reaction of iron and oxygen in the catalytic presence of water or air moisture. Rust consists of hydrous iron(III) oxides (Fe2O3·nH2O) and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3), and is typically associated with the corrosion of refined iron.

Silicon Chemical element, symbol Si and atomic number 14

Silicon is a chemical element with the symbol Si and atomic number 14. It is a hard, brittle crystalline solid with a blue-grey metallic lustre, and is a tetravalent metalloid and semiconductor. It is a member of group 14 in the periodic table: carbon is above it; and germanium, tin, lead, and flerovium are below it. It is relatively unreactive. Because of its high chemical affinity for oxygen, it was not until 1823 that Jöns Jakob Berzelius was first able to prepare it and characterize it in pure form. Its oxides form a family of anions known as silicates. Its melting and boiling points of 1414 °C and 3265 °C respectively are the second highest among all the metalloids and nonmetals, being only surpassed by boron. Silicon is the eighth most common element in the universe by mass, but very rarely occurs as the pure element in the Earth's crust. It is most widely distributed in space in cosmic dusts, planetoids, and planets as various forms of silicon dioxide (silica) or silicates. More than 90% of the Earth's crust is composed of silicate minerals, making silicon the second most abundant element in the Earth's crust, after oxygen. Silicon is a natural element, and when not previously present has a residence time of about 400 years in the world's oceans.

Sulfuric acid Chemical compound

Sulfuric acid (American spelling) or sulphuric acid (Commonwealth spelling), also known as oil of vitriol, is a mineral acid composed of the elements sulfur, oxygen and hydrogen, with the molecular formula H2SO4. It is a colorless, odorless and viscous liquid that is miscible with water.

Electrolysis Technique in chemistry and manufacturing

In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity".

Sodium hydroxide Chemical compound with formula NaOH

Sodium hydroxide, also known as lye and caustic soda, is an inorganic compound with the formula NaOH. It is a white solid ionic compound consisting of sodium cations Na+
and hydroxide anions OH

Redox Chemical reaction in which oxidation states of atoms are changed

Redox is a type of chemical reaction in which the oxidation states of atoms are changed. Redox reactions are characterized by the actual or formal transfer of electrons between chemical species, most often with one species undergoing oxidation while another species undergoes reduction. The chemical species from which the electron is removed is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. In other words:

Weathering Breaking down of rocks or other materials through exposure to the elements

Weathering is the breaking down of rocks, soils and minerals as well as wood and artificial materials through contact with water, atmospheric gases, and biological organisms. Weathering occurs in situ, and should not be confused with erosion, which involves the transport of rocks and minerals by agents such as water, ice, snow, wind, waves and gravity.

Base (chemistry) Type of chemical substance

In chemistry, there are three definitions in common use of the word base, known as Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances which react with acids as originally proposed by G.-F. Rouelle in the mid-18th century.

Corrosion Gradual destruction of materials by chemical reaction with its environment

Corrosion is a natural process that converts a refined metal into a more chemically stable form such as oxide, hydroxide, carbonate or sulfide. It is the gradual destruction of materials by chemical and/or electrochemical reaction with their environment. Corrosion engineering is the field dedicated to controlling and preventing corrosion.

A reducing agent is an element or compound that loses or "donates" an electron to an electron recipient in a redox chemical reaction. The reducer's oxidation state increases while the oxidizer's decreases; this is expressed by saying that reducers "undergo oxidation" and "are oxidized" while oxidizers "undergo reduction" and "are reduced". Thus, reducing agents "reduce" oxidizers by decreasing their oxidation state while oxidizing agents "oxidize" reducers by increasing their oxidation state.

A period 3 element is one of the chemical elements in the third row of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when the periodic table skips a row and a chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The third period contains eight elements: sodium, magnesium, aluminium, silicon, phosphorus, sulfur, chlorine, and argon. The first two, sodium and magnesium, are members of the s-block of the periodic table, while the others are members of the p-block. All of the period 3 elements occur in nature and have at least one stable isotope.

An acidic oxide is an oxide that either produces an acidic solution upon addition to water, or acts as an acceptor of hydroxide ions effectively functioning as a lewis acid. Acidic oxides will typically have a low pKa and may be inorganic or organic. A commonly encountered acidic oxide, carbon dioxide produces an acidic solution when dissolved.

Pyrometallurgy is a branch of extractive metallurgy. It consists of the thermal treatment of minerals and metallurgical ores and concentrates to bring about physical and chemical transformations in the materials to enable recovery of valuable metals. Pyrometallurgical treatment may produce products able to be sold such as pure metals, or intermediate compounds or alloys, suitable as feed for further processing. Examples of elements extracted by pyrometallurgical processes include the oxides of less reactive elements like iron, copper, zinc, chromium, tin, and manganese.

A combination reaction (also known as a synthesis reaction) is a reaction where two or more elements or compounds (reactants) combine to form a single compound (product). Such reactions are represented by equations of the following form: X + Y → XY.The combination of two or more elements to form one compound is called a combination reaction. In other words, when two or more elements or compounds react so as to form one single compound, then the chemical reaction that takes place is called a combination reaction. | a)- Between elements | C + O2 → CO2 | Carbon completely burnt in oxygen yields carbon dioxide |- | b) Between compounds | CaO + H2O → Ca(OH)2 | Calcium oxide (lime) combined with water gives calcium hydroxide (slaked lime) |- | c) Between elements and compounds | 2CO + O2 → 2CO2 | Oxygen combines with carbon monoxide,And carbon dioxide is formed. |}

Oxygen compounds

The oxidation state of oxygen is −2 in almost all known compounds of oxygen. The oxidation state −1 is found in a few compounds such as peroxides. Compounds containing oxygen in other oxidation states are very uncommon: −12 (superoxides), −13 (ozonides), 0, +12 (dioxygenyl), +1, and +2.


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