oxide, chemical compound containing oxygen and one other chemical element. Oxides are widely and abundantly distributed in nature. Water is the oxide of hydrogen. Silicon dioxide is the major component of sand and quartz. Carbon dioxide is given off during respiration by animals and plants. Carbon monoxide, sulfur dioxide, and oxides of nitrogen are among the waste gases of gasoline-burning internal-combustion engines. Nitrous oxide is an oxide of nitrogen often called laughing gas. Many of the metals form oxides. Some metal oxides, e.g., those of iron, aluminum, tin, and zinc, are important as ores. Litharge and red lead are lead oxides used as pigments in paint. A number of elements, e.g., arsenic, carbon, manganese, nitrogen, phosphorous, and sulfur, combine with oxygen to form more than one oxide. The inert gases do not form oxides. The halogens and inactive metals do not combine directly with oxygen, but their oxides can be formed by indirect methods. Oxides are usually named according to the number of oxygen atoms present in a molecule, e.g., monoxide (or simply oxide), dioxide, trioxide. In a molecule of carbon monoxide, CO, for example, there is one oxygen atom; in carbon dioxide, CO₂, there are two; and in phosphorus pentoxide, P₂O₅, there are five. Oxides are commonly classified as acidic or basic oxides or anhydrides. Sulfur trioxide is an acid anhydride; it reacts with water to form sulfuric acid. Phosphorus pentoxide reacts vigorously with water to form phosphoric acid. Many metal oxides react with water to form alkaline hydroxides, e.g., calcium oxide (lime) reacts with water to form calcium hydroxide (slaked lime). Some metal oxides do not react with water but are basic in that they react with an acid to form a salt and water. Others exhibit amphoterism; i.e., they react with both acids and bases. Still others are neutral and nonreactive.

Any naturally occurring inorganic compound with a structure based on close-packed oxygen atoms in which smaller, positively charged metal or other ions occur. Oxide minerals are common in all rock types, whether igneous, sedimentary, or metamorphic.

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Any of a large and important class of chemical compounds in which oxygen is combined with another element. Metal oxides contain a metal cation and the oxide anion (O₂⁻); they typically react with water to form bases or with acids to form salts. Oxides of nonmetallic elements are volatile compounds in which a covalent bond joins the oxygen and the nonmetal; they react with water to form acids or with bases to form salts. A few substances (e.g., aluminum, zinc) form amphoteric oxides, which form salts with both acids and bases. Certain organic compounds form oxides in which the oxygen is covalently bonded to an atom of nitrogen (amine oxides), phosphorus (phosphine oxides), or sulfur (sulfoxides) in the organic molecule.

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or laughing gas

Inorganic compound, one of the oxides of nitrogen. A colourless gas with a pleasantly sweetish odour and taste, it has an analgesic effect when inhaled; it is used as an anesthetic (often called just “gas”) in dentistry and surgery. This effect is preceded by mild hysteria, sometimes with laughter, hence the name laughing gas. It is also used as a propellant in food aerosols and as a leak detector.

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Colourless, toxic gas (NO), formed from nitrogen and oxygen by the action of electric sparks or high temperatures or, more conveniently, by the action of dilute nitric acid on copper or mercury. First prepared circa 1620 by Jan B. Helmont, it was first studied in 1772 by Joseph Priestley, who called it “nitrous air.” An industrial procedure for the manufacture of hydroxylamine is based on the reaction of nitric oxide with hydrogen in the presence of a catalyst. The formation of nitric oxide from nitric acid and mercury is applied in a volumetric method of analysis for nitric acid or its salts. The gas is synthesized via enzyme-catalyzed reactions in humans and other animals, where it serves as a signaling molecule. Among its numerous biological roles, it causes dilation of blood vessels and as such is an important regulator of blood pressure. Nitric oxide is one of the components of air pollution generated by internal-combustion engines.

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or deuterium oxide

Water composed of two atoms of deuterium (D; a heavy isotope of hydrogen) and one atom of oxygen (O), chemical formula D₂O. Water from most natural sources contains about 0.015percnt deuterium oxide; this can be enriched or purified by distillation, electrolysis, or chemical processing. Heavy water is used as a moderator in nuclear power plants, slowing down the fast neutrons so that they can react with the fuel in the reactor. Heavy water is also used in research as an isotopic tracer for chemical reactions and biochemical pathways. Water with tritium (T₂O) rather than deuterium may also be called heavy water.

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An oxide is a chemical compound containing at least one oxygen atom as well as at least one other element. Most of the Earth's crust consists of oxides. Oxides result when elements are oxidized by oxygen in air. Combustion of hydrocarbons affords the two principal oxides of carbon, carbon monoxide and carbon dioxide. Even materials that are considered to be pure elements often contain a coating of oxides. For example, aluminium foil has a thin skin of Al₂O₃ that protects the foil from further corrosion.

Virtually all elements burn in an atmosphere of oxygen. In the presence of water and oxygen (or simply air), some elements - lithium, sodium, potassium, rubidium, caesium, strontium and barium - react rapidly, even dangerously 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. Caesium 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 consist 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 aluminium oxide layer can be built to greater thickness by the process of electrolytic anodising. Although solid magnesium and aluminium react slowly with oxygen at STP, they, like most metals, will burn in air, generating very high temperatures. As a consequence, finely divided powders of most metals can be dangerously explosive in air.

In dry oxygen, iron readily forms iron(II) oxide, but the formation of the hydrated ferric oxides, Fe₂O_3−2x(OH)_x, that mainly comprise rust, typically requires oxygen and water. The production of free oxygen 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.

Due to its electronegativity, oxygen forms chemical bonds with almost all elements to give the corresponding oxides. So-called noble metals (common examples: gold, platinum) resist direct chemical combination with oxygen, and substances like gold(III) oxide must be generated by indirect routes.

Insolubility in water

The oxide ion, O²⁻, is the conjugate base of the hydroxide ion, OH⁻, and is encountered in ionic solid such as calcium oxide. O²⁻ is unstable in aqueous solution − its affinity for H⁺ is so great (pK_b ~ −22) that it abstracts a proton from a solvent H₂O molecule:

O²⁻ + H₂O → 2 OH⁻

Although many anions are stable in aqueous solution, ionic oxides are not. For example, sodium chloride dissolves readily in water to give a solution containing the constituent ions, Na⁺ and Cl⁻. Oxides do not behave like this. If an ionic oxide dissolves, the O²⁻ ions become protonated. Although calcium oxide, CaO, is said to "dissolve" in water, the products include hydroxide:

CaO + H₂O → Ca²⁺ + 2 OH⁻

In fact, no monoatomic dianion is known to dissolve in water - all are so basic that they undergo hydrolysis. Concentrations of oxide ion in water are too low to be detectable with current technology.

Authentic soluble oxides do exist, but they release oxyanions, not O²⁻. Well known soluble salts of oxyanions include sodium sulfate (Na₂SO₄), potassium permanganate (KMnO₄), and sodium nitrate (NaNO₃).

Nomenclature

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

Oxides are usually named after the number of oxygen atoms in the oxide. Oxides containing only one oxygen are called oxides or monoxides, those containing two oxygen atoms are dioxides, three oxygen atoms makes it a trioxide, four oxygen atoms are tetroxides, and so on following the Greek numerical prefixes. In the older literature and continuing in industry, oxides are named by contracting the element name with "a." Hence alumina, magnesia, chromia, are, respectively, Al₂O₃, MgO, Cr₂O₃.

Two other types of oxide are peroxide, O₂²⁻, and superoxide, O₂⁻. In such species, oxygen is assigned higher oxidation states than oxide.

Types of oxides

Oxides of more electropositive elements tend to be basic. They are called basic anhydrides; adding 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 acid; perchloric acid is a more hydrated form.

Some oxides can act as both acid and base at different times. They are amphoteric. An example is aluminium oxide. Some oxides do not show behavior as either acid or base.

The oxides of the chemical elements in their highest oxidation state are predictable and the chemical formula can be 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 that does not exist as expected as F₂O₇ but as OF₂. . Since F is more electronegative than O, OF₂ does not represent an oxide of fluorine, but instead represents a fluoride of oxygen. Phosphorus pentoxide, the third exception is not properly represented by the chemical formula P₂O₅ but by P₄O₁₀

List of all known oxides sorted by oxidation state

• Element in −1 oxidation state
• Oxygen difluoride (OF₂)

• Element in multiple oxidation states
• Antimony tetroxide (Sb₂O₄)
• Cobalt(II,III) oxide (Co₃O₄)
• Iron(II,III) oxide (Fe₃O₄)
• Lead tetroxide (Pb₃O₄)
• Manganese(II,III) oxide (Mn₃O₄)
• Silver(I,III) oxide (AgO)

• Element in +1 oxidation state
• Copper(I) oxide (Cu₂O)
• Dicarbon monoxide (C₂O)
• Dichlorine monoxide (Cl₂O)
• Lithium oxide (Li₂O)
• Potassium oxide (K₂O)
• Rubidium oxide (Rb₂O)
• Silver(I) oxide (Ag₂O)
• Thallium oxide (Tl₂O)
• Sodium oxide (Na₂O)
• Water (hydrogen oxide) (H₂O)

• Element in +2 oxidation state
• Aluminium monoxide (AlO)
• Barium oxide (BaO)
• Beryllium oxide (BeO)
• Cadmium oxide (CdO)
• Calcium oxide (CaO)
• Carbon monoxide (CO)
• Cobalt(II) oxide (CoO)
• Copper(II) oxide (CuO)
• Iron(II) oxide (FeO)
• Lead(II) oxide (PbO)
• Magnesium oxide (MgO)
• Mercury(II) oxide (O)
• Nickel(II) oxide (NiO)
• Nitrogen oxide (NO)
• Palladium(II) oxide (PdO)
• Strontium oxide (SrO)
• Sulphur monoxide (SO)
• Tin(II) oxide (SnO)
• Titanium(II) oxide (TiO)
• Vanadium(II) oxide (VO)
• Zinc oxide (ZnO)

• Element in +3 oxidation state
• Aluminium oxide (Al₂O₃)
• Antimony trioxide (Sb₂O₃)
• Arsenic trioxide (As₂O₃)
• Bismuth trioxide (Bi₂O₃)
• Boron oxide (B₂O₃)
• Chromium(III) oxide (Cr₂O₃)
• Dinitrogen trioxide (N₂O₃)
• Erbium(III) oxide (Er₂O₃)
• Gadolinium(III) oxide (Gd₂O₃)
• Gallium(III) oxide (Ga₂O₃)
• Holmium(III) oxide (Ho₂O₃)
• Indium(III) oxide (In₂O₃)
• Iron(III) oxide (Fe₂O₃)
• Lanthanum(III) oxide (La₂O₃)
• Lutetium(III) oxide (Lu₂O₃)
• Nickel(III) oxide (Ni₂O₃)
• Phosphorus trioxide (P₄O₆)
• Promethium(III) oxide (Pm₂O₃)
• Rhodium(III) oxide (Rh₂O₃)
• Samarium(III) oxide (Sm₂O₃)
• Scandium(III) oxide (Sc₂O₃)
• Terbium(III) oxide (Tb₂O₃)
• Thallium(III) oxide (Tl₂O₃)
• Thulium(III) oxide (Tm₂O₃)
• Titanium(III) oxide (Ti₂O₃)
• Tungsten(III) oxide (W₂O₃)
• Vanadium(III) oxide (V₂O₃)
• Ytterbium(III) oxide (Yb₂O₃)
• Yttrium(III) oxide (Y₂O₃)

• Element in +4 oxidation state
• Carbon dioxide (CO₂)
• Carbon trioxide (CO₃)
• Cerium(IV) oxide (CeO₂)
• Chlorine dioxide (ClO₂)
• Chromium(IV) oxide (CrO₂)
• Dinitrogen tetroxide (N₂O₄)
• Germanium dioxide (GeO₂)
• Hafnium(IV) oxide (HfO₂)
• Lead(IV) oxide (PbO₂)
• Manganese(IV) oxide (MnO₂)
• Nitrogen dioxide (NO₂)
• Plutonium dioxide (PuO₂)
• Ruthenium(IV) oxide (RuO₂)
• Selenium dioxide (SeO₂)
• Silicon dioxide (SiO₂)
• Sulfur dioxide (SO₂)
• Tellurium dioxide (TeO₂)
• Thorium dioxide (O₂)
• Tin dioxide (SnO₂)
• Titanium dioxide (TiO₂)
• Tungsten(IV) oxide (WO₂)
• Uranium dioxide (UO₂)
• Vanadium(IV) oxide (VO₂)
• Zirconium dioxide (ZrO₂)

• Element in +5 oxidation state
• Antimony pentoxide (Sb₂O₅)
• Arsenic pentoxide (As₂O₅)
• Dinitrogen pentoxide (N₂O₅)
• Niobium pentoxide
• Phosphorus pentoxide (P₂O₅)
• Tantalum pentoxide (Ta₂O₅)
• Vanadium(V) oxide (V₂O₅)

• Element in +6 oxidation state
• Chromium trioxide (CrO₃)
• Molybdenum(VI) oxide (MoO₃)
• Rhenium trioxide (ReO₃)
• Selenium trioxide (SeO₃)
• Sulphur trioxide (SO₃)
• Tellurium trioxide (TeO₃)
• Tungsten trioxide (WO₃)
• Uranium trioxide (UO₃)
• Xenon trioxide (XeO₃)

• Element in +7 oxidation state
• Dichlorine heptoxide (Cl₂O₇)
• Manganese(VII) oxide (Mn₂O₇)
• Rhenium(VII) oxide (Re₂O₇)
• Technetium(VII) oxide

• Element in +8 oxidation state
• Osmium tetroxide (OsO₄)
• Ruthenium tetroxide (RuO₄)
• Xenon tetroxide (XeO₄)

See also

• Other oxygen ions ozonide, O₃⁻, superoxide, O₂⁻, peroxide, O₂²⁻ and dioxygenyl, O₂⁺.
• Suboxide
• See Oxides for a list of oxides.

References