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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

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, or an oxygen rich environment. 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 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 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 grained 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, Fe2O3−2x(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.

Contents

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 (pKb ~ −22) that it abstracts a proton from a solvent H2O molecule:

O2− + H2O → 2 OH

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.[1] 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

See also

References

  1. ^ Fully Exploiting the Potential of the Periodic Table through Pattern Recognition Schultz, Emeric. J. Chem. Educ. 2005 82 1649.

1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

OXIDE, in chemistry, a binary compound of oxygen and other elements. In general, oxides are the most important compounds with which the chemist has to deal, a study of their composition and properties permitting a valuable comparative investigation of the elements. It is possible to bring about the direct combination of oxygen with most of the elements (the presence of traces of water vapour is generally necessary according to the researches of H. B. Baker), and when this is not so, indirect methods are available, except with bromine and fluorine (and also with the so-called inert gases - argon, helium, &c.), which so far have yielded no oxides. Most of the elements combine with oxygen in several proportions, for example nitrogen has five oxides: N 2 0, NO, N 2 0 3, NO 2, N205; for classificatory purposes, however, it is advantageous to assign a typical oxide to each element, which, in general, is the highest having a basic or acid character. Thus in Group I. of the periodic system, the typical oxide is M 2 O, of Group II. MO, of Group III. M 2 O 3, of Group IV. MO 2, of Group V. M 2 O 5, of Group VI. M03.

Five species of oxides may be distinguished: (I) basic oxides, (2) acidic oxides, (3) neutral oxides, (4) peroxides, (5) mixed anhydrides and salts. Basic oxides combine with acids or acidic oxides to form salts; similarly acidic oxides combine with basic oxides to form salts also. The former are more usually yielded by the metals (some metals, however, form oxides belonging to the other groups), whilst the latter are usually associated with the non-metals. An oxide may be both acidic and basic, i.e. combine with bases as well as acids; this is the case with elements occurring at the transi tion between basigenic and oxygenic elements in the periodic classification, e.g. aluminium and zinc. Neutral oxides combine neither with acids nor bases to give salts nor with water to give a base or acid. A typical member is nitric oxide; carbon monoxide and nitrous oxide may also be put in this class, but it must be remembered that these oxides may be regarded, in some measure at least, as the anhydrides of formic and hyponitrous acid, although, at the same time, it is impossible to obtain these acids by simple hydration of these oxides. Peroxides may in most cases be defined as oxides containing more oxygen than the typical oxide. The failure of this definition is seen in the case of lead dioxide, which is certainly a peroxide in properties, but it is also the typical oxide of Group IV. to which lead belongs. All peroxides have oxidizing properties. Peroxides may be basic or acidic. Some basic oxides yield hydrogen peroxide with acids, others yield oxygen (these also liberate chlorine from hydrochloric acid), and may combine with lower acidic oxides to form salts of the normal basic oxide with the higher acidic oxide. Examples are Ba02-1-H2S04=BaS04+H202; 2Mn02-I-2H2S04=2MnS04-1-2H20+02; Mn02-{-4HC1=MnC12+ 2H 2 0 +C1 2; Pb02+S02=PbS04 (i.e. Pb0+503) Two species of basic peroxides may be distinguished: (I) the superoxides or peroxidates, containing the oxygen atoms in a chain, e.g. Na 0 0 Na, Bab, which yield hydrogen peroxide with acids; and (2) the polyoxides, having the oxygen atoms doubly linked to the metallic atom, e.g. 0: M n: 0,0: Pb: O, and giving oxygen with sulphuric acid, and chlorine with hydrochloric. L. Marino (Zeit. anorg. Chem., 1907, 5 6, p. 2 33) pointed out that manganese and lead dioxide behaved differently with sulphur dioxide, the former giving dithionate and the latter sulphate, and suggested the following formulae: O:Mn :0, O Pb: as explaining this difference. A simpler explanation is that the manganese dioxide first gives a normal sulphite which rearranges to dithionate, thus: Mn0 2 +2S0 2 = Mn(S03)2-MnS206, whilst the lead dioxide gives a basic sulphite which rearranges to sulphate, thus: PbO-l-S0 2 = PbOS03 - > PbS04. Acidic peroxides combine with basic oxides to form "per" salts, and by loss of oxygen yield the acidic oxide typical of the element. Mixed anhydrides are oxides, which yield with water two acids, or are salts composed of a basic and acidic oxide of the same metal. Examples of mixed anhydrides are C10 2 and N02, which give chlorous and chloric acid, and nitrous and nitric acid: 2C102+ H20=HC102+HC103, 2N02+H20=HN02+HN03; and of mixed salts Pb203 and Pb304, which may be regarded as lead metaand ortho-plumbate: Pb0 Pb02, 2PbO Pb02.

Oxidation and Reduction

In the narrow sense "oxidation" may be regarded as the combination of a substance with oxygen, and conversely, "reduction" as the abstraction of oxygen; in the wider sense oxidation includes not merely the addition of oxygen, but also of other electro-negative elements or groups, or the removal of hydrogen or an electro-positive element or group. In inorganic chemistry oxidation is associated in many cases with an increase in the active valency. Ignoring processes of oxidation or reduction simply brought about by heat or some other form of energy, we may regard an oxidizing agent as a substance having a strong affinity for electro-positive atoms or groups, and a reducing agent as having a strong affinity for electro-negative atoms or groups; in the actual processes the oxidizing agent suffers reduction and the reducing agent oxidation.

Many substances undergo simultaneous oxidation and reduction when treated in a particular manner; this is known as selfor auto-oxidation. For example, on boiling an aqueous solution of a hypochlorite, a chlorate and a chloride results, part of the original salt being oxidized and part reduced: 3NaOC1= NaC103-2NaC1. Similarly phosphorous and hypophosphorous acids give phosphoric acid and phosphene, whilst nitrous acid gives nitric acid and nitric oxide: 4H3P03=3H3P04+PH3; 2H 3 PO 2 =H 3 PO 4 +PH 3 i 3HN02= HNO 3 +2NO--H20. In organic chemistry, a celebrated example is Cannizzaro's reaction wherein an aromatic aldehyde gives an acid and an alcohol: 2C E H 5 CHO+H 2 O =C6H5C02H+C6H5CH20H.

The important oxidizing agents include: oxygen, ozone, peroxides, the halogens chlorine and bromine, oxyacids such as nitric and those of chlorine, bromine and iodine, and also chromic and permanganic acid. The important reducing agents include hydrogen, hydrides such as those of iodine, sulphur, phosphorus, &c., carbon, many metals, potassium, sodium, aluminium, magnesium, &c., salts of lower oxyacids, lower salts of metals and lower oxides.


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Simple English

An oxide is a chemical compound containing an oxygen atom and other elements. Most of the earth's crust consists of oxides. Oxides result when elements are oxidized by air (when oxygen in the air react with the element).

Famous oxides are:

Aluminium oxide

Lead oxide

Calcium oxide

Iron(III) oxide (Rust)

Carbon dioxide

Carbon monoxide

Magnesium oxide (Magnesia)

Phosphorus pentoxide

Sulfur dioxide

Sulfur trioxide

Silicon dioxide

Nitrogen oxide

Zinc oxide

Copper oxide

Hydrogen oxide









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