An oxide is a chemical compound that contains at least one oxygen atom and one other element
in its chemical formula.
Metal oxides 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.
Hydrocarbon combustion affords the two principal carbon oxides: carbon monoxide and carbon
dioxide.
Even materials considered pure elements often develop an oxide coating.
For example, aluminium foil develops a thin skin of Al2O3 that protects the foil from
further corrosion.
Different oxides of the same element are distinguished by Roman numerals denoting their oxidation
number, e.g. iron(II) oxide versus iron(III) oxide.
Formation
Due to its electronegativity, oxygen forms stable chemical bonds with almost all elements
to give the corresponding oxides.
Noble metals 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, some elements— sodium—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.
Though solid magnesium and aluminium 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.
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.
Structure Oxides of most metals adopt polymeric structures
with M-O-M crosslinks.
Because these crosslinks are strong, the solids tend to be insoluble in solvents, though they
are attacked by acids and bases.
The formulas are often deceptively simple.
Many are nonstoichiometric compounds.
In these oxides, the coordination number of the oxide ligand is two for most electronegative
elements and 3–6 for most metals.
Molecular oxides Although most metal oxides are polymeric,
some oxides are molecules.
The most famous molecular oxides are carbon dioxide and carbon monoxide.
Phosphorus pentoxide is a more complex molecular oxide with a deceptive name, the formula being
P4O10.
Some polymeric oxides when heated depolymerize to give molecules, examples being selenium
dioxide and sulfur trioxide.
Tetroxides are rare, and there are only five known examples: ruthenium tetroxide, osmium
tetroxide, hassium tetroxide, iridium tetroxide, and xenon tetroxide.
Many oxyanions are known, such as polyphosphates and polyoxometalates.
Oxycations are rarer, an example being nitrosonium.
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.
Reactivity Oxides can be attacked by acids and bases.
Those attacked only by acids are basic oxides; those attacked only by bases are acidic oxides.
Oxides that react with both acids and bases 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.
This reactivity is the basis of many practical processes such, as the extraction of some
metals from their ores in the process called hydrometallurgy.
Reduction
Metals are "won" from their oxides by chemical reduction.
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 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 centre transfers an oxide ligand to the organic compound followed
by regeneration of the metal oxide, often by oxygen in air.
Hydrolysis 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 acid; perchloric acid is a more 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 solid
such as calcium oxide.
O2− is unstable in aqueous solution − its affinity for H+ is so great 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.
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 contracting
the element name with "a."
Hence alumina, magnesia, chromia, are, respectively, Al2O3, MgO, Cr2O3.
Special types of oxides are peroxide, O22−, and superoxide, O2−.
In such species, oxygen is assigned higher 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.
Since fluorine is more electronegative than oxygen, oxygen difluoride 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.
See also Other oxygen ions ozonide, O3−, superoxide,
O2−, peroxide, O22− and dioxygenyl, O2+.
Suboxide Oxohalide
Oxyanion See Category:Oxides for a list of oxides.
Salt References