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, Fe
2O
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 Fe
2O
3 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, O
2−, is the
conjugate base of the
hydroxide ion, OH
−, and is encountered in
ionic solid such as
calcium oxide. O
2− is unstable in
aqueous solution − its affinity for H
+ is so great (p
Kb ~ −22) that it abstracts a
proton from a solvent H
2O 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, Al
2O
3, MgO, Cr
2O
3.
Two other types of oxide are
peroxide, O
22−, and
superoxide, O
2−. 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 O
4,
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 F
2O
7 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 P
2O
5 but by P
4O
10.
List of all known oxides sorted by oxidation state
- Element in -1 oxidation state
- Element in multiple oxidation states
- Element in +1 oxidation state
- Element in +2 oxidation state
- Element in +3 oxidation state
- Element in +4 oxidation state
- Element in +5 oxidation state
- Element in +6 oxidation state
- Element in +7 oxidation state
- Element in +8 oxidation state
