The term allotrope was first suggested by Swedish chemist J. J. Berzelius (1779-1848).
Allotropes are two or more forms of the same element in the same physical state (solid, liquid, or gas) that differ from each other in their physical and sometimes chemical properties. The most notable examples of allotropes are found in groups 14, 15, and 16 of the periodic table.
Gaseous oxygen, for example, exists in three allotropic forms: monatomic oxygen (O), a diatomic molecule (O2), and in a triatomic molecule known as ozone (O3).
An important example of different physical properties among allotropes is the case of carbon. Solid carbon exists in two allotropic forms: diamond and graphite. Diamond is the hardest naturally occurring substance and having the highest melting point of more than 3,502°C. Graphite is a very soft material, the substance which is used as a lead in lead pencils. The Braggs used x-ray diffraction to show that diamond and graphite differ from each other in their atomic structure.
The allotropes of phosphorus have different chemical properties that may occur among such forms. White phosphorus is a waxy white solid that bursts into flame spontaneously when exposed to air. It is highly toxic. On the other hand, a second allotrope of phosphorus known as red phosphorus is very stable. It does not react with air, and is nontoxic.
Allotropes differ from each other structurally depending on the number of atoms in a molecule of the element. There are allotropes of sulphur (sulfur), for example, that contain 2, 6, 7, 8, 10, 12, 18, and 20 atoms per molecule (formulas S2 to S20).