Carbon Allotropes

Due to its extreme hardness, diamond is widely used in industrial applications such as cutting tools, drill bits, and abrasives.

Graphite's layered structure allows the layers to slide past each other easily, which makes it effective as a solid lubricant.

Graphite consists of stacked layers of carbon atoms arranged in a hexagonal lattice. The layers are held together by weak intermolecular forces, allowing them to slide.

When carbon atoms only form three bonds (as in graphite), the fourth outer electron becomes delocalized. These delocalized electrons are free to move and carry charge, allowing the substance to conduct electricity.

In both graphite and fullerenes, each carbon atom is covalently bonded to three other carbon atoms, leaving one electron per atom delocalized.

In the diamond crystal lattice, each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral geometry. This strong, three-dimensional network structure is what gives diamond its extreme hardness.

Diamond is widely recognized as the hardest naturally occurring substance found on Earth. This property is a direct result of its rigid, covalent network structure.

Carbon atoms typically form covalent bonds by sharing electrons with other atoms. This is the primary bonding mechanism in organic chemistry and carbon allotropes like diamond and graphite.

Allotropes are different physical forms of the same element. Buckminsterfullerene, diamond, and graphite are all forms of pure carbon, whereas sand is primarily silicon dioxide.

Allotropes are defined as different structural forms of the same element. Therefore, carbon allotropes are composed entirely of carbon atoms arranged in different patterns.