Diamond and graphite are both allotropes of carbon. Although both are composed of carbon atoms, their atomic organization is completely different. When a carbon atom forms covalent bonds with surrounding carbon atoms, there are two different bond formations. In diamond, the carbon atoms are arranged in a regular tetrahedral pattern, while in graphite, the carbon atoms are arranged in hexagonal layers.
In the late 18th century, Lavoisier identified diamond as a crystalline form of carbon, but the crystal structure of diamond was first determined by Bragg in 1913. He irradiated diamond with X-rays and obtained a photograph of the X-ray diffraction pattern of the crystal lattice, demonstrating the crystalline properties of diamond.
In addition to tetrahedrons, natural diamonds also include cubes (CD), hexahedrons (HD), octahedrons (OD), and dodecahedrons. The carbon atoms that make up a diamond crystal are covalently bonded, with bond lengths of equal length, 0.15 nm. Diamond is anisotropic, meaning that many of its properties are directionally dependent.
In a tetrahedron, the carbon atoms at the center of symmetry have weaker bonding forces, making diamond susceptible to cleavage there. In the typical crystal faces of diamond and the flaky structure of graphite, carbon bonds within the graphite sheets have a length of 0.142 nm, making them very strong. However, the bonds between the sheets are only 0.335 nm long, making these long bonds very weak. This gives graphite excellent lubricity. Diamond crystals are composed of carbon atoms, which are held together by covalent bonds.