If a carbon atom enters into the structure of diamond its two 2 and 2 electrons redistribute into four new equal-energy-level orbitals called 2 hybrid orbitals. It require a loss of energy but this effect is compensated by a very profitable covalent bonding. The angular distribution of the wave functions for these four 2 orbitals can be illustrated by drawing four lobes whose axes are at 10928 to each other, the axes of these lobes thus extend toward the corners of an imaginary tetrahedron centered around the carbon atom [Fig.2.1].
Quantum-mechanical calculations indicate that greater overlap between orbitals results in a stronger covalent bond. The diamond structure represents a three-dimensional network of strong covalent bonds [Fig.2.2], which explains why diamond is so hard.
The diamond structure is cubic with a cube edge length of Å and can be viewed as two interpenetrating FCC structures displaced by (1/4,1/4,1/4). The diamond crystal is highly symmetric with a cubic space group .
Since all the valence electrons contribute to the covalent bond, they are not free to migrate through the crystal and thus, diamond is a poor conductor with a bandgap of 5.48 eV.