Graphite is a good conductor of electricity. It is because of its

Graphite's electrical conductivity is a result of its hexagonal multilayer structure. In graphite, each carbon atom is covalently bonded to three other carbon atoms in the same plane (sp² hybridization), forming a series of flat hexagonal rings. This arrangement leaves one valence electron per carbon atom unbonded and "delocalized". These free electrons are mobile and can move throughout the layers, allowing graphite to conduct electricity. In contrast, diamond has a tetrahedral structure (sp³ hybridization) where all four valence electrons are involved in strong covalent bonds, leaving no free electrons and making it an insulator. While the local geometry around each carbon in graphite is triangular planar, the bulk property of conductivity is attributed to the delocalized electron system within its layered hexagonal lattice.