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Why does graphite conduct electricity while diamond acts as an insulator?

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graphene: atomic force microscopy — graphene Atomic force microscopy image of water structures that were trapped beneath a sheet of graphene one atom thick.

The difference in electrical conductivity between graphite and diamond arises from their distinct crystalline structures and bonding arrangements.1

In diamond, each carbon atom is covalently bonded to four other carbon atoms in a rigid, three-dimensional tetrahedral lattice. Because all valence electrons are localized within these strong covalent bonds, there are no free electrons available to carry an electrical current, which makes diamond an insulator.12

In contrast, graphite consists of parallel layers of carbon atoms arranged in fused hexagonal rings. Within these layers, the bonding involves complete π-bonding, which allows for electron mobility. While the layers themselves are held together by weak van der Waals forces, the delocalized electrons within the sheets enable graphite to conduct electricity. This conductivity is anisotropic, meaning it is significantly higher within the plane of the sheets than perpendicular to them. The individual layers of graphite are known as graphene.12