Graphite is a much better conductor of heat and electricity than diamond. This is due to the fact that each carbon atom in graphite:

Graphite's superior conductivity compared to diamond is fundamentally due to its atomic bonding and hybridization. In graphite, each carbon atom undergoes sp2 hybridization, forming three strong sigma bonds with three neighboring carbon atoms in a trigonal planar arrangement. This leaves one unhybridized p-orbital containing a fourth valence electron. These unhybridized electrons overlap to form a delocalized pi-electron system across the hexagonal layers. These delocalized electrons are free to move, facilitating the conduction of electricity and heat. In contrast, diamond undergoes sp3 hybridization where all four valence electrons are localized in strong tetrahedral sigma bonds, leaving no free electrons for conduction [1]. While graphite's layers are held by weak van der Waals forces, its conductivity is specifically a result of the sp2 hybridization and the resulting delocalized electrons.

Sources

1. [1] https://www.sciencedirect.com/science/article/abs/pii/S0301010422000581