The form of carbon known as graphite

Graphite and diamond are allotropes of carbon with vastly different physical properties due to their distinct crystal structures. Diamond is the hardest known natural substance because of its rigid, three-dimensional tetrahedral network where each carbon atom is sp3 hybridized and bonded to four others [2]. In contrast, graphite has a hexagonal layered structure where each carbon atom is sp2 hybridized and bonded to only three others [1]. This arrangement leaves one valence electron per carbon atom delocalized, forming a 'sea of electrons' that allows graphite to be an excellent conductor of electricity [1]. Diamond, lacking free electrons, acts as an electrical insulator. Furthermore, graphite's bond distances are not equal in all directions; while the covalent bonds within layers are strong and short, the distance between layers is much larger due to weak van der Waals forces.

Sources

1. [1] Science , class X (NCERT 2025 ed.) > Chapter 4: Carbon and its Compounds > Allotropes of carbon > p. 61
2. [2] https://www.scientificamerican.com/article/how-can-graphite-and-diam/