Graphite and diamond are both giant covalent structures made entirely of carbon atoms, yet they have very different properties. Explain why graphite can conduct electricity but diamond cannot, and why graphite is soft whilst diamond is very hard.

In graphite, each carbon atom uses three of its four valence electrons to form strong covalent bonds with three other carbons in the same plane, leaving one electron delocalised over the layer. This delocalised electron can move freely, allowing graphite to conduct electricity.

In diamond, every valence electron of each carbon atom is involved in a covalent bond with four neighbouring carbons. There are no free electrons, so diamond does not conduct electricity.

The layers in graphite are held together only by weak van der Waals forces, so they can slide over one another, giving graphite a soft, greasy feel.

In diamond, each carbon atom is tetrahedrally bonded to four others in all directions, forming a rigid, three‑dimensional network that resists deformation, making diamond extremely hard.