Why can graphite conduct electricity but diamonds cannot?

Though structurally very different, both are enormous carbon structures.

The structure of graphite is essentially layers of carbon atoms stacked on top of one another. In each layer, a carbon atom is bonded to three other carbon atoms in a trigonal planar shape, with London forces holding each sheet to the other.

Since there are only three bonds, every carbon atom will have a nonbonding electron in the p-orbital that will form π bonds with the other carbon atoms' nonbonding electrons; however, this particular π bond is unique because it is delocalized, or shared among the atoms, and these delocalized electrons act as charge carriers that allow electricity to flow through.

In contrast, every carbon atom in a diamond is bonded to four other carbon atoms, ensuring that all of the atoms and electrons are fixed in place and that electricity cannot easily pass through.

Diamonds cannot conduct electricity because of their strong covalent bonding, whereas graphite can because it has delocalized electrons that are free to move throughout its layers.