Conductor of Electricity

Graphite, unlike diamond and amorphous carbon, can conduct electricity due to its unique covalent bonding. Each carbon atom in graphite is bonded to three other carbon atoms in a trigonal planar structure. The fourth electron is free to move between layers, creating delocalized electrons that form a current. This property sets graphite apart from other covalent network substances. Now, let's look at the physical properties of graphite. Yes, graphite, diamond, amorphous carbon, etc. are all made up of only carbon atoms, but that doesn't mean that the bonding between the atoms are the same in each allotrope. Graphite is the only allotrope of carbon that is able to conduct electricity, and this is because of the unique covalent bonding between each carbon atom. Take that, diamond. Anyway, in graphite, each carbon atom is covalently bonded to three other carbon atoms in a trigonal planar geometry. All carbon atoms in the graphite lattice have four valence electrons, three of which are used to form covalent bonds with three neighboring carbon atoms, and the fourth electron is free to move about between layers. The delocalized electrons are considered mobile charged species and so form a current. This allows graphite to conduct electricity. Such is not the case for most covalent network substances.