Conductor of Heat And Electricity

Graphite, unlike other forms of carbon like diamond, is able to conduct electricity due to its unique covalent bonding between carbon atoms. Each carbon atom in the graphite lattice is bonded to three other carbon atoms, and the fourth electron is free to move between layers, allowing for the conduction of electricity. Graphite is also a great conductor of heat due to the close proximity of its carbon layers, which allows for rapid transfer of vibrations between atoms. 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 carbon atoms. Take that, diamond! Anyway, in graphite, each carbon atom is covalently bonded to three other carbon atoms in a trigonal plane in geometry. All carbon atoms in the graphite lattice have four valence electrons, three of which are used to form covalent bonds with three neighbouring carbon atoms, and the fourth electron is free to move about between layers. The delocalised 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 networked substances. Additionally, heat conduction works by the transfer of vibrations between particles in a substance. The carbon layers in the graphite are close to one another, allowing vibrations of carbon atoms to be rapidly transferred to atoms in layers above. Thus, graphite is also a fantastic conductor of heat.