How do diamond and graphite differ in terms of electrical conductivity?

Diamond is an electrical insulator, while graphite is a good conductor of electricity.

Diamond and graphite are both forms of carbon, known as allotropes, but they have different structures and properties. The difference in their electrical conductivity is due to the way the carbon atoms are arranged in each allotrope.

In diamond, each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral structure. This means that all of the electrons are involved in bonding and there are no free electrons available to carry an electric current, making diamond an excellent insulator.

On the other hand, in graphite, each carbon atom is covalently bonded to three other carbon atoms, forming layers of hexagonal rings. This leaves one electron per carbon atom free to move. These free electrons can carry an electric current, making graphite a good conductor of electricity. The layers in graphite are also able to slide over each other, which is why graphite is used as a lubricant.

Furthermore, the difference in electrical conductivity between diamond and graphite can be explained by band theory. In diamond, the valence band is fully occupied and the conduction band is empty, meaning there is a large energy gap between them. This makes it difficult for electrons to move from the valence band to the conduction band, hence diamond is an insulator. In contrast, in graphite, the valence band and conduction band overlap, meaning electrons can easily move from the valence band to the conduction band, hence graphite is a conductor.

In summary, the difference in electrical conductivity between diamond and graphite is due to the different arrangements of carbon atoms and the availability of free electrons in each allotrope.