Diamond and graphite are different forms of the element carbon. They both have giant structures of carbon atoms, joined together by covalent bonds. However, their structures are different so some of their properties are different.
Diamond is a giant covalent structure in which:
The rigid network of carbon atoms, held together by strong covalent bonds, makes diamond very hard. This makes it useful for cutting tools, such as diamond-tipped glass cutters and oil rig drills.
Diamond has a very high melting point because a lot of energy is required to break the strong covalent bonds between the atoms. It does not conduct electricity because it has no free electrons.
Graphite has a giant covalent structure in which:
Graphite has delocalised electrons, just like metals. These electrons are free to move between the layers in graphite, so graphite can conduct electricity. This makes graphite useful for electrodes in batteries and for electrolysis.
The forces between the layers in graphite are weak. This means that the layers can slide over each other. This makes graphite slippery, so it is useful as a lubricant.
Explain why diamond does not conduct electricity and why graphite does conduct electricity.
Diamond does not conduct electricity because it has no charged particles that are free to move. Graphite does conduct electricity because it has delocalised electrons which move between the layers.