Why does graphite conduct electericity but diamond does not if both substances have a giant covalent structure?

Graphite consists of carbon atoms joined together. Since carbon has 4 unpaired electrons in it's outer shell, it can form a maximum of 4 covalent bonds by sharing each of it's unpaired electrons with another atom. However, each carbon atom forms only three covalent bonds with other carbon atoms in graphite. This means that there will be one unpaired electron that remains unbonded in each carbon atom. This electron is referred to as delocalised as it is unbonded. When a current is passed through graphite, the delocalised electrons are able to freely move which allows the current to continue doen the graphite, making graphite a conductor of electricity even if it has a giant covalent structure. Bonded electrons however, cannot move and therefore cannot pass on a current. Diamond, which also has a giant covalent structure, consists of carbon atoms but this time, each carbon atom forms four covalent bonds with other atoms. Therefore, all the electrons are bonded and there are no delocalised or unpaired electrons. When a current is passed, there are no free electrons to allow the current to continue through so diamond does not conduct electricity.