Graphite is an excellent conductor of electricity due to the presence of delocalised electrons that can move freely along its layered structure.

As one of the allotropes of carbon, graphite represents a distinct form in which carbon exists. Its unique structure contributes significantly to its electrical conductivity. In graphite, each carbon atom forms covalent bonds with three neighboring carbon atoms, resulting in layers of hexagonal rings that create a two-dimensional sheet-like structure.

The fourth electron of each carbon atom is not involved in these covalent bonds, allowing it to be free to move. These electrons are referred to as delocalised or free electrons. In graphite, these delocalised electrons can move along the layers of carbon atoms, which is the fundamental reason electricity can flow through the material.

The layers within graphite are held together by weak Van der Waals forces, which enable the layers to slide past one another easily. This characteristic not only gives graphite its slippery texture—making it ideal for use in pencils—but also facilitates the movement of delocalised electrons between the layers.

It is crucial to recognize that not all carbon allotropes exhibit electrical conductivity. For instance, diamond, another allotrope of carbon, does not conduct electricity. This is primarily because all four electrons of each carbon atom in diamond participate in covalent bonding, leaving no free electrons available to carry an electric charge.

In summary, the ability of graphite to conduct electricity arises from its unique structure, which features layers of carbon atoms and free, delocalised electrons capable of moving along and between these layers.