Chemistry – Atoms, elements and compounds - Giant covalent structures | e-Consult

Diamond: In diamond, each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement. All the valence electrons of the carbon atoms are involved in these strong covalent bonds, forming a fully filled octet. Therefore, there are no free electrons available to carry an electrical charge. Hence, diamond is a poor conductor of electricity.

Graphite: In graphite, the carbon atoms are arranged in layers, and within each layer, the carbon atoms are covalently bonded. However, the bonds between the layers are weak Van der Waals forces. Within each layer, one valence electron of each carbon atom is delocalised and forms a 'sea' of electrons that are free to move throughout the layer. These delocalised electrons are responsible for the excellent electrical conductivity of graphite. When a voltage is applied, these electrons can move freely through the layers, carrying an electrical current.