(a) Graphite
– Black, soft, greasy
– Conducts electricity
– Most common form
– Used as lubricant, conductor in drycells and in pencil ‘leads’
(b) Diamonds
– Transparent, hard
– Does not conduct electricity
– Uncommon
– Used in jewelry, blades
Graphite
Make sure you know about the layer structure of graphite. Within each layer each carbon atom is bonded to three others by strong covalent bonds. Each layer is therefore a giant molecule. Between these layers there are weak forces of attraction (=van der Waal’s forces) and so the layers will pass over each other easily.
With only 3 covalent bonds formed between carbon atoms within the layers, an unbonded electron is present on each carbon atom. These spare electrons form electron clouds between the layers and it is because of these spare electrons that graphite conducts electricity.
Diamond
You also need to know the structure of diamond. Each of the carbon atoms in the giant structure is covalently bonded to four others. They form a tetrahedral arrangement. This bonding scheme gives rise to a very rigid, three dimensional structure and accounts for the extreme hardness of the substance. All the outer energy level electrons of the carbon atoms are used to form covalent bonds, so there are no electrons available to enable diamond to conduct electricity.
These two examples of allotropes are important. You must be able to
give account of differences.