Question

Why is graphite a good conductor of electricity but not a diamond?

Answer 1

Graphite

• Graphite and diamond are allotropes of carbon.
• Each carbon in graphite is $s{p}^2$ hybridized.
• A hexagonal array in graphite is formed when each carbon atom is connected to three other carbon atoms in the same plane.
• Structure of graphite:

Electricity conductivity of graphite

• In graphite, only three carbon atoms are directly bonded to one another through covalent bonds.
• As a result, bonding requires just three of a carbon atom's four valence electrons, while the fourth is relatively free and can move from one carbon atom to another.
• With these free electrons, graphite is a good conductor of electricity.

Diamond

• Each carbon atom in a diamond is connected to four other carbon atoms by strong covalent bonds, forming a huge covalent structure.
• In diamond a regular tetrahedral network structure is formed by the carbon atoms.
• Structure of diamond:

Electricity conductivity of Diamond

• Diamonds have no free electrons that can move and conduct electricity because of their tetrahedral arrangement of covalently bound carbon atoms.
• Since the existence of free electrons is required for electrical conduction.
• As a result, diamond is a poor electrical conductor.