Diamond & Graphite (CIE IGCSE Chemistry)

• Diamond and graphite are allotropes of carbon which have giant covalent structures
• Both substances contain only carbon atoms but due to the differences in bonding arrangements they are physically completely different
• Giant covalent structures contain billions of non-metal atoms, each joined to adjacent atoms by covalent bonds forming a giant lattice structure

Diamond

• In diamond, each carbon atom bonds with four other carbons, forming a tetrahedron
• All the covalent bonds are identical, very strong and there are no intermolecular forces

Diagram showing the structure and bonding arrangement in diamond

Graphite

• Each carbon atom in graphite is bonded to three others forming layers of hexagons, leaving one free electron per carbon atom which becomes delocalised
• The covalent bonds within the layers are very strong, but the layers are attracted to each other by weak intermolecular forces

The structure and bonding in graphite

Properties of Diamond

• Diamond has the following physical properties:
  • It does not conduct electricity
  • It has a very high melting point
  • It is extremely hard and dense

• All the outer shell electrons in carbon are held in the four covalent bonds around each carbon atom, so there are no freely moving charged particles to carry the current thus it cannot conduct electricity
• The four covalent bonds are very strong and extend in a giant lattice, so a very large amount of heat energy is needed to break the lattice thus it has a very high melting point
• Diamond ́s hardness makes it very useful for purposes where extremely tough material is required
• Diamond is used in jewellery due to its sparkly appearance and as cutting tools as it is such a hard material
• The cutting edges of discs used to cut bricks and concrete are tipped with diamonds
• Heavy-duty drill bits and tooling equipment are also diamond-tipped

Properties of Graphite

• Each carbon atom is bonded to three others forming layers of hexagonal-shaped forms, leaving one free electron per carbon atom
• These free (delocalised) electrons exist in between the layers and are free to move through the structure and carry charge, hence graphite can conduct electricity
• The covalent bonds within the layers are very strong but the layers are connected to each other by weak forces only, hence the layers can slide over each other making graphite slippery and smooth
• Graphite thus:
  • Conducts electricity
  • Has a very high melting point
  • Is soft and slippery, less dense than diamond 

• Graphite is used in pencils and as an industrial lubricant, in engines and in locks
• It is also used to make non-reactive electrodes for electrolysis