Giant Covalent Structures
- Diamond and graphite both consist of carbon atoms and 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
Diamond
- In diamond, each carbon atom bonds with four other carbons, forming a tetrahedron
- All the covalent bonds are identical and very strong
Diagram to show the structure of diamond
In diamond, each carbon is bonded to four other carbon atoms
Linking the Bonding & Properties
- Diamond does not conduct electricity
- All the outer shell electrons in carbon are held in the four covalent bonds around each carbon atom
- As a result, there are no freely moving particles to carry a charge
- Diamond has a very high melting point
- Diamond has a giant covalent structure
- There are strong covalent bonds between the carbon atoms
- These need lots of energy to break
- It is extremely hard and dense
- It has strong covalent bonds and each carbon atom is bonded to four other carbon atoms
- Diamond's hardness makes it very useful in cutting tools like drills
Examiner Tip
Diamond is the hardest naturally occurring mineral, but it is by no means the strongest. Students often confuse hard with strong, thinking it is the opposites of weak. Diamonds are hard, but brittle – that is, they can be smashed fairly easily with a hammer. The opposite of saying a material is hard is to describe it as soft.
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
Diagram to show the bonding and structure in graphite
Each carbon atom is bonded to three other carbon atoms
Linking the Bonding & Properties
- Graphite conducts electricity
- Each carbon atom is bonded to three others leaving one free electron per carbon atom
- These free (delocalised) electrons exist in between the layers
- They are free to move through the structure and carry charge
- Graphite has a high melting point
- Graphite has a giant covalent structure
- There are strong covalent bonds between the carbon atoms
- These need lots of energy to break
- Graphite is slippery
- Graphite is arranged in layers
- Although the atoms within the layers are joined by strong covalent bonds, the layers have only weak intermolecular forces between them
- As a result the layers can slide over each other
- This property allows graphite to be used in pencils and as an industrial lubricant
Examiner Tip
Don’t confuse pencil lead with the metal lead – they have nothing in common. Pencil lead is actually graphite, and historical research suggests that in the past, lead miners sometimes confused the mineral galena (lead sulfide) with graphite; since the two looked similar they termed both minerals ‘lead’. The word graphite derives from the Latin word ‘grapho’ meaning ‘I write’, so it is a well named mineral!
