Diamond and graphite (AQA GCSE Chemistry Combined Science): Revision Notes
Diamond and graphite
What are giant covalent structures?
Some substances have giant covalent structures. This means many atoms are connected together in a big network by covalent bonds. The atoms form a pattern called a lattice.
Giant covalent structures are different from simple molecules because they contain thousands or millions of atoms all bonded together in one continuous network. This gives them very different properties compared to substances made of individual molecules.
Both diamond and graphite are made of carbon atoms, but they have very different structures and properties.
Diamond structure
In diamond, carbon atoms form a 3D lattice structure. Each carbon atom makes four covalent bonds with other carbon atoms around it. This creates a very strong network throughout the whole structure.
Key features of diamond:
- Every carbon atom bonds to four others
- Strong covalent bonds between all atoms
- Forms a rigid 3D structure
Diamond properties
Because of its strong bonding structure, diamond has these properties:
- Very hard - difficult to break or scratch
- Very high melting point - needs lots of energy to break the strong bonds
- Does not conduct electricity - all electrons are used in bonding, so none are free to move
Diamond's hardness comes from the fact that breaking the material requires breaking many strong covalent bonds throughout the 3D structure. This is why diamond is used in cutting tools and drill bits.
Graphite structure
In graphite, carbon atoms form layers. Each carbon atom makes three covalent bonds with other carbon atoms in the same layer. The fourth electron from each carbon atom becomes delocalised (free to move).
Key features of graphite:
- Each carbon atom bonds to three others
- Strong covalent bonds within each layer
- Weak forces between the layers
- Delocalised electrons can move freely
Graphite properties
Because of its layered structure, graphite has these properties:
- Soft and slippery - layers can slide over each other easily
- Very high melting point - strong bonds within layers need lots of energy to break
- Conducts electricity - delocalised electrons can carry electrical current
The layers in graphite are held together by weak intermolecular forces, which is why they can slide over each other easily. This is what makes graphite useful as a lubricant and in pencil leads.
Why can graphite conduct electricity but diamond cannot?
This is one of the most important differences between these two forms of carbon.
Electrical conduction explained:
In graphite, each carbon atom only uses three of its four outer electrons for bonding. The fourth electron becomes delocalised and can move freely between the layers. These moving electrons can carry electrical current.
In diamond, all four outer electrons from each carbon atom are used for bonding. There are no free electrons available to carry electrical current.
Key Points to Remember:
- Giant covalent structures are networks of many atoms joined by covalent bonds
- Diamond has each carbon bonded to four others in a 3D structure - this makes it very hard but unable to conduct electricity
- Graphite has each carbon bonded to three others in layers - this makes it soft and able to conduct electricity
- Both have high melting points because of strong covalent bonds
- Delocalised electrons in graphite allow it to conduct electricity like metals do