Chemistry
Graphite Learning objectives: • describe the structure and bonding of graphite • explain the properties of graphite • explain the similarity to metals.
Graphite is a carbon structure that is black and slippery and sticks to paper. It makes good ‘lead’ for pencils. The slipperiness also makes graphite a good lubricant, even though it is a solid. Powdered graphite is often used to lubricate door locks and is also used in furnaces and as brake linings. It is a good electrical conductor so it is used in batteries and as powder in carbon microphones.
Graphite layers Like diamond, graphite is also made from carbon atoms. In graphite the carbon atoms are arranged differently. The carbon atoms only make three bonds with other carbon atoms. The carbon atoms bond to make six-sided rings. The rings stack to make layers.
KEY WORDS delocalised electrons graphite hexagonal rings weak bonds
DID YOU KNOW? Graphite pencils were first made in 1565. All the graphite came from a mine near the Honister Pass in Cumbria. Clay was added to make different grades of pencil hardness. The graphite in pencils was called ‘lead’ because the Roman writing implement, the stylus (used for writing on wax tablets), was made of the metal lead.
Graphite can be easily cut across the layers but not through the layers. 1
State the number of bonds that carbon makes in (a) diamond and (b) graphite.
2
Graphite can be used as a lubricant. State what this tells you about the forces between layers.
Structure and properties In graphite, each carbon atom forms three strong covalent bonds with three other carbon atoms, forming layers of hexagonal rings which have no covalent bonds between the layers.
Figure 2.44 Graphite is slippery and can leave marks on paper.
As the bonds need a lot of energy to break, graphite has a high melting point. Only weak bonds hold the layers together so the layers are free to slide over each other. There are no covalent bonds between the layers and so graphite is soft and slippery. In graphite, only three electrons from each carbon atom form strong covalent bonds with electrons from other carbon atoms. The remaining one electron of each carbon atom is delocalised. These delocalised electrons allow graphite to conduct electricity.
Figure 2.45 Graphite is made of hexagonal rings that stack as layers.
The delocalised electrons move easily along the layers. This is similar to the way that delocalised electrons move in metals. This is why both graphite and metals can conduct electricity. Diamond has no delocalised electrons so cannot conduct electricity.
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AQA GCSE Chemistry for Combined Science: Trilogy: Student Book
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5/24/16 8:52 PM