Diamond & Graphite (Cambridge (CIE) IGCSE Chemistry)
Revision Note
Written by: Caroline Carroll
Reviewed by: Stewart Hird
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Structure of graphite & diamond
What is the structure of 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
The structure and bonding in graphite
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
What is the structure of 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 to show the formation of a covalent bond
Diagram showing the structure and bonding arrangement in diamond
Uses of graphite & diamond
What are the properties and uses of graphite?
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!
What are the properties and uses of diamond?
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
Diamond has the following physical properties:
It does not conduct electricity
It has a very high melting point
It is extremely hard and dense
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.
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