Diamond and graphite: giant covalent structures of carbon

Diamond and graphite are both made only of carbon atoms joined by strong covalent bonds in giant covalent structures, but the atoms are arranged differently. In diamond, each carbon atom bonds to four others, so diamond is very hard, has a very high melting point and doesn't conduct electricity. In graphite, each carbon atom bonds to three others in layers of hexagonal rings, with one delocalised electron per atom, so graphite conducts electricity and its layers slide over each other.

What is a giant covalent structure?

A giant covalent structure is a huge network of atoms all joined by covalent bonds, with no separate molecules. To melt one, you have to break a great many strong covalent bonds, so giant covalent substances have very high melting and boiling points. AQA names diamond, graphite and silicon dioxide (silica) as examples, and expects you to recognise them from diagrams of their bonding.

The structure of diamond

In diamond, each carbon atom forms four covalent bonds with other carbon atoms, in a rigid three-dimensional network.

• Very hard: every atom is locked in place by four strong bonds, so diamond is used in cutting tools and drill tips.
• Very high melting point: melting means breaking many strong covalent bonds.
• Doesn't conduct electricity: all four outer electrons of each carbon atom are used in bonds, so there are no delocalised electrons or ions to carry charge.

The structure of graphite

In graphite, each carbon atom forms three covalent bonds with three other carbon atoms, making flat layers of hexagonal rings. There are no covalent bonds between the layers.

• Conducts electricity and heat: each carbon atom has one outer electron that isn't in a bond. These electrons are delocalised and move along the layers. AQA points out that this makes graphite similar to metals.
• Soft and slippery: the layers are held together only weakly, so they slide over each other. That is why graphite is used in pencils and as a lubricant.
• Very high melting point: melting still means breaking the strong covalent bonds within the layers.

Diamond vs graphite

Diamond, graphite, graphene and fullerenes are all forms of carbon, known as allotropes.

Graphene and fullerenes

• Graphene is a single layer of graphite, one atom thick. AQA says its properties make it useful in electronics and composites; like graphite, it has delocalised electrons.
• Fullerenes are molecules of carbon atoms with hollow shapes, based on hexagonal rings but sometimes containing rings of five or seven carbon atoms. The first discovered was buckminsterfullerene, C₆₀, which is spherical.
• Carbon nanotubes are cylindrical fullerenes with very high length to diameter ratios, useful in nanotechnology, electronics and materials.

Identify the type of bonding in a substance in our Chemical Bonding course, then press Mark to check your answer in the course:

Exam questions with answers

Explain why graphite conducts electricity but diamond does not. (4 marks)

Each carbon atom in graphite is bonded to only three others, which leaves one of its outer electrons delocalised. Delocalised electrons are free to move along the layers, so graphite can carry an electric current. In diamond, every carbon atom uses all four outer electrons in covalent bonds, so no electrons are free to move.

Why does diamond need such a high temperature to melt? (3 marks)

Diamond is a giant covalent network: each carbon atom is held to four neighbours by strong covalent bonds. To melt it, a huge number of these bonds must break, and that takes a great deal of energy.

Why is graphite soft and slippery?

Graphite is made of layers with no covalent bonds between them. The forces between the layers are weak, so the layers can slide over each other.

Common mistakes

• Saying intermolecular forces are broken when diamond melts. Diamond has no molecules; melting breaks covalent bonds.
• Saying graphite conducts because of free ions. It is delocalised electrons, not ions.
• Saying the covalent bonds in graphite are weak. The bonds within layers are strong; it is the forces between layers that are weak.
• Forgetting the numbers. Four bonds per atom in diamond, three in graphite, and say so in your answer.

Practise with a short timed paper

Diamond, graphite and the other forms of carbon are in GCSE Chemistry. The daily GCSE Chemistry Higher paper is a short timed paper of 8 questions, with a new one every day. For how the type of bonding decides a substance's properties, see our Chemical Bonding course.

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Chemistry · Bonding and structureWhy are metals good conductors of electricity?Why metals are good conductors of electricity and heat: delocalised electrons in metallic bonding, how they carry charge, and why other materials don't.