AQA GCSE Combined Science (8464), Higher tier · Chemistry › Bonding, structure, and the properties of matter › Structure and bonding of carbon
Practise Graphite. 14 exam-style questions on this subtopic, at up to four difficulty levels, with full mark schemes and a progress tracker. Free, no account needed.
Graphite is a form of carbon made of layers of hexagonal rings. You need to explain why it is soft and slippery, has a very high melting point and conducts electricity, and why it is similar to metals. A 4 to 6 mark question comparing graphite with diamond is common.
Grade by grade
What you need to be able to do, from the first marks up to the top grade.
3
Describe graphite as layers of carbon atomsGraphite is made of layers of carbon atoms arranged in hexagonal rings.
4
State the number of bonds each carbon formsEach carbon atom forms three covalent bonds with three other carbon atoms.
5
Explain why graphite is soft and slipperyThere are no covalent bonds between the layers, so the layers can slide over each other.
6
Explain why graphite conducts electricityOne electron from each carbon atom is delocalised and can move through the structure, carrying charge.
6
Explain graphite's very high melting pointThe many strong covalent bonds within the layers must be broken.
8
Compare the structures and properties of diamond and graphiteFour bonds in a rigid 3D network compared with three bonds, layers and delocalised electrons.
Notes
Structure and bonding
Graphite is made of carbon atoms only, in a giant covalent structure.
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, only weak forces.
Each carbon atom has four outer electrons but uses only three in bonds. The fourth electron is delocalised: one electron from each carbon atom is free to move through the structure.
Properties explained
Soft and slippery: the forces between the layers are weak, so the layers can slide over each other. This is why graphite is used in pencils and as a lubricant.
Very high melting point: to melt graphite, the strong covalent bonds within the layers must be broken. There are many of them, so a lot of energy is needed.
Conducts electricity (and thermal energy): the delocalised electrons can move through the structure, carrying charge.
Graphite is similar to metals because it has delocalised electrons. This is one reason it is used for electrodes.
Graphite compared with diamond
Both are giant covalent structures of carbon atoms, and both have very high melting points.
Diamond: 4 bonds per atom in a rigid 3D network → very hard; no delocalised electrons → does not conduct.
Graphite: 3 bonds per atom, in layers with weak forces between them → soft; one delocalised electron per atom → conducts.
Cheatsheet
Graphite: each carbon forms 3 covalent bonds → layers of hexagonal rings
No covalent bonds between layers (only weak forces) → layers slide → soft and slippery
One electron per carbon atom is delocalised → conducts electricity and thermal energy
Melting breaks strong covalent bonds within the layers → very high melting point
Like metals, graphite has delocalised electrons
Uses: pencils, lubricants, electrodes
How to answer each type of question
Explain why graphite is soft and slippery
3 marks5
Say graphite is made of layers (of hexagonal rings of carbon atoms).
Say there are only weak forces (no covalent bonds) between the layers.
Say the layers can slide over each other.
Example. When you write with a pencil, layers of graphite rub off onto the paper. Explain why graphite is soft.
Show the model answer
Graphite is made of layers of carbon atoms (1). There are only weak forces between the layers / no covalent bonds between the layers (1). So the layers can easily slide over each other (1).
Explain why graphite conducts electricity
3 marks6
Say each carbon atom forms only three covalent bonds.
So one electron from each carbon atom is delocalised.
The delocalised electrons can move through the structure and carry charge.
Example. Graphite is used to make electrodes. Explain why graphite conducts electricity.
Show the model answer
Each carbon atom forms only three covalent bonds (1), so one outer electron from each carbon atom is delocalised (1). The delocalised electrons can move through the structure and carry charge (1).
Explain why graphite is soft but has a very high melting point
3 marks7
Softness: layers slide because the forces between the layers are weak.
Melting: the strong covalent bonds within the layers must be broken.
There are many of these bonds, so a lot of energy is needed.
Example. Graphite is soft, but it has a very high melting point. Explain why.
Show the model answer
Graphite is soft because there are only weak forces between the layers, so the layers slide over each other (1). To melt graphite, the strong covalent bonds within the layers must be broken (1). There are many of these bonds, so a lot of energy is needed (1).
6-mark: compare diamond and graphite
6 marks8
Structure: both are giant covalent structures of carbon; diamond has 4 bonds per atom in a 3D network, graphite has 3 bonds per atom in layers.
Hardness: diamond is rigid, so hard; graphite's layers slide, so it is soft.
Conductivity: diamond has no delocalised electrons; graphite has one per carbon atom.
Write point by point, linking each property to the structure, and mention both substances each time.
Example. Diamond and graphite are both forms of carbon. Compare the structure and bonding of diamond and graphite, and explain how they affect the hardness and electrical conductivity of each substance.
Show the model answer
Marked in levels. A top-level answer (5–6 marks) links structure to both properties for both substances, using points such as: • both are giant covalent structures made of carbon atoms • in diamond each carbon atom forms four covalent bonds; in graphite each forms three • graphite has layers of hexagonal rings, with weak forces (no covalent bonds) between the layers • diamond is very hard because it is a rigid 3D network of strong covalent bonds • graphite is soft because its layers can slide over each other • graphite has one delocalised electron per carbon atom, which can move and carry charge, so it conducts electricity • diamond has no delocalised electrons, so it does not conduct Level 2 (3–4 marks): some correct links between structure and properties. Level 1 (1–2 marks): simple correct statements with few links.
Pencil 'lead' is graphite: its layers slide off onto the paper, which reminds you why graphite is soft.
Soft but hard to melt: sliding only overcomes the weak forces between layers; melting has to break covalent bonds.
Where marks are lost
Saying graphite has a low melting point because it is soft. Melting breaks the strong covalent bonds within the layers.
Saying there are covalent bonds between the layers, or that the layers slide because the covalent bonds are weak.
Saying graphite conducts because of free ions, or because the layers move.
Writing that each carbon atom in graphite forms four bonds.
Calling graphite a metal. It is a non-metal that has delocalised electrons like a metal.
Exam technique
Say 'one electron from each carbon atom is delocalised', not just 'it has free electrons'.
For softness you need three ideas: layers, weak forces between them, layers slide.
In 6-mark comparisons, organise your answer by property (structure, hardness, conductivity) and mention both substances in each part.
Sample questions
Written for this site in the style of AQA exam questions. They are not taken from real past papers.
Question 1Easy4 marks
Graphite is a form of carbon.
(a) How many covalent bonds does each carbon atom form in graphite?[1]
(b) Graphite is used in pencils. Which property makes graphite suitable for this use? Tick (✓) one box.[1]
It is soft and slippery.
It is very hard.
It has a low melting point.
It is transparent.
(c) The layers in graphite are made of rings of carbon atoms. How many carbon atoms are in each ring? Tick (✓) one box.[1]
4
5
6
8
(d) Graphite conducts electricity. Which particles carry the charge through graphite?[1]
Show the answer and mark scheme
(a)Answer: 3
3
(b)Answer: It is soft and slippery.
(c)Answer: 6
(d)Answer: delocalised electrons
(delocalised) electrons
Question 2Medium6 marks
Graphite is a form of carbon.
(a) Explain why graphite is soft and slippery.[3]
(b) Explain why graphite conducts electricity.[3]
Show the answer and mark scheme
(a)Answer: Graphite is made of layers with no covalent bonds between them (only weak forces), so the layers can slide over each other.
graphite has layers (of carbon atoms)
there are no covalent bonds between the layers / only weak forces between the layers
so the layers can slide over each other
(b)Answer: Each carbon atom forms only three bonds, so one electron from each atom is delocalised; these electrons move through the structure carrying charge.
each carbon atom forms three covalent bonds
(so) one electron from each carbon atom is delocalised
the delocalised electrons move through the structure / along the layers, carrying charge
Question 3Hard6 marks
Diamond and graphite are both forms of carbon. Diamond is used to make the tips of drills. Graphite is used as a lubricant and to make electrodes.
Compare the structure and bonding of diamond and graphite. Use your answer to explain why each substance is suitable for its uses.[6]
Show the answer and mark scheme
Answer: Both are giant covalent carbon. Diamond: 4 bonds per atom, rigid 3D — very hard (drill tips), no delocalised electrons. Graphite: 3 bonds per atom, layers with no bonds between them — slippery (lubricant); one delocalised electron per atom — conducts (electrodes). Both have very high melting points.
both are giant covalent structures containing only carbon atoms
in diamond, each carbon atom forms four covalent bonds with other carbon atoms
diamond has a rigid three-dimensional structure so it is very hard, which makes it suitable for drill tips
in graphite, each carbon atom forms three covalent bonds, forming layers of hexagonal rings
there are no covalent bonds between the layers (only weak forces), so the layers slide over each other and graphite is slippery, which makes it a good lubricant
in graphite, one electron from each carbon atom is delocalised, so graphite conducts electricity and can be used for electrodes
diamond has no delocalised electrons, so it does not conduct electricity
both have very high melting points because many strong covalent bonds must be broken, so they are not damaged at high temperatures
Marked with levels of response: the full level descriptors are in the app.