Diamond — AQA GCSE Combined Science
Test yourself on Diamond with AQA GCSE practice questions.
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Diamond explained
Diamond is a giant covalent substance made of carbon atoms.
Read the full explanation
Each carbon atom forms four strong covalent bonds to four other carbon atoms, arranged tetrahedrally. This continuous three-dimensional network means there are no free electrons or ions, so diamond does not conduct electricity. The many strong covalent bonds must be broken for melting, giving diamond a very high melting point. The rigid network also makes diamond extremely hard. For example, diamond is used in drill tips and cutting tools. Its hardness and high melting point arise from the strength and number of covalent bonds, while its electrical insulating behaviour is due to the absence of delocalised electrons.
Students should be able to explain the properties of diamond in terms of its structure and bonding.
Diamond is a giant covalent lattice of carbon atoms. Every carbon atom forms four single covalent bonds, arranged tetrahedrally, so each atom is joined to four others and the whole crystal is one continuous network. There are no separate molecules and no free electrons. Melting requires breaking many strong covalent bonds, so the melting point is extremely high. Those same strong bonds resist deformation, making diamond very hard. Because all four outer electrons per carbon are locked into bonding, there are no delocalised electrons, so diamond does not conduct electricity. A useful method is to link each property to a structural feature: high melting point to many strong covalent bonds; hardness to a rigid three-dimensional network; non-conduction to the absence of delocalised electrons.
Your focus
- Describe the structure of diamond in terms of carbon atoms and covalent bonds.
- Explain why diamond is very hard and has a very high melting point.
- Explain why diamond does not conduct electricity.
Show all 6 objectives
- Describe the giant covalent lattice structure of diamond, including that each carbon atom forms four covalent bonds.
- Explain how the number and strength of covalent bonds in diamond account for its very high melting point and hardness.
- Explain why diamond does not conduct electricity in terms of the absence of delocalised electrons.
Diamond exam tips
Marking Points
- Diamond is a giant covalent structure of carbon atoms.
- Each carbon atom forms four covalent bonds with other carbon atoms.
- The bonds are arranged tetrahedrally, creating a rigid three-dimensional network.
- All outer electrons are involved in covalent bonding, so there are no delocalised electrons or ions to carry charge.
- A large amount of energy is needed to break the many strong covalent bonds, so diamond has a very high melting point.
- The rigid network of strong covalent bonds makes diamond very hard.
- Diamond is a giant covalent (macromolecular) lattice in which each carbon atom is covalently bonded to four other carbon atoms.
- The bonds are strong single covalent bonds arranged tetrahedrally, giving a rigid three-dimensional network with no separate molecules.
- A very large amount of energy is needed to break the many strong covalent bonds, so diamond has a very high melting point.
- The rigid network of strong bonds makes diamond extremely hard, which is why it is used in cutting tools and drill tips.
- All four outer electrons of each carbon atom are used in bonding, so there are no delocalised electrons and diamond does not conduct electricity.
- A good answer links each property explicitly to a structural feature rather than listing properties alone.
Examiner Tips
- 💡Always link the number of covalent bonds per carbon atom to the properties of hardness and high melting point.
- 💡For electrical conductivity, explicitly state that there are no free or delocalised electrons in diamond.
- 💡Use diagrams to show the tetrahedral arrangement and the continuous network when explaining properties.
- 💡For each property you give, immediately state the structural reason, for example high melting point because many strong covalent bonds must be broken.
- 💡Use precise terms such as giant covalent lattice, tetrahedral and delocalised electrons rather than vague words like strong structure.
- 💡If asked to compare with graphite, refer back to the stem so your answer stays focused on diamond's structure and bonding.
Common Mistakes
- Stating that diamond conducts electricity because it contains carbon: correct this by explaining that all four outer electrons per carbon atom are used in covalent bonds, leaving no free electrons to carry charge.
- Thinking diamond is soft because it is a non-metal: correct this by describing the rigid giant covalent lattice and the many strong covalent bonds that make it hard.
- Confusing diamond with graphite: correct this by noting that diamond has four bonds per carbon atom in a tetrahedral network, whereas graphite has three bonds per carbon atom in layers with delocalised electrons.
- Saying diamond is made of small molecules: correct this by stating it is a giant covalent lattice with no separate molecules.
- Claiming diamond conducts electricity because it contains carbon: correct this by explaining that all outer electrons are localised in covalent bonds, so there are no delocalised electrons.
- Stating that diamond is hard because the bonds are short: correct this by linking hardness to the rigid three-dimensional network of strong covalent bonds.