AQA · GCSE · Combined Science
Course version: 8464 · Version from Sept 2018
In diamond, each carbon atom forms four covalent bonds with other carbon…
Full specification wording
Diamond: In diamond, each carbon atom forms four covalent bonds with other carbon atoms in a giant covalent structure, so diamond is very hard, has a very high melting point and does not conduct electricity.
Diamond is a giant covalent substance made of carbon atoms.
Checked against the AQA 8464 specification (opens in a new tab) on .
- Point 1 of 2
- 2 key terms
Study notes
From the checked teaching notes for this specification point.
The idea
Diamond is a giant covalent substance made of carbon atoms. 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.
What earns the marks
- 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.
Common mistakes and how to put them right
Stating that diamond conducts electricity because it contains carbon.
Put it right 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.
Put it right by describing the rigid giant covalent lattice and the many strong covalent bonds that make it hard.
Confusing diamond with graphite.
Put it right 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.
Exam 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.
Try the lesson
The start of the Diamond: In diamond, each carbon atom forms four covalent bonds with other carbon atoms in a giant covalent structure, so diamond is very hard, has a very high melting point and does not conduct electricity. lesson, made from checked teaching notes for this specification point.
Have a go first
How many covalent bonds does each carbon atom form in diamond?
Listen
Diamond is a giant covalent substance made of carbon atoms. 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. A strong answer earns marks for these 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 common slip: Stating that diamond conducts electricity because it contains carbon. Instead, explaining that all four outer electrons per carbon atom are used in covalent bonds, leaving no free electrons to carry charge. Exam tip: Always link the number of covalent bonds per carbon atom to the properties of hardness and high melting point.
Key words
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Structure and properties of diamond
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Electrical conductivity and free electrons
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See every definition
- Structure and properties of diamond
- A rigid three-dimensional network of strong covalent bonds makes diamond very hard and gives it a very high melting point.
- Electrical conductivity and free electrons
- Diamond does not conduct electricity because there are no free electrons or ions to carry charge.
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