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Diamond melts at over 3500°C because melting it means breaking the strong covalent bonds in its giant structure. Free notes and 11 questions answered.
The Structure and Properties of Diamond is part of Structure and bonding of carbon in GCSE AQA Higher Chemistry.
3 min
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GCSE AQA Higher
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Diamond melts at over 3500°C because melting it means breaking the strong covalent bonds in its giant structure. Free notes and 11 questions answered.
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Use this lesson to stay inside Structure and bonding of carbon while connecting the detail back to the wider Topic 2: Bonding, structure, and the properties of matter topic.
Read the core explanations from The Structure and Properties of Diamond before testing yourself from memory.
Diamond is the hardest natural substance on Earth. You can scratch glass with it, be used in diamond-tipped tools that cut through steel, and it won't melt until temperatures reach over 3500°C. But here's the puzzle - diamond is made entirely of carbon atoms, the same element that makes up soft graphite in your pencil. How can the same atoms create such different materials?
The answer lies in how those carbon atoms connect to each other. In diamond, every single carbon atom forms four strong covalent bonds with its neighbours, creating a three-dimensional network that extends throughout the entire crystal. This arrangement gives diamond its extraordinary properties.
Picture a carbon atom bonded strongly to four other carbon atoms. This forms a giant three-dimensional covalent structure in which each carbon atom is joined to four neighbours.
But here's where it gets interesting - each of those four neighbouring carbon atoms also forms four bonds with other carbon atoms. This pattern repeats throughout the entire crystal, creating a giant three-dimensional network. Every carbon atom is covalently bonded to exactly four others, and this continues in all directions.
Now we can explain why diamond behaves so differently from other materials. The giant covalent structure creates three key properties that make diamond special.
Diamond is extremely hard because breaking or scratching it means breaking covalent bonds. Since every carbon atom is locked in place by four strong covalent bonds, enormous force is needed to move the atoms out of position. The three-dimensional network means there are no weak points - stress gets distributed throughout the entire structure.
Diamond has a very high melting point (over 3500°C) for the same reason. To melt diamond, you need to break the covalent bonds holding the structure together. These bonds are much stronger than the weak intermolecular forces in simple molecular compounds, so much more energy is required.
Diamond does not conduct electricity because all the electrons are used in covalent bonding. There are no free electrons available to carry electric current through the structure. Each carbon atom uses all four of its outer electrons to form bonds, leaving none free to move.
Understanding diamond becomes clearer when you compare it with graphite, another form of pure carbon. Both contain only carbon atoms, but their properties are completely different because of how those atoms are arranged.
Worked example: Comparing Diamond and Graphite Structures
Explain why diamond is hard while graphite is soft, even though both are made of carbon atoms.
Answer: Diamond is hard because its 3D covalent network resists deformation in all directions, while graphite is soft because its layers can slide over each other despite strong bonding within each layer.
Diamond's unique combination of properties makes it valuable for specific applications. Industrial diamonds are used in cutting tools, drill bits, and abrasives precisely because of their hardness. The structure-property relationship we've studied explains why diamond works so well in these roles.
In exam questions about diamond, you'll often be asked to explain its properties in terms of structure and bonding. The key is to make clear connections: start with the bonding (four covalent bonds per carbon atom), describe the structure (giant 3D network), then explain how this leads to the properties (hardness, high melting point, electrical insulation).
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Move to the related lessons or back to the topic page when you need the wider sequence before exam-style practice.
Diamond is extremely hard, has a very high melting point of over 3500°C and does not conduct electricity. These properties come from its giant covalent structure, in which every carbon atom is held by four strong covalent bonds and no electrons are free to move.
Diamond has a very high melting point, over 3500°C, because melting it means breaking the strong covalent bonds that hold its giant structure together. These bonds are much stronger than weak intermolecular forces, so much more energy is required.
Diamond melts at over 3500°C. Melting it means breaking the strong covalent bonds that hold its giant structure together, which needs far more energy than breaking the weak intermolecular forces in simple molecular compounds.
Diamond is extremely hard because each carbon atom is joined to four others by strong covalent bonds in a giant three-dimensional network. Scratching or breaking diamond means breaking covalent bonds, so enormous force is needed and there are no weak points.
Industrial diamonds are used in cutting tools, drill bits and abrasives because diamond is the hardest natural substance. Diamond-tipped tools can cut through steel, and diamond can scratch glass, because breaking its four covalent bonds per carbon atom needs enormous force.
Diamond has a giant covalent structure: a three-dimensional network in which every carbon atom is covalently bonded to exactly four others. This pattern repeats in all directions throughout the crystal, linking millions of carbon atoms in a continuous lattice.
Each carbon atom in diamond forms four strong covalent bonds with four neighbouring carbon atoms. All four outer electrons are used in bonding, so every atom is locked in place in a rigid three-dimensional network with no weak points.
Diamond is hard because each carbon atom forms 4 covalent bonds in a rigid 3D network that resists deformation in all directions. Graphite is soft because each carbon atom forms 3 covalent bonds in flat layers, and the weak forces between layers let them slide.
Diamond is used in cutting tools, drill bits and abrasives because of its hardness, which comes from its giant covalent structure. Every carbon atom is locked in place by four strong covalent bonds, so enormous force is needed to scratch or deform it.
Diamond does not conduct electricity because each carbon atom uses all four of its outer electrons to form covalent bonds. There are no free electrons available to carry electric current through the structure.
Diamond rates 10 on the Mohs hardness scale, the maximum value, making it the hardest natural substance. Scratching it means breaking covalent bonds, and every carbon atom is locked in place by four strong covalent bonds, so enormous force is needed.