Ever wondered why diamond is so incredibly hard whilst pencil...
Chemistry213 views·Updated Jun 18, 2026·2 pagesUnderstanding Giant Covalent Structures
Zofia@zofia_577
1
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Diamond and Graphite - Two Sides of Carbon
Diamond is basically carbon's superhero form - each carbon atom bonds to four others, creating an incredibly strong 3D network. This structure makes diamond the hardest natural material on Earth, perfect for cutting tools and jewellery.
However, there's a trade-off. Because all electrons are locked up in bonds, diamond can't conduct electricity at all. It also has a ridiculously high melting point because breaking those covalent bonds requires massive amounts of energy.
Graphite takes a completely different approach. Each carbon atom only bonds to three others, forming flat layers of hexagonal rings that look a bit like chicken wire. The fourth electron from each atom is free to move around, which means graphite conducts electricity brilliantly.
Quick Tip: Remember that graphite's layers can slide over each other because there are no bonds between them - that's why it works so well in pencils and as a lubricant!
2
of 2
Modern Carbon Wonders - Graphene and Fullerenes
Graphene is essentially a single layer of graphite that's been peeled off. Don't let its thinness fool you though - it's incredibly strong because of those covalent bonds holding the carbon atoms together. Plus, it conducts electricity amazingly well thanks to those delocalised electrons zipping across its surface.
Fullerenes are carbon's football-shaped cousins. These hollow molecules are made of carbon atoms arranged in hexagonal rings (and sometimes pentagons or heptagons). The most famous is buckminsterfullerene, which contains exactly 60 carbon atoms arranged like a football.
Unlike diamond and graphite, buckminsterfullerene molecules are held together by weak forces between molecules rather than strong covalent bonds. This makes them slippery with low melting points - perfect for applications like lubricants and drug delivery systems.
Remember: All these structures use the same carbon atoms, but their different arrangements give them completely different properties - it's chemistry's way of showing that structure really does determine function!
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Chemistry213 views·Updated Jun 18, 2026·2 pagesUnderstanding Giant Covalent Structures
Zofia@zofia_577
Ever wondered why diamond is so incredibly hard whilst pencil graphite is soft and slippery? It's all down to how carbon atoms bond together in different ways. These giant covalent structures show how the same element can create materials with...
1
of 2
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Diamond and Graphite - Two Sides of Carbon
Diamond is basically carbon's superhero form - each carbon atom bonds to four others, creating an incredibly strong 3D network. This structure makes diamond the hardest natural material on Earth, perfect for cutting tools and jewellery.
However, there's a trade-off. Because all electrons are locked up in bonds, diamond can't conduct electricity at all. It also has a ridiculously high melting point because breaking those covalent bonds requires massive amounts of energy.
Graphite takes a completely different approach. Each carbon atom only bonds to three others, forming flat layers of hexagonal rings that look a bit like chicken wire. The fourth electron from each atom is free to move around, which means graphite conducts electricity brilliantly.
Quick Tip: Remember that graphite's layers can slide over each other because there are no bonds between them - that's why it works so well in pencils and as a lubricant!
2
of 2Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Modern Carbon Wonders - Graphene and Fullerenes
Graphene is essentially a single layer of graphite that's been peeled off. Don't let its thinness fool you though - it's incredibly strong because of those covalent bonds holding the carbon atoms together. Plus, it conducts electricity amazingly well thanks to those delocalised electrons zipping across its surface.
Fullerenes are carbon's football-shaped cousins. These hollow molecules are made of carbon atoms arranged in hexagonal rings (and sometimes pentagons or heptagons). The most famous is buckminsterfullerene, which contains exactly 60 carbon atoms arranged like a football.
Unlike diamond and graphite, buckminsterfullerene molecules are held together by weak forces between molecules rather than strong covalent bonds. This makes them slippery with low melting points - perfect for applications like lubricants and drug delivery systems.
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