C2.3 Comparison of Graphite and Diamond

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  • Created by: Sophie
  • Created on: 22-11-14 15:03

C2.3 Comparison of Graphite and Diamond

Similarities

  • Both are giant covalent structures (macromolecules)
  • Both formed of only carbon atoms
  • Strong covalent bonds between atoms

Differences

  • Diamond has a regular three-dimensional giant structure, whilst graphite has a giant two-dimensional structure
  • One electron from each carbon atom in graphite is delocalised. These delocalised electrons mean that graphite can conduct electricity and heat. As there are no delocalised electrons in diamond, diamond is not a conductor of electricity or heat
  • Whilst diamond forms a 3D structure, graphite forms a 2D one. There are weak intermolecular forces between the layers in graphite so the layers can slide over each other. This makes graphite slippery and grey
  • In graphite, every C atom is bonded covalently to 3 C atoms, whilst in diamond, every C atom is bonded covalently to 3 C atoms

Overall comparison

Graphite can conduct electricity and heat whilst diamond cannot. The layers in graphite slide over each other. Graphite has a sea of delocalised electrons but diamond does not. Graphite is used in lubricants and pencils, diamonds are used in jewellery and industrial drills (because it is a hard material). Diamond is hard and clear whilst graphite is slippery and grey. Diamond is a good insulator of heat and electricity, whilst graphite is not.

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