Diamond/Graphite

• Giant covalent structures contain very many atoms, each joined to adjacent atoms by covalent bonds. The atoms are usually arranged into giant regular lattices – extremely strong structures because of the many bonds involved.
• Properties of Giant Covelant Compounds:

  • Very high melting points – this is because a lot of strong covalent bonds must be broken. Graphite, for example, has a melting point of more than 3,600°C.
  • Variable electrical conductivity – diamond does not conduct electricity, whereas graphite contains free electrons so it does conduct electricity. Silicon is a semi-conductor – it is midway between non-conductive and conductive.
• Graphite:

  
• Diamond:

  
• Diamond:

  
  

  • Very hard and has a high melting point
  • It does not conduct electricity as there are no delocalised electrons in the structure.
  • It is an insulator
  • No free electrons
  • Very hard structure because all of the 4 outer atoms are bonded to other 4 carbon atoms (goes on and on)
  • Uses of diamond= Jewellery, adds hardness to tools etc...
  • Shape= Tetrahedral structure (4)
• Graphite:

  • Hexagonal rings (6 sides)
  • Graphite is a form of carbon in which the carbon atoms form covalent bonds with three other carbon atoms
  •  each carbon atom has a ‘spare’ electron (as carbon has four outer electrons) which is delocalised between layers of carbon atoms meaning it's got capacity to travel electricity
  • These layers can slide over each other, so graphite is much softer than diamond.
• Graphite:

  • used in pencils, and as a lubricant
  • Graphite conducts electricity due to the ‘spare’ electrons being delocalised between the layers
  • This conductivity makes graphite useful as electrodes for electrolysis.
  • However, graphite still has a very high melting and boiling point because the strong covalent bonds that hold the carbon atoms together in the layers require a lot of heat energy to break.