Diamond & Graphite

• What are the two allotropes of carbon discussed in today's video?

  Diamond and graphite
• What is an allotrope?

  Different structural forms of the same element in the same physical state
• What are some examples of allotropes of carbon?

  Diamond, graphite, and fullerenes
• What type of structure do both diamond and graphite have?

  Giant covalent structures
• How are the atoms arranged in diamond?

  Each carbon atom is covalently bonded to four other carbon atoms
• Why is diamond very strong and has a high melting point?

  Because it has strong covalent bonds that require a lot of energy to break
• Why doesn't diamond conduct electricity?

  Because it has no free electrons or ions that can move
• How does the bonding in graphite differ from that in diamond?

  Each carbon in graphite is bonded to only three other carbon atoms
• What is the arrangement of atoms in graphite?

  Atoms are arranged into hexagons forming large flat sheets
• Why is graphite relatively soft compared to diamond?

  Because the layers in graphite are held together weakly and can slide over one another
• What allows graphite to conduct electricity and heat?

  The presence of delocalized electrons that are free to move
• What is a single layer of graphite called?

  Graphene
• What can scientists do with isolated layers of graphene?

  They can use them to make other structures such as spheres and tubes
• What are the key differences between diamond and graphite?

  • Diamond:
  • Each carbon bonded to four others
  • Strong covalent bonds
  • High melting point
  • Does not conduct electricity
  • Graphite:
  • Each carbon bonded to three others
  • Arranged in layers of hexagons
  • Layers can slide over each other
  • Conducts electricity due to delocalized electrons
• What will be covered in the next video about carbon allotropes?

  • Graphene
  • Fullerenes
  • Their structures and properties