Diamond and Graphite notes

Allotropes of Carbon: Diamond and Graphite

Introduction to Allotropes

• Allotropes are different structural forms of the same element in the same physical state.
• Carbon in the solid state can exist as diamond, graphite, fullerenes, etc.

Giant Covalent Structures

• Both diamond and graphite are giant covalent structures: regular lattice of covalently bonded atoms.
• This makes them both very strong.

Diamond

• Each carbon atom is covalently bonded to four other carbon atoms (the maximum number of bonds carbon can make).
• Forms a regular three-dimensional pattern.
• Each covalent bond is very strong, so it takes loads of energy to break them.
• Properties:
  • Very strong
  • High melting point
  • Does not conduct electricity because it has no free electrons or ions that can move around.

Graphite

• Each carbon atom is bonded to only three other carbon atoms.
• Atoms are arranged into hexagons, forming large flat sheets.
• Sheets are arranged on top of one another to form layers.
• Individual layers are only held together weakly (no covalent bonds between them).
• Layers are free to slide over one another, making graphite relatively soft compared to diamond.
• Properties:
  • High melting point (individual layers are strongly held together with covalent bonds)
  • Each carbon atom has one spare electron (not used in bonding).
  • Spare electron becomes delocalized (free to move around).
  • Conducts electricity and heat (due to free electrons).

Graphene and Fullerenes

• A single layer of graphite is known as graphene.
• Scientists can isolate these individual layers and use them to make other structures, such as spheres and tubes, called fullerenes.

Allotropes of Carbon: Diamond and Graphite ### Introduction to Allotropes - Allotropes are different structural forms of the same element in the same physical state. - Carbon in the solid state can exist as diamond, graphite, fullerenes, etc. ### Giant Covalent Structures - Both diamond and graphite are giant covalent structures: regular lattice of covalently bonded atoms. - This makes them both very strong. ### Covalent Bonds - Covalent bonds are chemical bonds that involve the sharing of electron pairs between atoms. - These bonds occur when atoms share electrons in order to achieve a stable electron configuration, typically resembling that of a noble gas. - Covalent bonds are strong bonds, and substances held together by covalent bonds often have high melting and boiling points. ### Diamond - Each carbon atom is covalently bonded to four other carbon atoms (the maximum number of bonds carbon can make). - Forms a regular three-dimensional pattern. - Each covalent bond is very strong, so it takes loads of energy to break them. - ">>Properties:":

• Very strong - High melting point - Does not conduct electricity because it has no free electrons or ions that can move around. ### Graphite - Each carbon atom is bonded to only three other carbon atoms. - Atoms are arranged into hexagons, forming large flat sheets. - Sheets are arranged on top of one another to form layers. - Individual layers are only held together weakly (no covalent bonds between them). - Layers are free to slide over one another, making graphite relatively soft compared to diamond. - ">>Properties:":
• High melting point (individual layers are strongly held together with covalent bonds) - Each carbon atom has one spare electron (not used in bonding). -