4. Carbon allotropes - comparing diamond, graphite and graphene

• DIAMOND C_n where n is an extremely large number of carbon atoms!
  • Diamond is an allotrope of carbon - a distinct molecular form of solid carbon (1st allotrope of carbon discussed).
    • Note: Allotropes are defined as different physical forms of the same element in the same physical state, so the atoms in allotropes are identical, but bonded in different ways, as you will see with the four forms of carbon.
  • In diamond every carbon atom is strongly linked to four other carbon atoms by strong directional single covalent bonds giving a very three dimensional (3D) strong lattice. This results in a very rigid strong structure.
    • Theoretically in a diamond crystal, all the carbon atoms are linked together - as shown in the ball and stick diagram of diamond on the right.
    • This is a 3D tetrahedral lattice of carbon atoms.
    • The result is a very pure almost colourless crystal structure with a high refractive index that gives diamonds a lustrous look - they give quite a sparkle as light passes through it.
  • The strength and hardness of carbon in the form of diamond enables it to be used as the 'leading edge' on cutting tools, the hardness is derived from the very strong rigid three–dimensional carbon–carbon bond network.
    • It is possible to manufacture synthetic industrial diamonds, though not as perfect or as valuable as naturally occurring diamonds, they are just as useful and cheaper for use in cutting tools.
  • Diamond also has a very high melting point because of this very strong giant covalent lattice in which every carbon atom is strongly bonded to four other carbon atoms (see diagram above on right).
    • It takes a lot of energy to break (overcome) the carbon-carbon bonds and cause melting.
    • The more thermal kinetic energy needed, the higher the melting point.
    • Under very high pressure, diamond will melt at a staggering 4700^oC!
  • The strong bond network in diamond (and graphite and silica) prevents these materials from dissolving in any conventional solvent.
  • Diamond does not conduct electricity  - there are no free-delocalised electrons or ions that can move through the structure to carry an electric current.
  • Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances is given in a section of the Energetics Notes.

• GRAPHITE consists of multilayers of C_n sheets where n is an extremely large number of carbon atoms all joined together!
  • Graphite is another form of carbon (2nd allotrope of carbon to be discussed)
    • The carbon atoms form joined hexagonal rings forming layers 1 atom thick in graphite.
    • This is a 2D hexagonal lattice of carbon atoms - joined up hexagonal rings of carbon atoms.
    • Each carbon atom is strongly covalently bonded to three other carbon atoms.
    • So every carbon atom in a layer is strongly bonded to three other carbon atoms.
    • A crystal of graphite contains millions of layers of these sheets of carbon atoms.
    • Although graphite is almost black and opaque (unlike diamond), it does look a bit shiny and smooth.
  • There are three strong covalent bonds per carbon atom in graphite (3 C–C bonds in a planar arrangement from 3 of its 4 outer electrons). So three of the electrons are tightly held in three directed covalent bonds, BUT, the fourth outer electron is 'delocalised', that is one electron per carbon atom is shared between the carbon atoms to form the equivalent of a 4th bond per carbon atom.
    • This 4th electron is free to move around ('delocalised'), hence graphite's ability to conduct electricity, giving graphite a similarity with metals.
    • This electrical conductivity of carbon in the form of its graphite allotrope is quite unusual for giant covalent structures.
    • This situation requires advanced level concepts to fully explain the structure of graphite, and this bonding situation also occurs in fullerenes described below, and in aromatic compounds you deal with only at advanced level.
  • The layers are only held together by weak intermolecular forces indicated by the dotted lines NOT by strong covalent bonds, so graphite, for a giant covalent structure, is unusually physically weak.
    • There are no strong covalent bonds between carbon atoms of adjacent layers.
  • Like diamond and silica (above) the large molecules of the layer ensure graphite has typically very high melting point because of the strong 2D bonding network (note: NOT a 3D network).
    • It takes a lot of energy to break (overcome) the carbon-carbon bonds within individual layers of graphite, hence its very high melting point.
  • Graphite will not dissolve in solvents because of the strong bonding in the layers.
  • BUT there are two crucial differences compared to diamond ...
    • The 4th shared electron, can move freely through each layer, so graphite is a good conductor of heat AND electricity like a metal (unusual for a non-metal).
      • Contrasting with diamond which is an electrical insulator and a poor thermal conductor.
      • For a non-metal, graphite is a relatively good conductor of heat and electricity, which gives it some similarity with metals.
        • The electric charge OR thermal kinetic energy is quickly conveyed by the delocalised electrons.
      • Graphite is used in electrical contacts e.g.
        • brushes in electric motors,
        • electrodes in electrolysis and can withstand high temperatures e.g. in the extraction of aluminium - the electrolysis of aluminium oxide (bauxite ore) requires quite a high temperature.
    • The weak forces enable the layers to slip over each other so where as diamond is hard material graphite is a 'soft' crystal, it feels slippery.
      • Graphite can act as a lubricant for the same reason, the slipperiness of the layers!
      • This enables graphite to be used as a lubricant in its own right, but it is also an important additive that is sometimes added to other oil based lubricants to enhance their performance.
        • Graphite is used in some types of brake linings - thermally very stable and a good conductor of heat (also unusual for a non-metal).
      • Carbon in the form of graphite is the only non–metal that is a significant electrical conductor.
      • Graphite is used in pencils (often wrongly called lead pencils!) because the weak structure allows the graphite layers to slide off onto paper when pressure is applied on rubbing the pencil over paper.
  • These two different characteristics of graphite described above are put to a common use with the electrical contacts in electric motors and dynamos.
    • These contacts (called brushes) are made of graphite sprung onto the spinning brass contacts of the armature.
    • The graphite brushes provide good electrical contact and are self–lubricating as the carbon layers can slide over each other on the rotating metal contacts.
• The difference in structures of diamond and graphite is also highlighted by the differences in density.
  • Carbon (graphite) is 2.25 g/cm³ (2250 kg/m³)
  • Carbon (diamond) is 3.51 g/cm³ (3510 kg/m³)

Graphene consists of a single layer of graphite and has properties that make it useful in electronics and composites.

• GRAPHENE

  • Graphene is a single sheet of graphite, therefore it is another form of the non-metallic element carbon.

    • You can consider it as a 3rd allotrope of carbon to be described here.

    • So, again, this is a 2D hexagonal lattice of carbon atoms - joined up hexagonal rings of carbon atoms.

  • Graphene is a 2D nanomaterial because each layer is only one carbon atom thick made up of a huge number of linked hexagonal rings of carbon atoms.

  • Like graphite, graphene conducts electricity because delocalised electrons can readily run through the one atom thick layer - its the 4th of carbon's outer electrons that is free to move through the network of carbon atoms.

  • Because of the strong carbon-carbon bonding, it is a very strong light-weight material with a higher tensile strength than steel.

  • Each layer is only one carbon atom in thickness and is transparent to light.

  • Graphene is actually a better conductor of thermal energy ('heat') as copper !!!
    • The 4th shared electron, can move freely through the layer so well. that graphene is a VERY good conductor of heat AND electricity like a metal (and better than copper!).
    • Again, this contrasts with diamond which is an electrical insulator and a poor thermal conductor.
    • The electric charge OR thermal kinetic energy is quickly conveyed by the delocalised electrons.

  • More notes on graphene, properties and uses

A quick comparison of carbon in the form of diamond, graphite and graphene

Compare and contrast their physical properties - its all about the nature of the bonding situation!

	Diamond	Graphite	Graphene
Structure ball and stick 2D/3D models	3D giant covalent network of C-C bonds	2D giant covalent network of C-C bonds	2D giant covalent network of C-C bonds
Bonding	Very strong 3D network of strong C-C covalent bonds. NO weak intermolecular forces between layers of atoms involved.	Very strong 2D network of C-C covalent bonds within layer Weak intermolecular forces between layers of atoms involved (dotted lines).	Very strong 2D network of C-C covalent bonds within layer. Its basically a single layer of graphite - but can be fabricated into a very strong material
Physical nature	Very hard, very strong structure	Brittle and slippery - weak structure because layers can slide over each other	Very strong across the layer
Melting point	Very high, great thermal stability	Very high, great thermal stability	Very high, great thermal stability
Electrical conduction	Very poor, no delocalised electrons	Moderately good conductor, delocalised electrons to carry current	Very good conductor, delocalised electrons to carry current - its as good as copper!
Thermal conductivity	Very poor, no delocalised electrons to carry current	Good conductor - delocalised electrons to convey kinetic energy quickly - more so than fixed atoms	Very good conductor, delocalised electrons to convey (thermal) kinetic energy
Solubility	Won't dissolve in any solvent - bonding to strong for solvent to break down structure	Won't dissolve in any solvent - bonding to strong for solvent to break down structure	Won't dissolve in any solvent - bonding to strong for solvent to break down structure
Example of use related to structure	Hard edge of a cutting tool.	Acts as a very thermally stable lubricant. Electrical contacts including electrodes in electrolysis processes.	Fabricated layers that need to disperse heat. Strong and used to strengthen other materials.

You can also compare the similarity of carbon (diamond), silicon Si (right diagram and silicon dioxide SiO₂ (left diagram) - the bonding diagrams say it all.

All three have a similar 3D giant covalent structure, like the one above or on the right, which could be carbon (diamond) or the element silicon.

They will therefore share a similar set of physical properties described above for diamond.