Structure and bonding of carbon

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  1. Different forms of carbon
  2. What is an allotrope?
  3. Types of carbon: diamond
  4. Types of carbon: graphite
  5. Types of carbon: Buckminster-Fullerene / carbon nanotubes
  6. Summary of carbon allotropes.
  7. Bonding, structure and properties using carbon allotropes.
Topic Notes

In this lesson, we will learn:

  • To recall the definition of allotrope and name some carbon allotropes.
  • How the bonding of a material leads to its specific 3d-structure.
  • How the 3d structure of a material explains its unique properties using carbon allotropes as an example.


  • The carbon atom can make four covalent bonds to other atoms to fill its outer electron shell. Partly because of this high number, it has multiple forms or allotropes when samples of carbon are found in nature. An allotrope is the unique bonding arrangement and structure that atoms of an element make.
  • There are three important examples of carbon allotropes: diamond, graphite, and Buckminster Fullerene (and carbon nanotubes).
    • Diamond is an allotrope of carbon where each carbon atom is covalently bonded to four other carbon atoms.
      • The four strong covalent bonds around each carbon atom makes a tetrahedral shape around each atom.
      • This bonding creates a large lattice structure because every carbon atom is connected to four others. Diamond is practically one giant molecule because every single carbon atom is (eventually) connected to all the other atoms in the lattice through this bonding. For this reason, we say diamond has a giant covalent structure.
      • This structure gives diamond its unique properties:
        • Diamond is clear and colorless, as the carbon atoms in this structure reflect visible light. This also makes it shiny and lustrous which is why it is desirable in jewelry.
        • Diamond is extremely hard, as deforming diamond would need you to deform the entire giant lattice that every 'diamond carbon' atom is part of. This makes it useful in cutting tools and drills.
        • Diamond is insoluble in water, as interactions with the water molecules are not nearly strong enough to 'pull apart' the giant covalent lattice, so it does not dissolve.
        • Diamond cannot conduct electricity, as no free electrons are found in the diamond lattice. All the electrons of all carbon atoms in the lattice are locked up in covalent bonds to each other, so no carrying of electric charge through the lattice can take place.
    • Graphite is an allotrope of carbon where each carbon atom is covalently bonded to three others in layers of 2d sheets.
      • The three strong covalent bonds on each carbon atom are equally spaced in 2d 120o120^{o} apart from each other. There is one electron on each carbon atom still unbonded or 'free'.
      • This bonding gives graphite a structure of layers of 2-dimensional carbon atom sheets. These sheets stack on top of each other with weak stabilising interactions due to the spare electron of each carbon atom.
      • This unique structure of graphite gives it its unique properties that are quite different from diamond:
        • Graphite is a dark grey/black colour and is opaque as it absorbs visible light that interacts with it.
        • Graphite is a smooth, slippery material because the stabilising forces between the sheets of carbon atoms are quite weak. This means applying some pressure to graphite makes the layers slide over each other quite easily. This is how pencils work: graphite layers slide off of the pencil and onto the paper we write on when we press the pencil down! It’s also used in lubricants, which are chemicals that deliberately reduce friction.
        • Graphite is a good conductor of electricity,because each carbon atom has a spare electron. The spare electrons of all the carbon atoms are delocalised - they are capable of moving and carrying electric charge throughout the sheet that it is part of and through weak interactions that hold the layers close together. For this reason, graphite is used in electrodes for electrolysis experiments.
    • Buckminster Fullerene is an allotrope of carbon where each carbon atom is bonded to other carbons to make a 3d spherical ball of 60 carbon atoms known as a 'buckyball'.
      • Buckminster Fullerene is one of a larger category called Fullerenes.
      • Some Fullerenes have tube-like structures which have a very large surface area to volume ratio. These fullerenes are called nanotubes and they have unique desirable properties such as conducting electricity and high strength combined with lightness. Many are also useful catalysts – the high surface area to volume ratio is a common property seen in nanomaterials.
  • Carbon is just one example of an element which has multiple allotropes. We have only looked at three here, but another is coal which is a very important fuel. The difference between coal, graphite, diamond and Buckminster-Fullerene is simply how the carbon atoms are bonded and arranged together, they are all 'types of carbon'. A summary of the allotropes and their features are below:




    Buckminster-Fullerene / carbon nanotubes


    Giant covalent tetrahedral structure

    Giant covalent 2d layered sheets

    3d hollow sphere / 3d hollow cylinders

    Melting point

    Very high

    Very high


    Conducts electricity



    Yes (nanotubes)


    Very high



    Cutting tools, jewelry

    Pencils, electrodes

    Catalysts, medical science.

  • The different carbon allotropes are a good example of the bonding β†’\to structure β†’\to properties link in material chemistry:
    • The nature of bonding in a substance leads to the material's structure.
      The bonding will tell us how the negatively-charged electrons are interacting with the atoms which contain the positively-charged nucleus. This could lead to strong electrostatic forces (large positive charges attracting large negative charges like in ionic or metallic bonding), a giant covalent lattice, or a simple covalent molecule with only weak forces keeping molecules close together.
      • The structure of a substance leads to explaining the properties that we observe of the substance:
        • Can the particles in the structure move freely and interact with electric charge? Charged particles will interact with other charged particles. If the structure has charged particles with free movement, they will be able to carry electric charge and therefore conduct electricity. If there are no charged particles or the particles are unable to move throughout the structure, this won't be possible.
          This might affect the hardness and strength of the material too – if particles can/will move when a force is applied, the general structure will change shape!
        • What forces of attraction are keeping the whole structure together? Recall the particle model; in any substance the solid state has particles packed together with 'low energy' because the current energy is not enough to overcome the attractive forces keeping the particles together.
          In a gas state, particles are far apart with 'high energy' as that high energy has overcome the attractive forces that were holding them together - that's why the particles are far apart and gases take up more volume than solids!
          If those forces of attraction are strong then a lot of energy will be needed to overcome them. These structures will have a high melting/boiling point.