₁H

_ZSymbol, z = atomic or proton number

highlighting position of Group 4/14 elements

carbon

silicon

germanium

tin

lead

flerovium

3.5 (diamond)

The structure of the three allotropes of carbon (diamond, graphite and fullerenes), silicon and silicon dioxide (silica)

DIAGRAMS

• It is possible for many atoms to link up to form a giant covalent structure or lattice. The atoms are usually non–metals.
• This produces a very strong 3–dimensional covalent bond network or lattice.
• This gives them significantly different properties from the small simple covalent molecules mentioned above.

• In carbon dioxide, the smaller C atom can form a double bond with oxygen, but in silicon dioxide (silicon(IV) oxide, silica, quartz) the larger Si atom can only form single Si–O bonds.

• The result is SiO₂ has a giant 3D covalent lattice structure or network in which each silicon atom forms 4 bonds to an oxygen in a 'tetrahedral' spatial arrangement.

• The structure is therefore held together by strong covalent bonds (not weak intermolecular forces) and so it is far more thermally stable giving a high melting point, and insoluble in any solvent, because solvation energies are much lower than covalent bond energies.

• Carbon and silicon are two elements which form giant covalent structures i.e. they are high melting and insoluble solids.

• Carbon (diamond) and silicon form networks based on tetrahedral arrangements of C–C or Si–Si bonds around each atom. Both are very hard substances because of the strong bonding, diamond is harder because the smaller C atoms give shorter stronger bonds. Both are poor conductors of electricity because the outer electrons are strongly held and localised between the two atoms of any bond.

• However, carbon in the form of graphite, forms hexagonal ring layers in which the three C–C single bonds are supplemented by delocalised electron bonding from the 4th out electron of carbon. This makes graphite a moderately good electrical conductor as the electrons can move freely through a layer. The layers are held together by weak inter–molecular forces and easily slip over each other making graphite a 'slippery' brittle solid. But as a giant covalent structure it is still high melting and insoluble.

• Relatively recently (and another case of serendipity!) a 3rd form of carbon has been discovered in the form of the 'ball shaped' fullerenes.

• TYPICAL PROPERTIES of GIANT COVALENT STRUCTURES
• This type of giant covalent structure is thermally very stable and has a very high melting and boiling points because of the strong covalent bond network (3D or 2D in the case of graphite below).
• A relatively large amount of energy is needed to melt or boil giant covalent structures. Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.
• They are usually poor conductors of electricity because the electrons are not usually free to move as they can in metallic structures.
• Also because of the strength of the bonding in all directions in the structure, they are often very hard, strong and will not dissolve in solvents like water. The bonding network is too strong to allow the atoms to become surrounded by solvent molecules
• Silicon dioxide [silicon(IV) oxide, silica, SiO₂] has a similar 3D structure and properties to carbon (diamond) e.g. very hard, very high melting point and virtually insoluble in anything!
  • This contrasts sharply with the structure and properties of the gas carbon dioxide which is a small covalent molecule.
  • With only weak intermolecular forces between the O=C=O molecules it consequently has a very low melting/boiling point (actually it sublimes at –78^oC). Carbon dioxide readily dissolves in solvents such as water and organic polar solvents.
  • Carbon dioxide has two polar bonds, C^δ+=O^d–, but because of the linearity of the molecule the two permanent dipoles cancel out to give overall a non–polar molecule
  • Note that carbon + oxygen, instead of forming a 3D network of O–C–O single bonds, with the smaller carbon atom, it is energetically more favourable to form C=O double bonds and thus forming a small triatomic molecule.
• The hardness of diamond enables it to be used as the 'leading edge' on cutting tools.
• 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.
• Many naturally occurring minerals are based on –O–X–O– linked 3D structures where X is often silicon (Si) and aluminium (Al), three of the most abundant elements in the earth's crust.
  • Silicon dioxide is found as quartz in granite (igneous rock) and is the main component in sandstone – which is a sedimentary rock formed the compressed erosion products of igneous rocks.
  • Many some minerals that are hard wearing, rare and attractive when polished, hold great value as gemstones.

Carbon–DIAMOND and silicon

  SILICA

silicon dioxide

• Carbon also occurs in the form of graphite. The carbon atoms form joined hexagonal rings forming layers 1 atom thick.
• There are three strong covalent bonds per carbon (3 C–C bonds in a planar arrangement from 3 of its 4 outer electrons), BUT, the fourth outer electron is 'delocalised' or shared between the carbon atoms to form the equivalent of a 4th bond per carbon atom (this situation requires advanced level concepts to fully explain, and this bonding situation also occurs in fullerenes described below, and in aromatic compounds you deal with at advanced level).
• The layers are only held together by weak intermolecular forces shown by the dotted lines NOT by strong covalent bonds.
• 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 3D network)..
• Graphite will not dissolve in solvents because of the strong bonding
• BUT there are two crucial differences compared to diamond ...
  • Electrons, from the 'shared bond', can move freely through each layer, so graphite is a conductor like a metal (diamond is an electrical insulator and a poor heat conductor). Graphite is used in electrical contacts e.g. electrodes in electrolysis.
  • 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 is used as a lubricant.
• These two different characteristics 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 slide over each other.

GRAPHITE

• A 3rd form of carbon are fullerenes or 'bucky balls'! It consists of hexagonal rings like graphite and alternating pentagonal rings to allow curvature of the surface.
• Buckminster Fullerene C₆₀ is shown and the bonds form a pattern like a soccer ball. Others are oval shaped like a rugby ball. It is a black solid insoluble in water.
• They are NOT considered giant covalent structures and are classed as simple molecules. They do dissolve in organic solvents giving coloured solutions (e.g. deep red in petrol hydrocarbons, and although solid, their melting points are not that high.
• They are mentioned here to illustrate the different forms of carbon AND they can be made into continuous tubes to form very strong fibres of 'pipe like' molecules called 'nanotubes'. These 'molecular size' particles behave quite differently to a bulk carbon material like graphite.
• Uses of Nanotubes:
  • They can be used as semiconductors in electrical circuits.
  • They act as a component of industrial catalysts for certain reactions whose economic efficiency is of great importance (time = money in business!).
    • The catalyst can be attached to the nanotubes which have a huge surface are per mass of catalyst 'bed'.
    • They large surface combined with the catalyst ensure two rates of reaction factors work in harmony to increase the speed of the industrial reaction.
  • Nanotube fibres are very strong and so they are used in 'composite materials' e.g. reinforcing graphite in carbon fibre tennis rackets.
  • Nanotubes can 'cage' other molecules and can be used as a means of delivering drugs in controlled way to the body.

FULLERENES