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CHEMICAL BONDING Part 4 Covalent Bonding – giant covalent structures and polymers

Doc Brown's Chemistry Chemical Bonding GCSE/IGCSE/O/AS/A Level  Revision Notes

DIAGRAMS of GIANT COVALENT STRUCTURES and their PROPERTIES EXPLAINED – This section describes how covalent bonds can lead to large linear ('1D') giant molecular covalent structures e.g. thermoplastic polymer macromolecules, two dimensional ('2D') structures like graphite layers and three dimensional ('3D') giant covalent structured molecules like diamond, silica and thermosetting plastics. The physical properties of diamond, graphite, fullerenes,  silica (silicon dioxide) are described and explained using models of their molecular structure.  Examples of the uses diamond, graphite, fullerenes are explained. These notes on giant covalent structures are designed to meet the highest standards of knowledge and understanding required for students/pupils doing GCSE chemistry, IGCSE chemistry, O Level chemistry, KS4 science courses and a basic primer for AS/A Level chemistry courses. These revision notes on giant covalent structures should prove useful for the new AQA, Edexcel and OCR GCSE (9–1) chemistry science courses.


Part 1 Introduction – why do atoms bond together? (I suggest you read 1st)

Part 2 Ionic Bonding – compounds and properties

Part 3 Covalent Bonding – small simple molecules and properties

Part 4 Covalent Bonding – macromolecules and giant covalent structures (this page)

Part 5 Metallic Bonding – structure and properties of metals

Full CHEMICAL BONDING INDEX including Part 6 advanced concepts for advanced level chemistry

Part 4. COVALENT BONDING – macromolecules & giant covalent structures

 giant network bonding – giant molecules e.g. carbon C–diamond/graphite, silicon Si/silica SiO2

properties of giant covalent structures * polymers/plastics * properties of polymers

carbon (diamond), carbon (graphite), carbon (buckminsterfullerene/fullerenes), silica/silicon dioxide SiO2

 BIG!(c) doc b4. Large Covalent Molecules and their Properties

Macromolecules – giant covalent networks and polymers. What is the bonding, structure and properties of the carbon allotropes diamond, graphite & buckminsterfullerenes (fullerenes)?, why does diamond have such a high melting point? why is silica (silicon dioxide) a giant covalent structure, thermosets, thermoplastics? Because covalent bonds act in a particular direction i.e. along the 'line' between the two nuclei of the atoms bonded together in an individual bond, strong structures can be formed, especially if the covalent bonds are arranged in a strong three dimensional giant covalent lattice.

Its a good idea to have some idea of where the elements forming giant covalent structures are in the periodic table

The black zig–zag line 'roughly' divides the metals on the left from the non–metals on the right of the elements of the Periodic Table.

Pd metals Part of the modern Periodic Table

Pd = period, Gp = group

metals => non–metals
Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0

1H  Note that H does not readily fit into any group

2 3Li 4Be atomic number Chemical Symbol eg 4Be 5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba Transition Metals 81Tl 82Pb 83Bi 84Po 85At 86Rn
Gp 1 Alkali Metals  Gp 2 Alkaline Earth Metals  Gp 7 Halogens  Gp 0 Noble Gases

Chemical bonding comments about the selected elements highlighted in white

The non–metallic elements carbon and silicon form giant covalent structures

Materials that consist of giant covalent structures are solids with very high melting points and usually physically hard materials (not graphite). All of the atoms in these structures are linked to other atoms by strong covalent bonds in specific directions eg a grain of sand (silica) is one giant molecule! These substances usually have an extended 3D network of strong covalent bonds. These bonds must be overcome to melt or boil these giant covalent substances and this requires very high temperatures to give the particles sufficient kinetic energy to weaken the bonds and cause melting of the substance. Diamond and graphite (forms of carbon) and silicon dioxide (silica) are examples of giant covalent structures. You should be able to recognise giant covalent structures from diagrams showing their bonding and structure. You also need to be able to explain the properties of giant covalent substances in terms of their molecular structure. Most giant covalent structures don't have freely moving charged particles like ions or electrons, so they are poor conductors of electricity.

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


  • It is possible for many atoms to link up to form a giant covalent structure or lattice.
    • The structures of giant covalent structure are usually based on non–metal atoms like carbon, silicon and boron.
    • The atoms in a giant covalent lattice are held together by strong directional covalent bonds and every atoms is connected to at least 2, 3 or 4 atoms.
    • What you might call 'atomic networking'!
  • This very strong 3–dimensional covalent bond network or lattice gives the structure great thermal stability e.g. very high melting point and often great physical strength.
    • This is because it takes so much thermal kinetic energy to weaken the bonds sufficiently to allow melting.
  • This gives them significantly different properties from the small simple covalent molecules (see simple molecular substances).
  • This is illustrated by carbon in the form of diamond (an allotrope of carbon). Carbon has four outer electrons that form four single bonds, so each carbon bonds to four others by electron pairing/sharing.
    • Pure silicon, another element in Group 4, has a similar structure.
    • NOTE: Allotropes are different forms of the same element in the same physical state. They occur due to different bonding arrangements and so diamond, graphite (below) and fullerenes (below) are the three solid allotropes of the element carbon.
      • Oxygen (dioxygen), O2, and ozone (trioxygen), O3, are the two small gaseous allotrope molecules of the element oxygen.
      • Sulphur has three solid allotropes, two different crystalline forms based on small S8 molecules called rhombic and monoclinic sulphur and a 3rd form of long chain ( –S–S–S– etc.) molecules called plastic sulphur.
  • 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 because strong chemical bonds must be broken (and not just weakening intermolecular forces as in the case of small covalent molecules like water).
    • 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 are in metallic structures (and they are NOT made up of ions).
    • All the valency bonding electrons are tightly held and shared by the two atoms of any bond, so in giant covalent structures they are rarely free to move through the lattice and not even when molten either, since these giant molecular covalent structures do NOT contain ions.
  • 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 (silica, SiO2) has a similar 3D structure and properties to carbon (diamond) shown below and also pure silicon itself.
  • DIAMOND Cn where n is an extremely large number of carbon atoms!
    • Diamond is an allotrope of carbon - a distinct molecular form of solid carbon
    • In diamond every carbon atom is strongly linked to four other carbon atoms by strong directional 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.
      • The result is a very pure crystal structure with a high refractive index that gives diamonds quite a sparkle as light passes through it.
    • The 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.
    • 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.
      • The more energy needed, the higher the melting point.
      • Under very high pressure, diamond will melt at a staggering 4700oC!
    • The strong bond network in diamond (and graphite and silica) prevents these materials from dissolving in any conventional solvent.
    • 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.
    • Pure elemental silicon (not the oxide) has the same molecular structure as diamond and similar properties, though not as strong or high melting.
      • The molecular diagram is the same for Sin, where n is a huge number!
      • Silicon melts at 1400oC, and has poor electrical conductivity and won't dissolve in any solvent.
      • The silicon in the transistors of electronic devices is 'doped' with other elements to increase its electrical conductivity.
  • SILICON DIOXIDE (SILICA) (SiO2)n where n is an extremely large number of silicon and oxygen atoms!
    • 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 ('silica') 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.
      • Looking at the diagram on the right of silica, each silicon atom (black blobs) forms four strong covalent bonds with the linking oxygen atoms (yellow blobs).
      • You can also see from the diagram that there are two oxygen atoms to every silicon atoms giving the empirical formula SiO2.
      • Again like diamond, theoretically all the atoms in a silica crystal are linked together by a strong 3D covalent bond network giving a strong rigid structure.
      • It takes a lot of energy to break (overcome) the strong silicon-oxygen bonds in the giant covalent lattice of silicon dioxide (silica), hence the high melting point of 1600oC.
    • Therefore Silica (SiO2) is a very hard substance with a very high melting point and won't dissolve in any solvent.
    • There are no free electrons so silicon dioxide doesn't conduct electricity.
    • Many more minerals that are hard wearing, rare and attractive when polished, hold great value as gemstones, but sand is also mainly silica, but not quite as valuable!























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a 3D diagram





silicon dioxide

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a 3D diagram

  • GRAPHITE consists of multilayers of Cn sheets where n is an extremely large number of carbon atoms all joined together!
    • Graphite is another form of carbon (another 'allotrope' of carbon)
      • The carbon atoms form joined hexagonal rings forming layers 1 atom thick in graphite.
      • 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' or 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 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 shown by the dotted lines NOT by strong covalent bonds, so graphite, for a giant covalent structure, is unusually weak physically.
      • 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 in the 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 ...
      • 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).
        • For a non-metal, graphite is a relatively good conductor of heat and electricity, which gives it some similarity with metals.
        • 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.
        • 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.
        • 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.
        • Graphite can act as a lubricant for the same reason, the slipperiness of the layers!
    • 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.


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2D layers

a 3D section of a graphite crystal

Graphene consists of a single layer of graphite and has properties that make it useful in electronics and composites. You should be able to explain the properties of graphene in terms of its structure and bonding. Know that fullerenes are molecules of carbon atoms with hollow shapes. The structure of fullerenes is often based on hexagonal rings of carbon atoms but fullerenes may also contain rings with five or seven carbon atoms. The first fullerene to be discovered was Buckminsterfullerene (a C60 molecule of hexagonal and pentagonal rings) which has a spherical shape. Carbon nanotubes are long cylindrical fullerenes with very high length to diameter ratios. Again, their properties make them useful for nanotechnology, electronics and materials. You should be able to recognise graphene and fullerenes from diagrams and descriptions of their bonding and structure and give examples of the uses of fullerenes, including carbon nanotubes.
    • Graphene is a single sheet of graphite, therefore it is another form of the non-metallic element carbon.

    • Graphene is a 2D nanomaterial because it is only one carbon atom thick made up of hexagonal rings of carbon atoms.

    • Like graphite, graphene conducts electricity because delocalised electrons can run through the layer.

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

    • More notes on graphene, properties and uses

  • A 3rd form of carbon (another allotrope of carbon) are fullerenes or 'bucky balls'!
    • They consists of hexagonal rings like graphite and alternating pentagonal rings of carbon atoms to allow curvature of the spherical surface, in fact curved sufficiently to form 'football' or 'rugby ball' shapes.
    • Some fullerenes have rings of seven carbon atoms, again to allow curvature of the surface of the hollow sphere of the 'bucky ball'..
  • Buckminster Fullerene C60 ,the first to be discovered, is shown on the right 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.
    • All of the fullerenes are hollow with the rings of carbon atoms forming the surface.
    • These 'molecular size' fullerene particles behave quite differently to a bulk carbon materials like graphite or diamond.
  • Fullerenes are NOT considered giant covalent structures and are classed as simple molecules.
  • Fullerenes 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.
  • Fullerene molecules can be used for drug delivery into the body, as lubricants, as catalysts and in the form of carbon nanotubes can be used for reinforcing composite materials, eg sports equipment like tennis rackets
  • Fullerenes 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'.
    • Carbon nanotubes are basically long cylindrical fullerenes.
    • Carbon nanotubes have a very high tensile strength, very good electrical conductivity and a relatively high thermal conductivity - good conductors of electricity and heat.
    • They are used as a component in strong composite materials.
  • Uses of carbon nanotubes – carbon nanotechnology – examples of nanochemistry
    • 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 an industrial reaction so making the process more efficient and more economic.
    • Carbon nanotube fibres are very strong and so they are used in 'composite materials' e.g. reinforcing graphite in carbon fibre tennis rackets.
      • This is partly due to carbon nanotubes have a high length to diameter ratio.
    • Carbon nanotubes or long fullerenes can 'cage' other molecules and can be used as a means of delivering drugs in controlled way to the body because the thin carbon nanotubes can penetrate cell walls.
    • Carbon nanotubes or fullerenes are an important additive in other oil based lubricants to enhance their performance.
      • Additives are added to lubricating oils to improve their effectiveness in reducing friction and as a chemical stabiliser eg to inhibit thermal degradation of the oil in high temperature situation, but I'm not sure what the function of carbon nanotubes is in this case? I suspect the reasons involve some complex physics of viscosity well beyond the scope of these notes!
    • More on carbon nanotubes, properties and uses
  • I've written NEW pages with more examples and more details on


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(c) doc bBonding in polymers and 1–3 'dimension' concepts in macromolecules

  • The bonding in polymers or plastics is no different in principle to the examples described above, but there is quite a range of properties and the difference between simple covalent and giant covalent molecules can get a bit 'blurred'.

    • Bonds between atoms in molecules, e.g. C–C in polymer molecule chains are called intramolecular bonds and very strong.

    • The much weaker electrical attractions between individual molecules are called intermolecular forces.

  • In thermosoftening plastics like poly(ethene) the bonding is like ethane except there are lots of carbon atoms linked together to form long chains. They are moderately strong materials but tend to soften on heating and are not usually very soluble in solvents. The structure is basically a linear 1 dimensional strong bonding networks. The polymer molecules are held together by weak intermolecular forces and NOT strong chemical bonds. The long polymer molecules mean the intermolecular forces are appreciable but the material is flexible and softens on heating.

  • Graphite structure is a layered 2 dimensional strong bond network made of 2D layers of joined hexagonal rings of carbon atoms with weak inter–molecular forces between the layers. (more details on graphite)

  • Thermosetting plastic structures like melamine have a 3 dimensional cross–linked giant covalent structure network similar to diamond or silica in principle, but rather more complex and chaotic!

    • Because of the strong 3D covalent bond network they do not dissolve in any solvents and do not soften and melt on heating and are much stronger than thermoplastics.

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the formation of poly(ethene) doc b oil notes

The repeating unit shown in brackets for poly(ethene) is -(CH2-CH2)-  and n is a very large number, and can represent hundreds or thousands of repeating units in one molecule.


a section of poly(chloroethene), PVC

The repeating unit shown in brackets for poly(chloroethene) is -(CH2-CHCl)-

All these examples above are 2D representations of the molecular structure of the polymer-plastic material


Other representations of a poly(ethene) molecule, typically a long chain molecule formed from lots of repeating units joined together by strong carbon-carbon bonds.

The intermolecular forces between polymer molecules are bigger than those between small molecules like water and great enough to ensure plastic polymers like poly(ethene) and PVC are solid at room temperature. The greater the intermolecular forces the greater the energy needed to overcome them and melt a material. The higher the temperature, the greater the kinetic energy of particles to overcome the molecular attractive forces. These intermolecular forces are much weaker than ionic or covalent bonds so the melting points of polymers are still much less than those of giant covalent structures like diamond or silica or ionic compounds like sodium chloride. Also, giant covalent structures like diamond and silica have very strong 3D covalent bond networks making them much more thermally stable.

A couple of Advanced Level 'scribbles', yet to be typed up!




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