GIANT COVALENT STRUCTURES
CHEMICAL BONDING Part 4 Covalent Bonding
– giant covalent structures and polymers
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)
Ionic Bonding – compounds and properties
Covalent Bonding – small simple molecules and properties
Covalent Bonding – macromolecules and giant covalent structures
Metallic Bonding – structure and properties of metals
Part 6 More advanced concepts for
advanced level chemistry (in preparation, BUT a lot on
intermolecular forces in Equilibria Part 8)
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 *
* properties of polymers
(diamond), carbon (graphite), carbon
(buckminsterfullerene/fullerenes), silica/silicon dioxide
BIG!4. 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
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
Part of the modern Periodic Table
Pd = period,
Gp = group
metals => non–metals
that H does not readily fit into any group
Chemical Symbol eg 4Be
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
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
- The structures of giant covalent
structure are usually based on non–metal atoms like carbon, silicon
- 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
- 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
- This is because it takes so much
thermal kinetic energy to weaken the bonds sufficiently to allow
- This gives
them significantly different properties from the small simple
covalent molecules (see simple molecular
- 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,
and fullerenes (below)
are the three solid allotropes of the
- 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
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
- 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
- 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.
- 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.
where n is an extremely large number of carbon atoms!
- 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.
- 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
- 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
- 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.
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
each silicon atom (black blobs) forms four strong covalent bonds
with the linking oxygen atoms (yellow blobs).
- Again like diamond, theoretically
all the atoms in a silica crystal are linked together by a strong 3D
covalent bond network.
- It takes a lot of energy to break (overcome)
the strong silicon-oxygen bonds in the giant covalent lattice of
silicon dioxide (silica).
- 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
a 3D diagram
a 3D diagram
multilayers of Cn sheets where n is an extremely
large number of carbon atoms all joined together!
- Carbon also occurs in the form of
- 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.
- 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 AND is free to move around - hence graphite's
ability to conduct electricity.
- This situation requires advanced level concepts to fully explain
the structure of graphite,
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
- 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.
- 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
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.
a 3D section of a graphite crystal
Graphene is a single sheet of
graphite, therefore it is another form of the non-metallic
Graphene is a 2D nanomaterial
because it is only one carbon atom thick.
It 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
- 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
- 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
- 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
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
- 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
- Carbon nanotubes have a very high tensile
strength, very good electrical conductivity and a relatively high
thermal conductivity - good conductors of electricity and heat.
- 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
- 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.
- 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 because the thin carbon nanotubes can
penetrate cell walls.
- More on
- I've written NEW pages with more
examples and more details on
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
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.
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)
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!
More on polymers in Oil
notes on polymers and Extra
organic chemistry notes on polymers.
the formation of poly(ethene)
a section of poly(chloroethene), PVC
All these examples are 2D representations of the molecular
structure of the polymer-plastic material
A couple of Advanced Level 'scribbles',
yet to be typed up!
keywords–phrases formulae: giant covalent lattice structures
carbon diamond graphite SiO2 explaining their physical properties
WHAT NEXT? and other associated Pages