Chapter 7 – Covalent Molecules,
Networks and Lattices
Term 2, Week 1,
Lesson 1
Three Different Forms of Carbon
• Charcoal, Diamond and Graphite all consist of
the non-metal carbon.
• They all have the chemical composition, C.
• Charcoal, Diamond and Graphite are different
physical forms of carbon.
• These are called allotropes of carbon.
• Their physical properties are very different.
Diamond
• Diamond was found in India thousands of year
ago.
• It is such a prized jewel because when cut and
polished it sparkles with brilliance.
• It is the hardest naturally occurring substance.
Charcoal
• Charcoal is used to fuel fires for many years.
• It is also used for drawing.
• Charcoal is a black powdery substance.
Graphite
• Graphite is used:
– In ‘lead’ pencils
– As an electrode in some torch cells (batteries),
which have a rod of graphite in the centre
– As a dry lubricant in place of grease in machinery
that runs at high temperatures
– As an additive to make rubber and plastic
materials more flexible
– As fibres in recently developed sports equipment.
Graphite cont…
• The properties that make graphite so useful
are its ability to conduct electricity in solid
state, its soft, greasy nature and high melting
temperature.
• It is a grey opaque solid.
Carbon Bonds
• A diamond consists of carbon atoms bonded very strongly to other
carbon atoms in a 3D lattice.
• Diamond is hard and crystalline and has a very high melting
temperature.
• Graphite is very hard in 2 dimensions but slippery and soft in the
other direction.
• It consists of layers of carbon atoms.
• Within each layer the carbon atoms are very strongly bonded to the
other 3 carbons atoms but between layers there is very weak
bonding.
• The structure of carbon is currently the subject of research.
• It appears that charcoal may be made up of fragments of
buckyballs, which are particles shaped roughly like a soccer ball and
made up of groups of 60 carbon atoms.
Compounds of Carbon
• Carbon readily forms many compounds with
other non-metals and many of these
compounds consist of very small molecules
which are gases at room temperature.
Carbon Dioxide
• Carbon dioxide consists of a carbon atom
bonded to 2 oxygen atoms.
• The CO2 molecule is small with very strong
bonds holding the carbon and oxygen atoms
together but very weak forces between the
molecules.
• This means that the molecules are easy to
separate and so dry ice sublimes (goes from
solid to gas with no liquid phase).
Carbon & Bonding
• The type of bonding between the carbon
atoms in diamond and graphite and between
carbon and oxygen in carbon dioxide is all the
same type.
• They are all covalently bonded, this means
that electrons are shared to produce full outer
shells.
Covalent Bonds
• Atoms in a covalent bond are atoms of nonmetallic elements.
• Some other covalently bonded substances,
such as diamond and graphite, exist as
covalent lattices, in which a huge number of
atoms are joined together.
• Most substances that exist as small covalent
molecules are liquids or gases at room
temperature.
Covalent Bonds etc…
• There are only weak forces of attraction
between the molecules.
• These forces are relatively easy to overcome
and so the melting or boiling temperature of
these substances tend to be low.
• All the substances that exist as covalent
lattices are solids at room temperature
because all the atoms in the lattice are held to
one another by strong covalent bonds.
Comparing Diamond, Graphite and
Charcoal
Bonding Models
• When a covalent molecular substance melts,
the molecules separate for each other yet the
atoms within the molecules remain bound to
one another.
• This means that the bonding model for
molecular substances must describe two kinds
of forces:
– The strong forces of attraction which holds atoms
together within molecules;
– The weak force of attractions between molecules.
Covalent Bonding within Molecules
• When non-metals combine electrons are
shared so that each atom has eight electrons
in its outer shell.
• In the case of hydrogen, two outer electrons is
a stable configuration.
• Molecules formed in this way are more stable
than the separate atoms.
Small Molecules with Single
Covalent Bonds
• Hydrogen (H2)
– The simplest molecule is that formed when
hydrogen atoms bond.
– The melting temperature of solid hydrogen is -259
degrees, indicating that the forces between
hydrogen molecules are extremely weak.
– In contrast, an extremely large amount of energy
is required to separate the two hydrogen atoms of
a H2 molecule.
Hydrogen cont…
– When two hydrogen atoms form a hydrogen
molecule, two electrons, one from each H, are
shared between the two atoms.
– These two electrons spend their time between the
two nuclei and are considered to be localised.
– This region is where the –ve charged electrons will
have the greatest attraction to both +ve charged
nuclei.
– The strong force of attraction involving a shared
pair of electrons is called a single covalent bond
Small Molecules cont – Chlorine
(Cl2)
• A chlorine atoms needs one more electron to
achieve a stable outer shell.
• The two atoms in a chlorine molecule are
joined by a single covalent bond.
• Of all of the outer shell electrons, only two are
shared.
• The other electrons are distributed around the
two atoms.
Chlorine (Cl2) cont…
• The electrons that are shared are known as the
bonding electrons.
• The outer shell electrons not shared are called the
non-bonding electrons.
• Pairs of non-bonding electrons are called lone pairs.
• This means that chlorine has one bonding pair and
three lone pairs.
• Chlorine needs to share one electron, so we say that
it has a covalency of 1.
• The covalency of an atom is generally the number of
electrons it shares when covalently bonded.
Chlorine (Cl2) cont…
• However, this cannot be assumed for all
situations.
• There are a number of different compounds in
which atoms do not achieve an octet of
electrons or exhibit different covalencies in
different molecules for various reasons.
• For example, sulfur forms a number of stable
oxides, SO2 and SO3, each with sulfur
exhibiting a different covalency.
Chlorine (Cl2) cont…
Charge Cloud Diagrams
• Are more realistic ‘pictures’ of hydrogen and
chlorine molecules.
• They attempt to represent the fact that the
electrons within the molecules are not
stationary but are continuously moving.
• This gives the appearance of negative charge
clouds in the molecules.
Charge Cloud Diagrams cont…
Small Molecules cont – Hydrogen
Chloride (HCl)
•
•
•
•
•
H electron configuration is 1
Needs one electron for stable shell
So forms one covalent bond by sharing 1 electron
Chlorine electron configuration is 2,8,7
So forms one covalent bond by sharing 1 electron, leaving six
non-bonding electrons (three non-bonding lone pairs)
• In a hydrogen chloride molecule, each atom contributes one
electron to form a shared pair of electrons.
• The hydrogen atom can be considered to have a stable outer
shell of two electrons and the chlorine a stable outer shell of
eight electrons.
Hydrogen Chloride (HCl)
A representation of a hydrogen
chloride molecule, showing the
outer-shell electrons.
A charge cloud
diagram of hydrogen
chloride.
Your Task
• Look over pages 115-118
• Come up with your own summaries for
– Molecules with more than two atoms
– Some simple hydrocarbons
– Some small molecules with double or triple bonds
– Representing molecules
Week 1, Lesson 2
Shapes of Molecules
• The shape of a molecule describes the way in
which the atoms are arranged.
• For a simple molecule such as hydrogen, or
hydrogen chloride the shape is linear.
• The two atoms must be in a straight line.
• For molecules with more than two atoms the
situation is not so simple.
Shapes of Molecules cont…
• Consider a molecule of methane.
• The central carbon is surrounded by four
separate pairs of electrons in its outer shell.
• Each pair represents a single covalent bond
between the carbon atom and a hydrogen
atom.
• These electrons pairs are negatively charged
and so repel each other.
Shapes of Molecules cont…
• Molecules are most stable when the electron
pairs, both bonding and non-bonding, are as
far apart as possible.
• The maximum separation possible for the four
pairs of valence electrons is described as
tetrahedral.
Tetrahedral Formation
Water’s Formation
• Looking at the valence structure for water, the oxygen
atom has four separate pairs of electrons around it.
• These four pairs adopt the same arrangement as
methane.
• But, when describing the shape of a water molecule, the
most important thing is the position of the atoms.
• A water molecule may be described as angular or vshaped.
• Although the lone pairs help to determine the shape of
the molecule, they are not included in the description.
• Shape describes the positions of the atoms only.
Ammonia
• Is a different shape from that of methane.
• This is because one point of the tetrahedron
contains a lone pair of electrons.
• The four separate electron pairs around the
nitrogen atom assume a tetrahedral
orientation, however, only three points of the
tetrahedron are occupied by hydrogen atoms.
• Therefore, the molecular shape is a triangular
pyramid.
Hydrogen Fluoride
• The orientation of the four pairs of electrons
(three non bonding lone pairs and one
bonding pair) around the fluorine atom is also
tetrahedral.
• Because there are only two atoms in the
molecule, the shape of the molecule is linear.
Molecule Shapes
VSPER Model
• The valence shell electron pair repulsion
model states that electron pairs in the outer
shell of an atom in a molecule will repel one
another due to their negative charge and
assume positions as far apart as possible,
while remaining attached to the atom.
VSPER Model cont…
• In all four molecules, CH4, NH3, H2O and HF, the
central atom has four pairs of electrons in its outer
shell.
• So in each case, the electron pair geometry is
tetrahedral which the molecular shape is determined
by the positions of the atoms.
• The VSPER model applies equally to molecules where
the central atom has either less or more than four
electron pairs around it.
• Some examples of simple molecules where the
central atom does not obey the octet rule are given
on the next slide.
Exceptions…
(a) The valence structure of BeH2 is linear because the central atom, Be, is
surrounded by two pairs of electrons only.
(b) BF3 has the shape of a planar triangle. B has three pairs of valence
electrons which are arranged at 120° from one another to achieve maximum
stability.
(c) In SF6, the six F atoms are arranged at 90° from one another, the
maximum distance possible for six electron pairs. Its shape is octahedral.
Carbon Dioxide (CO2)
• Carbon has an electron configuration of 2,4 so it needs four
electrons for a stable outer shell.
• Oxygen has an electron configuration of 2,6 and so need to
electrons for a stable outer shell.
• The central atom, carbon, forms a double bond with each
oxygen.
• The four electron pairs are concentrated in two regions of
negative charge only, each one consisting on a double bond.
• These two regions repel and so assume positions as far from
one another as possible.
• That is, on opposite sides of the carbon atom.
• See figure 7.20, page 121
Week 1, Lesson 3
Forces between Molecules
• If you rub a plastic or glass rod with a piece of
fur or nylon and hold the rod next to a thin
stream of water from a tap, you will see how
the water is attracted to the charged rod.
• This suggests that the water molecule has
some charge although the model of a
molecule does not show this.
• Since some molecules other than water also
behave as though they have some charge the
model will need to be modified.
Polarised Bonds
• When two different non-metals form a covalent
bond, one atom usually attracts the bonding
electrons more strongly.
• In HCl, the chlorine atom attracts the bonding
electrons more strongly than the hydrogen atom.
• Chlorine is said to have a greater electronegativity
than hydrogen.
• This means that the bonding electrons spend more
time closer to the chlorine atoms.
• The chlorine end of the molecule becomes slightly
negative and the hydrogen atom is slightly positive.
Polarised Bonds cont…
• The chlorine only has a partial charge, not a whole negative
charge, and hydrogen a partial positive charge.
• The covalent bond in the hydrogen chloride molecule is said
to be polarised and the molecule is a polar molecule.
• The molecule is said to be a dipole, because it has two
charged ends, or poles.
• In solid hydrogen chloride, the molecules will be arranged so
that the oppositely charged ends of the dipoles are next to
each other.
• It is these forces that hold them together.
• This type of attraction is called dipole-dipole attraction.
Polarised Bonds cont…
Intermolecular Forces and Melting
• Dipole-dipole attraction forces are strong enough to hold
hydrogen chloride molecules together in a solid lattice at low
temperatures.
• These forces are much weaker than the covalent bonds within
the molecule.
• When this sort of bond is heated to its melting temperature,
the molecules gain sufficient energy to overcome some of the
forces of attraction and the lattice breaks to form a liquid,
however, each molecule remains intact because of the strong
covalent bonds.
• The dipole-dipole forces have not totally disappeared, they
are still present and hold the molecules in a liquid state.
Identifying Polar Molecules
• Compare the structural formulas below…
Formaldehyde has the polar bonds arranged in
such a way that each end has an opposing
charge. The oxygen end has a negative charge,
while the hydrogen end has a positive charge.
The molecule is therefore a dipole.
The carbon dioxide, however is symmetrical. Both
ends have a negative charge, with the positive in
the centre. Both the negative ends are equal and
therefore is not a dipole.
Identifying Polar Molecules cont…
• The molecule tetrafluormethane (CF4) is also non-polar.
Although each of the covalent C-F bonds are polarised, the
molecule itself does not have a negative and a positive side.
• So two conditions must apply if a molecule is said to be
dipole:
– It must have polar bonds
– The partial charges must be distributed asymmetrically
across the molecule.
– Polar molecules will attract each other with the negative
side of one molecule attracting the positive side of
another. This weak dipole-dipole attraction helps hold the
molecules together in a polar covalent molecular solid.
Hydrogen Bonding
• Is the most significant kind of dipole-dipole
attraction.
• Hydrogen bonding occurs between molecule in
which hydrogen is bonded to N, O or F.
• It occurs because:
– N, O and F are highly electronegative atoms
– The bonds between hydrogen atoms and each of these
other atoms are highly polarised
– There is a relatively large partial positive charge on the
hydrogen atom when it is bonded to any of these atoms
– N, O and F atoms each have at least one lone pair of
electrons when they form molecules.
Hydrogen Bonding cont…
• The H-bonding between the partially positive hydrogen atom
of one molecule and the lone pair on a N, O or F atom of a
neighbouring molecule is much stronger than other cases of
dipole-dipole bonding.
• Hydrogen bonding in not another type of bonding such as
ionic, metallic and covalent bonding.
• Hydrogen bonding is an intermolecular force and only occurs
from one covalent molecule to another under the conditions
mentioned.
• Hydrogen bonds are much weaker than ionic, metallic or
covalent bonds.
• They are more significant then other forms of dipole-dipole
attraction.
Hydrogen Bonding cont…
Weakest Bonding Forces
• The weakest bonding force operates in all
substances.
• It is most important when it is the only force
between particles.
• It arises as a result of the constant movement of
electrons in atoms and molecules.
• At one instant, it is very unlikely that all electrons in
the atom or molecule are distributed evenly around
the nucleus. It’s more likely that at one point there
are more electrons at one ‘side’ than the other.
Weakest Bonding Forces cont…
• This causes the atom or molecule to be
momentarily polar and so electrostatically
interact with neighbouring atoms or
molecules.
• Because the electrons are moving so rapidly,
the direction of these instantaneous
interactions are constantly changing.
• These instantaneous dipoles results in a net,
but weak force of attraction between all
particles.
Weakest Bonding Forces & Noble
Gases
• The effect of this force can be seen if you examine the melting
and boiling temperature of the noble gases.
• Bonding models are based on the observations that the atoms
of noble gases rarely combine with other atoms.
• They rarely gain, lose, or share electrons by reacting with
other substances.
• As elements, the particles present in the solid, liquid and gas
states of noble gases are the individual atoms.
• The fact that liquid and solid states of these elements can be
formed means that there must be forces between the atoms.
• The very low values for their melting and boiling
temperatures indicate that these forces are extremely weak.
Weakest Bonding Forces & Noble
Gases cont…
• The melting and boiling temperatures of the noble gases
increase as the sizes of the atoms increase.
• Even though these forces are weak they become more
significant as the number of electrons, and size of the atom,
increases.
• These forces are only considered when no significantly
stronger forces exist.
• They are known as dispersion forces or van der Waal’s forces
and are responsible for holding the molecules together in
non-polar molecular substances.
• Examples of these substances include molecular oxygen,
hydrogen, nitrogen and methane, as well as the noble gases.
Weakest Bonding Forces
(a) How an instantaneous dipole forms in a chlorine molecule.
(b) How an instantaneous dipole in one xenon atom interacts with a
neighbouring xenon atom.
Week 2, Lesson 1
Covalent Lattices
• Diamond and graphite are giant covalent lattices.
• Carbon dioxide is a small covalent molecule.
• Carbon dioxide has strong bonds between the
carbon and two oxygen atoms. However, the bond
between molecules is weak. This is why it is a gas at
room temperature.
• Diamond and graphite form covalent bonds
continuously throughout the lattice.
• The structure of diamond is known as a network
lattice.
• The structure of graphite is described as a layer
lattice.
Covalent Network Lattices
• Diamonds are the hardest naturally occurring
substance and is also used widely in machinery for
drilling, cutting and polishing.
• In diamond, each carbon atom is covalently bonded
to four other carbons, and the bonding extends
through the lattice.
• Four bonds are directed away from each carbon
atom at the angle of a regular tetrahedron.
• The strong covalent bonds are continuous
throughout the lattice.
• There are no weak links.
• See page 129, Figure 7.32 for structure of diamond
Graphite: A Covalent Layer Lattice
• Graphite is very hard in one direction, but
quite slippery and soft in another direction.
• It has a layered structure.
• The carbon atoms within the layers are held
together by covalent bonds.
• These bonds within the layer are very strong.
• The forces between layers are weak dispersion
forces.
• This structure is known as a covalent layer
lattice.
Covalent Layer Lattice cont…
• Within a layer, each carbon atom is bonded to
three other carbon atoms.
• The 4th electron from each carbon is
delocalised, it is free to move within the layer.
• These electrons allow graphite to conduct
electricity.
• In graphite, the distance between layers is
much greater than the carbon to carbon
length within the layer.
• See page 129, Figure 7.33 for structure of graphite.
Using the Models of Bonding
• Covalent molecular, network and layer lattice
models of bonding can be used to explain the
similarities and differences in properties of
non-metal elements and compounds.
Melting and Boiling Temperatures
• When a substance melts the particles break
out of their positions in the lattice and move
randomly in the liquid state.
• For substances composed of small molecules,
the forces between the molecules are weak.
• Only a small amount of energy is required to
disrupt the intermolecular forces.
• Therefore, substances with small molecules
have a low melting temperature.
Melting and Boiling Temperatures
cont…
• In network and layer lattices, a large amount
of energy is needed to disrupt the lattice.
• They have strong covalently bonding
throughout the lattice and have high melting
temperatures.
• At high temperatures, all bonds break and
atoms are free to move independently.
Electrical Conductivity
• An electric current will run through a substance
when there are charged particles are free to move.
• Covalent molecules have no overall charge.
• Because the molecules are neutral and there are no
delocalised electrons, covalent molecular substances
cannot conduct electricity in solid or molten state.
• In network lattices, all electrons are localised, as such
there are no free moving electron to carry an
electrical current.
• In graphite, one electron from each carbon atom is
delocalised and is free to move. This allows
electricity to be conducted.
Hardness and Softness
• Network lattices are usually very hard, eg, diamond. They
have strong covalent bonds throughout the lattice with all
atoms being a fixed positions.
• In graphite, the forces between layers are weak and so layers
can slide over one another. Graphite appears soft and greasy
in one direction.
• In molecular solid, the molecules adopt fixed positions and
form a regular lattice. The lattice is crystalline and can be
hard. For example, ice where the water molecules are held in
rigid positions to maximise the possibility of hydrogen
bonding.
• However, because of the intermolecular forces of attraction
being weak, covalent molecular substances are often liquid or
gases at room temperature.
Chemical Reactivity
• It is very hard to disrupt a strong bond.
• Substances composed of covalent network lattices like
diamond or silica do not react readily.
• Graphite, a covalent layer lattice, is also very unreactive.
• Small covalent molecules are much more reactive.
• This reactivity depends on the strength and stability of the
bonds. For example, triple bonded nitrogen is not very
reactive, where a double bonded oxygen is more reactive.
• The reactivity of small covalent molecules also depends on a
number of factors including polarity of bonds and nature of
the arrangement of the atoms in the molecule.