Covalent Bonding

Chapter Four - How atoms combine: Covalent bonding and
Intermolecular bonding forces
Chapter Story - Fullerenes (Bucky balls)
Until quite recently there were considered to be only two allotropes (different
structural forms of the same element) of carbon: graphite and diamond. In the late
1970’s a new group of carbon compounds was manufactured by the Australian
scientist Dr Bill Burch at the Australian National University in Canberra, though he
did not go on to investigate their structure. In 1985, Dr Richard Smalley and his
colleagues were the first to determine the structure of these new compounds - a
roughly circular group of carbon atoms arranged in a series of pentagons and
hexagons in a pattern similar to that seen in the leather swatches which make up a
soccer ball.
[Insert graphic of a standard fullerene, ensuring that alternate hexagons and pentagons
are clearly visible. Place diagram adjacent to a similarly sized photograph of a soccer
Smalley was to name these most unusual compounds ‘fullerenes’ though they were to
quickly become known as ‘Bucky balls’ after the inventor of geodesic domes, a
gentleman by the name Buckminster Fuller. Close inspection of the bonding in
fullerenes indicates that each carbon atom is bonded to only three other atoms,
resulting in the presence of one free, or delocalised, electron per carbon atom. This
characteristic enables fullerenes to conduct electricity and also confers on them a
number of other unusual and useful properties. Whilst initially little more than a
scientific curiosity, fullerenes have swiftly developed roles in a wide range of
potential applications involving superconductivity, the production of microscale
semiconductors and electromechanical devices and even broad-spectrum lasers.
The most stable fullerene involves 60 carbon atoms bonded into an approximately
spherical shape and is known as ‘buckminster fullerene C60’. Since this first structure
has been isolated and described a range of other fullerenes have been manufactured,
including those shaped into tubes, hemispheres and even one shaped like a saddle! It
is very likely that Buckyballs will play an important role in developing technologies
into the next century and beyond.
4.1 - Introduction
In Chapter Three the important concept of the chemical bond was introduced - the
means by which atoms come together to form more chemically stable structures such
as molecules and ionic crystals. We recognised that all bonding results from
electrostatic interactions (the attractive forces which act between particles of opposite
charge and the repulsive ones between particles of like charge). Whilst the donation of
electrons between metals and nonmetals causes the bonding type broadly classified as
“ionic bonding” the sharing of electrons between nonmetal atoms is classed as
“covalent” bonding and it is to this type of bonding that we shall address our energies
in this chapter.
4.2 - The Molecule
A molecule may be formally defined as “a discrete group of nonmetal atoms
covalently bonded to each other”. Of particular significance here is the fact that
molecules contain specific numbers of atoms in a set ratio. For example, a molecule of
water contains two atoms of hydrogen bonded to one atom of oxygen; it is of no
consequence whether the water is present in the solid, liquid or gaseous phase - the
molecule is always the same. Another important molecule is ‘glucose’ with the
chemical formula C6H12O6. Each glucose molecule contains a total of 24 atoms
chemically bonded to one another: 6 carbon, 12 hydrogen and 6 oxygen atoms.
Why then cannot we classify sodium chloride as a molecule, for we write its formula
as NaCl? The answer lies in the way the atoms (or ions in this case) arrange
themselves. Sodium chloride forms an ionic crystal lattice; an ‘infinite’ array of
sodium ions and chloride ions arranged in a cubic lattice in such a way that each
sodium ion is surrounded by 6 chloride ions and vice versa. The formula ‘NaCl’ is
simply the simplest whole-number ratio of Na+ to Cl- ions within the crystal lattice.
Similarly, it is not valid to describe composite materials such as wood, paper or
natural fibres as molecules; the atoms which make up these substances are not present
in discrete, countable units but rather as complex mixtures of many different
The ways in which these separate molecules interact with each other is most important
and, in part, depends on the shape and structure of the molecule. We shall look at
these intermolecular bonding forces in greater detail later in this chapter.
4.3 - The Covalent bond
Nonmetal atoms tend to have high eletronegativities - that is to say, they have a strong
tendency to attract electrons in order to fulfil the octet rule and so gain a filled outer
shell of 8 electrons when bonding. When nonmetal atoms form bonds with each other
both atoms ‘want’ the outer shell electrons of the other atom - thus sharing of
electrons takes place. When a pair of electrons are shared between two atoms a
covalent chemical bond has been formed and the pair of electrons are known as a
bonding pair. Other pairs of outer shell electrons which do not actually take part in the
bond are known as nonbonding pairs.
Just as we were able to draw electron shell diagrams for ionically bonded structures,
so too can we construct similar diagrams for molecules. Note that only the outer shell
electrons are shown in these representations as it is (generally) only the outer shell
electrons which actually take part in bonding.
4.3.1 - Diatomic molecules
Diatomic molecules are those which consist of only two atoms covalently bonded to
one another. This group includes all of the elemental nonmetals, with the exception of
the Inert Gases of Group VIII as the elements of Group VIII already have 8 electrons
in the outer shell and so have very little tendency to form bonds with other atoms. It
should be noted that certain compounds have been produced in recent years which
involve Group VIII elements forming bonds with extremely reactive gases such as
fluorine, though the generation of theses compounds is beyond the scope of this
Elements such as hydrogen, fluorine and chlorine have seven electrons in the outer
shell and so will tend to share one further electron in order to obtain eight electrons.
This sharing of one pair of electrons results in the formation of a single covalent bond
between the atoms.
Example 4.1 - Hydrogen gas (H2)
Example 4.2 - Chlorine gas (Cl2)
Example 4.3 - Hydrogen chloride (HCl)
The nonmetal elements in Group VI of the Periodic Table, such as oxygen and sulfur,
have only six electrons in their outer shells and so must gain a share of two more
electrons in order to fulfil the octet rule. In the case of oxygen gas a diatomic
molecule can form between two atoms of oxygen in such a way that four electrons, or
two pairs, are found between them - a double bond has been formed.
Example 4.4 - Oxygen gas (O2)
What type of bond might we expect between two nitrogen atoms, considering that
nitrogen is to be found in Group V of the Periodic Table and so atoms of this element
have only five electrons in their outer shell?
Example 4.5 - Nitrogen gas (N2)
The triple covalent bond between nitrogen atoms is very strong and helps us to
recognise why nitrogen is such an inert, or unreactive, gas. It requires a great deal of
energy to overcome the triple bond between the atoms and so allow nitrogen to form
new bonds with other atoms.
4.3.2 - Polyatomic molecules
Polyatomic molecules are those which consist of more than two atoms covalently
bonded to one another. This group encompasses the great majority of molecules
known and includes significant ones such as water, carbon dioxide and methane.
When determining the ways in which these molecules may share bonding electrons we
must consider the outer shell electron configuration of each atom involved in the
Example 4.6 - Water (H2O)
Oxygen has six electrons in its outer shell and so must gain a share of two more to
satisfy the octet rule. Hydrogen has only one electron and will tend to share another
electron in order to gain a filled outer shell (in its case, only two electrons). Thus,
oxygen will form separate covalent bonds with two hydrogen atoms in order to satisfy
the electron sharing requirements of all three atoms.
Example 4.7 - Methane (CH4)
Carbon is in Group IV and so has four electrons in its outer shell. In order to gain a
share of eight electrons it must form a covalent bond with four hydrogen atoms, as
shown in Example 4.7.
The carbon dioxide molecule is slightly different to the above examples, as in this
case multiple bonds are formed. Carbon needs to share four more electrons and
oxygen two in order that both elements’ atoms comply with the octet rule. The way in
which the electrons are shared is shown in Example 4.8.
Example 4.8 - Carbon dioxide (CO2)
4.3.4. - Electron-dot formulae
As covalent bonding generally only involves the outer shell electrons it is often more
convenient to represent the sharing of electrons to form covalent bonds in electron-dot
formulae. Using this technique, only outer shell electrons are represented by dots or
crosses. As electrons tend to occur in pairs (orbitals) the dots (crosses) are also paired
wherever possible. As a covalent bond consists of a pair of bonding electrons between
the two atoms involved, a pair of dots (crosses) between the atoms is representative of
a bond whilst other pairs of dots (crosses) will represent nonbonding pairs of
electrons. When drawing these representations we need to remember that electrons
carry a negative charge and so pairs of electrons will tend to repel each other in threedimensional space. This fact is most important when we begin to look at the shapes of
molecules in the next section.
Exercise 4.1
Write electron-dot formulae for each of the examples listed above:
a) H2
b) HCl
c) H2O
d) O2
e) N2
a) H2
b) HCl
H Cl
f) NH3
c) H2O
d) O2
e) N2
4.4 - Valence structures and shapes of molecules
The most useful and widely employed representation of molecules is the valence
structure. In this representation each pair of electrons, bonding and nonbonding pairs
alike, is shown by a simple line. Of considerable significance, however, is that the
actual shape of the molecule is also shown. To enable us to achieve an accurate
indication we must recall that each pair of electrons will be repelled from the others as
far as possible in three dimensional space for the simple reason that all electrons carry
a negative charge. This is known as the “valence shell electron pair repulsion
hypothesis”, or the VSEPR theory, and was first put forward by Sidgewick and
In the simplest diatomic molecules there are only two atoms to consider and so the
molecule will invariably be linear. In the unusual case of three pairs of electrons being
present, as is the case in boron trifluoride, BF3, the electron pairs will repel in such a
manner as to form an equilateral triangle. Note that the shape of the molecule is
defined by the locations of the atoms that make up the molecule. Thus, whilst
nonbonding pairs of electrons can be important in determining the overall shape of a
molecule they are not actually considered as a part of the designated shape.
[Insert graphic of linear and trigonal planar molecular shapes using both charge-cloud
representations and valence structures. Show H2, HCl, O2, N2 and BF3]
The most common situation encountered in molecules is the existence of four pairs of
electrons, either bonding or nonbonding, surrounding each atom. Whilst it may be
tempting to assume that the four pairs of electrons will arrange themselves in a
square-planar structure, with each pair of electrons at 900 to the others this is not
actually the case. In fact, the most widely spaced arrangement of four pairs of
electrons in three-dimensional space is known as the tetrahedral arrangement, where
each atom is placed at the vertex of a regular triangular pyramid. The bond angle in
this arrangement is approximately 1090.
[Insert graphic of tetrahedral and triangular pyramidal molecular shapes using both
charge-cloud representations and valence structures. Show CH4 and NH3]
It is possible for the atoms of elements in Period 3 and beyond to form molecules
which do not comply with the octet rule insofar as they can form quite large molecules
where up to 12 electrons (rather than just eight) may be involved in bonding. Whilst a
detailed treatment of the reasons for the existence of such molecules as these being
possible is beyond the realms of this course, in essence it can occur because the third
shell can hold up to 18 electrons.
Five pairs of electrons, as is found in phosphorus pentafluoride, PF5, arrange
themselves into what is known as a trigonal bipyramid. This shape is effectively two
triangular pyramids joined together at the base. Six pairs of electrons generate a shape
known as an octahedron as the shape (a pair of square-based pyramids attached at the
base in this instance) has 8 sides altogether. An example of a molecule which takes up
this octahedral arrangement of atoms is sulfur hexafluoride, SF6.
[Insert graphic of trigonal bipyramidal and octahedral molecular shapes using both
charge-cloud representations and valence structures. Show PF5 and SF6]
The orientation that successively larger numbers of pairs of electrons will take in
space, and the names of the shapes generated in the process, is summarised in the
Table below:
Table 4.1 - Shapes of molecules
No. of pairs of
Name of shape
Triangular planar
4.5 - Intermolecular bonding forces
We have seen how the arrangement of the outer-shell, or valence, electrons of the
constituent atoms of a particular molecule can affect its shape. But what of the actual
bonding orbitals themselves? Does each atom have an even share of the electrons
which constitute these orbitals, or does one attract the electrons more strongly? If the
sharing is uneven, how does the resultant polarity of the bonds within the molecule
affect its general physical and chemical properties?
To answer these most important questions we must first revisit the concept of
electronegativity - defined as the ‘electron attracting power of an atom’; and recall
that electronegativity increases across a Period and up a Group. Thus, elements such
as fluorine, oxygen and nitrogen (all toward the top right-hand corner of the Periodic
Table) have high electronegativities whereas the metals tend to be quite low. As every
element has, by definition, a different electronegativity value it becomes obvious that
the only molecules which will contain bonding electrons which are shared completely
evenly will be those made up of atoms of only one element, such as the diatomic
molecules N2, O2, Cl2, H2 etc. The great majority of molecules are made up of atoms
of more than one element and so will contain polar bonds.
[Note that the Table of electrovalencies shown is identical to Table 2.1 used earlier in
this book.]
Table 4.2 - Electronegativities of selected elements
H 2.1
Li 1.0
Na 0.9
K 0.8
Rb 0.8
Cs 0.7
Fr 0.7
Be 1.5
Mg 1.2
Ca 1.0
Sr 1.0
Ba 0.9
Ra 0.9
B 2.0
Al 1.5
Ga 1.6
In 1.7
Tl 1.8
C 2.5
Si 1.8
Ge 1.8
Sn 1.8
Pb 1.8
N 3.0
P 2.1
As 2.0
Sb 1.9
Bi 1.9
O 3.5
S 2.5
Se 2.4
Te 2.1
Po 2.0
F 4.0
Cl 3.0
Br 2.8
I 2.5
At 2.2
The term dipole is used to describe the uneven charge distribution which results from
electrons spending more time on average in the vicinity of one particular atom than
another. If the molecules which make up a substance are themselves polar (that is to
say, they contain dipoles) then these molecules will align themselves in such a manner
as the positive end of one molecule will interact with the negative end of another.
[Insert graphic to indicate the nature of dipole-dipole interactions.]
Intermolecular bonding (bonding between molecules) is the general name given to
this important class of interactions. But why are they considered to be of such
significance? As an experimental science any theory that we may put forward needs to
have an observable and experimental foundation. What evidence do we have to
suggest that these intermolecular bonds actually exist?
The fact that all atomic and molecular species can be converted into the liquid, and in
the great majority of cases even the solid state indicates that there must be attractive
bonding forces acting between the atoms or molecules. Most molecular substances
exist as gases at room temperature and are only converted to liquid, and subsequently
solid, states at very low temperatures when the kinetic energy of the particles is very
low. This in turn tells us that these intermolecular bonding forces must be relatively
weak compared to the strong intramolecular covalent bonds within the molecule, as
they are relatively easily overcome by the kinetic energy of the molecules as the
temperature increases. It is important that we recognise that the melting and boiling
points of a substance give us valuable information about the strength of the bonding
forces acting between its constituent particles.
4.5.1 - The ‘pure’ covalent bond: Dispersion forces
Covalent bonding results from the sharing of electrons between two or more nonmetal
atoms. In the so-called ‘pure covalent bond’ this electron sharing on average between
the atoms is exactly even, a situation which can only occur between atoms of the same
electronegativity - the same element.
[Insert simple graphic displaying shared electron cloud between H2, Cl2 and O2
However, even diatomic molecules such as these can exist in the liquid state if the
temperature is sufficiently low (See Table 4.2).
Table 4.2 - Boiling points of selected diatomic molecules
Boiling Point 0C
Hydrogen - H2
Nitrogen - N2
Oxygen - O2
Fluorine - F2
Chlorine - Cl2
Bromine - Br2
To appreciate the source of the weak attractive forces acting between these molecules
we must consider the orientation of all electrons, both bonding and nonbonding, at a
particular instant rather than over a period of time. If we take the simplest of these
molecules, hydrogen, and consider the orientation of the two bonding electrons the
situation becomes clearer. Whilst it is possible that these fast-moving electrons will be
symmetrically oriented around the two hydrogen nuclei it is far more likely that, for an
instant of time, the two electrons will be found at one side of the molecule thus
creating an instantaneous dipole for that moment (See Figure 4.1).
Molecule of H2 at some instant.
Same molecule an instant later.
symmetrical (no dipole)
nonsymmetrical (instantaneous dipole
Once that instantaneous dipole has been generated it will influence the orientation of
the electrons within molecules close to it to generate a weak polar interaction. A
moment later, the electrons have attained a different orientation and a new set of
interacting dipoles will be generated. Over time, the orientation in three dimensional
space of the electrons of a specific molecule will average to produce no permanent
dipole. However, the weak interactions generated by billions of instantaneous dipoles
will have resulted in a weak overall attractive force. These weak bonding forces are
known collectively as dispersion forces (also known as ‘van der Waal’s forces’). The
larger the molecule, and so the greater the number of electrons present, the more
frequent and pronounced will be these dispersion forces. By referring to Table 4.2 it
can be noted that the boiling points of these elemental substances increases with
increasing atomic number.
4.5.2 - The ‘polar’ covalent bond: Dipole-dipole attractions
The great majority of molecules consist of two or more atoms of different elements
and, as a consequence, will contain polar bonds - they are said to possess ‘permanent
dipole moments’. For example, the hydrogen chloride molecule, HCl, is polar in
nature as the chlorine atom is more electronegative than the hydrogen and so the
bonding electrons will tend to be located more often nearer to it than the hydrogen.
[Insert graphic to indicate polarisation of bonding orbital toward Cl. Indicate + for H
and - for the chlorine atom].
The symbol  (delta) is used to indicate a partial charge of positive or negative on
each atom. In all cases the more electronegative atom will draw the electrons towards
it more strongly and so will attain a slight negative charge whilst the less
electronegative atom will lose its share of the electrons and so will carry a slight
positive charge.
It is also possible to have molecules which contain polar bonds but overall are
nonpolar - the dipole moments cancel each other out. An example of such a molecule
is methane, CH4. This molecule contains four polar C-H bonds but they cancel each
other out due to the symmetrical nature of the molecule.
[Insert graphic to indicate polarisation of bonding orbitals in methane. Indicate + for
the hydrogens and - for the carbon atom. Try to indicate the idea that the dipole
moments cancel in three dimensions.]
Chloromethane, CH3Cl is a similar molecule but in this case it is polar. Can you see
[Insert graphic to indicate polarisation of bonding orbitals in chloromethane. Indicate
+ for the hydrogens and - for the chlorine atom. Emphasise the asymmetry of the
molecule in three dimensions.]
Note that this molecule is not symmetrical in three dimensions and so the dipole
moments cannot cancel each other out. The significantly greater strength of the
intermolecular bonding forces between chloromethane molecules compared to
methane can be seen in their respective boiling points: 40 0C compared to -162 0C!
Another important example of a nonpolar molecule which nevertheless contains polar
bonds is carbon dioxide; known as ‘dry ice’ in the solid state as it sublimes directly
from solid to gas without passing through the liquid state. When cooled sufficiently to
overcome the kinetic energy of the molecules, the carbon dioxide molecules’ dipoles
will align themselves in such a way as to form dipole attractions between adjacent
[Insert graphic to indicate polarisation of bonding orbitals in carbon dioxide. Indicate
+ for the carbon and - for the oxygen atoms. Ensure the linear nature of the
molecule and the orientation of dipoles is clearly evident.]
Carbon dioxide sublimes at the moderately high temperature of -78 0C. It should be
noted that dispersion forces are always present between molecules, though if the
molecules are significantly polarised their effects may be relatively small.
4.5.3 - A special case of dipole attractions: Hydrogen bonds
The dipole-dipole attractions between hydrogen atoms and the three most
electronegative elements of fluorine, oxygen and nitrogen are particularly strong and,
as a consequence, have been given the special name of hydrogen bonds. [To help
remember which atoms combine in molecules to exhibit hydrogen bonding the
memory mnemoic H-FON may be usefully employed]. Water exhibits hydrogen
bonding between its molecules and it is these unusually strong intermolecular
attractions which enable us to explain some of the unusual properties of water, such as
its relatively high melting and boiling points and its low density when frozen.
We know from consideration of its valence structure that water is a ‘V-shaped’
molecule and so will contain permanent dipoles. The water molecules will orient
themselves in three dimensional space in such a manner as to maximise the dipole
attractions between them.
[Insert graphic to indicate polarisation of bonding orbitals in water. Indicate + for the
hydrogen and - for the oxygen atoms. Ensure the v-shape nature of the molecule and
the orientation of dipoles is clearly evident. Use dotted lines between molecules to
represent hydrogen bonds.]
When cooled sufficiently the water molecules take up what is known as an ‘open
lattice structure’ which explains why ice is less dense than water. The ability of ice to
float on water is most unusual as almost all other molecular substances are more dense
as the solid than as the liquid. The significant polarity of the water molecules also
underlies its excellence as a solvent of polar solutes, such as sodium chloride (Table
salt) and sugar. The properties of water are examined in significantly more detail in
Chapter 10.
Another significant molecule which exhibits hydrogen bonding is ammonia, NH3. Its
valence structure shows us that ammonia is trigonal pyramidal in shape and, as such,
the molecule is asymmetric and contains permanent dipoles. Ammonia has a relatively
high boiling point at -33 0C though this is much lower than water due to the lesser
electronegativity of nitrogen in comparison to oxygen; it is also much more volatile.
Ammonia is extremely soluble in water due to the polar nature of both substances.
[Insert graphic to indicate polarisation of bonding orbitals in ammonia. Indicate + for
the hydrogen and - for the nitrogen atoms. Ensure the trigonal pyramidal shape of the
molecule and the orientation of dipoles is clearly evident. Use dotted lines between
molecules to represent hydrogen bonds.]
4.6 - The ‘giant’ molecules: diamond and graphite
Diamond has always been highly prized for its exceptional brilliance and beauty when
cut and polished. Diamonds are formed deep within the earth as a result of extremes
of temperature and pressure acting on pieces of carbon, generally in the form of
charcoal. The result is the formation of dull, generally clear stones with a waxy sheen.
The largest rough diamond ever discovered was named the ‘Cullinan’, a stone
weighing 3,106 carats, or in the vicinity of 620 grams! Like most of the world’s
diamonds, this stone was discovered in the Kimberley region of South Africa. Before
any diamond can reveal its inner beauty it must first be cut and polished. The most
usual technique employed is the so-called ‘brilliant cut’ where the diamond is cut and
polished to produce a total of 58 facets. This style of diamond cutting was first
introduced in the 17th century.
[Insert photograph of one of the larger and more spectacular diamonds. Perhaps
include a small comment and/or photo from the Argyle diamond mines in the
Kimberley mountain ranges of NW Western Australia].
A molecule is defined as ‘a discrete (or countable) group of covalently bonded
nonmetal atoms’. There are a small group of covalently bonded substances that do not
meet this strict definition and yet are still generally classed as molecules, these being
the giant covalent lattice structures such as diamond, graphite, silicon dioxide and
silicon carbide. Like ionic solids, these substances form ‘infinite’ arrays of atoms
bonded to each other in a regular manner; it is not possible to count the actual number
of atoms involved and so the criterion of being a ‘discrete group’ has not been met.
Each of these giant molecular lattices involves the Group IV elements carbon and
silicon - each element being capable of forming four bonds with other atoms.
Diamond and graphite are both allotropes of carbon; different structural forms of the
same element. Their properties, however, differ markedly as a direct consequence of
the significantly different ways in which the carbon atoms are bonded to each other. In
the case of diamond, each carbon atom is surrounded by three other carbon atoms
bonded into a tetrahedral lattice structure. The bonding forces are immensely strong in
all three dimensions, resulting in diamond being an extremely hard substance of very
high sublimation point (approximately 3550 0C). At these extremely high
temperatures the bonds between the carbon atoms are overcome and they have so
much energy that the atoms move straight into the gaseous phase.
Diamond is the hardest known naturally occurring substance and for this reason very
small or flawed stones, collectively known as industrial diamonds, are commonly
employed on the tips of industrial cutting equipment, such as drills. Due to the fact
that all of the bonding electrons in the outer shell of the carbon atoms are engaged in
bonding orbitals diamond is an excellent insulator (as there are no free electrons
available to conduct electric current) and is chemically inert. It is interesting to note
that the phrase “Diamonds are forever” made famous by the James Bond movie of the
same name, has a solid basis in chemical fact!
[Insert graphic to show the structural formula of both diamond and graphite. Ensure
the tetrahedral lattice structure and the hexagonal ring structure with free electron
(respectively) is clearly indicated. Include appropriate photographs.]
Graphite displays many fundamentally different properties to diamond even though
they are both allotropes of the same element. It is a soft, greasy solid which is a good
conductor of electricity and makes an excellent lubricant. Some of the more common
uses of graphite are listed below:
 the ‘lead’ in greylead pencils
 as a dry lubricant to replace oil in certain industrial applications, such as the
‘dry sump’ used in some high-performance racing car engines
 as an inert electrode in the common dry cell (often incorrectly described as
a ‘battery’)
 as a chemical additive to rubber and certain plastics to make them more
 converted into a fibre and used to produce high-strength, light and flexible
composite materials employed in such items as tennis racquets and fishing
[Insert photo/s of dry cells, with at least one cut open to display graphite rod, pencils
and fishing rods etc]
If we look carefully at the structural diagram of graphite (above) it can be noted that
each carbon atom is bonded to only two others in a series of hexagons which make up
a layer. As each carbon atom has four valence electrons this leaves one electron
unattached for each carbon atom in the lattice. The delocalised electrons are able to
move relatively freely between the layers of strongly bonded carbon atoms, explaining
why graphite is such a good conductor of electricity. Furthermore, with only weak
dispersion forces holding the layers of carbon atoms together it is easy for one layer to
slide past another, explaining the lubricant properties of the material. Because the
bonding within layers is very strong, graphite has the very high melting point of
3730 0C and so can be readily used as a lubricant in the high temperatures generated
within engines and other moving mechanical structures.
Two other common examples of substances which display covalent lattice structures
are silicon dioxide (SiO2) and silicon carbide (SiC). Like diamond, both of these form
tetrahedral lattices and as a consequence share many of diamond’s properties. Silicon
carbide is used as an abrasive for industrial cutting tools and silicon dioxide is the
major component of sand, also commonly used as an abrasive agent.
[Insert graphic to show the structural formula of both silicon carbide and silicon
dioxide. Ensure the tetrahedral lattice structure in both is clearly indicated. Include
appropriate photographs to show uses; perhaps photo of beach sands and cutting
wheel on drill.]
Table 4.2 - Summary of properties of selected giant molecules
[Insert graphic to show
tetrahedral structure]
[Insert graphic to show layer
lattice structure]
1. Extremely hard
2. Very high sublimation point
3. Nonconductor of electricity
4. Chemically inert
1. Soft and greasy
2. Very high melting point
3. Conductor of electricity
4. Good lubricant
Very strong covalent bonding in
three dimensional lattice
structure results in hardness and
high sublimation point.
Nonconductor as no free
electrons, which also explains
lack of reactivity.
Strong bonding into layers with
weak dispersion forces between
these layers. Each carbon atom
has one free electron as only
three bonds formed between
adjacent carbon atoms; hence
good conductor of electricity.
4.7 - Summary/ Objectives
At the end of this chapter you should:
 note that covalent bonding occurs between nonmetal atoms and involves the
sharing of one or more electrons between these atoms to (generally) ensure eight
outer shell electrons are shared by each atom
 recall that molecules are formed as a result of covalent bonding
 recall that a molecule is defined as a discrete group of covalently bonded nonmetal
 be able to accurately employ electron dot, valence and structural formulae to
represent selected molecules; including water, ammonia, methane, carbon dioxide
and the diatomic molecules hydrogen, oxygen and nitrogen
 understand how multiple bonds may be formed to ensure the octet rule is met
 appreciate that some larger molecules do not conform to the octet rule insofar as
they may have more than four pairs of bonding electrons; phosphorus pentafluoride
and sulfur hexafluoride being examples
 be able to apply the VSEPR (valence shell electron pair repulsion) model to
determine the shapes of molecules and accurately represent these shapes in two
 recognise the significance of the nature of the intermolecular bonding forces which
may be of significance when determining such physical properties as melting and
boiling points
 appreciate that the symmetry of a molecule will be a major factor in determining
the intermolecular bonding forces in action
 recall that dispersion forces occur between all molecules, though they are the only
intermolecular bonds present in symmetrical molecules
 note that dispersion forces are the weakest of the intermolecular bonding forces,
though their strength increases with the size of the molecule and the number of
electrons present
 recognise that dipole-dipole attractions result from the interactions between polar
molecules and that these forces are stronger than dispersion forces, though they
exist in parallel with each other
 appreciate that a molecule may contain polar bonds but still be nonpolar overall if
the dipoles cancel each other out
 note that hydrogen bonding is a particularly strong form of dipole-dipole attraction
which occurs between molecules which contain hydrogen and one or more of the
atoms fluorine, oxygen or nitrogen (remembered as ‘H-FON’)
 recall the structural features and properties of the giant molecular lattice structures
of diamond and graphite
 be able to employ a knowledge of these different structures to explain the different
properties of diamond and graphite
4.8 - Chapter Four: Questions
By first drawing the electron dot formulae determine the number of bonding
and nonbonding electrons for each of the following molecules:
a) H2
b) HF
c) O2
d) CH2Cl2
e) SiCl4
f) SF6
Draw electron dot, valence and structural formulae for each of the following
a) F2
b) HCl
c) Cl2
d) O2
e) N2
f) H2O
g) CH4
h) CH3Cl
i) NH3
For each of the substances listed in Question 2 determine whether the
molecule is symmetrical or nonsymmetrical and polar or nonpolar.
Name the shapes of the following molecules by first drawing their valence
a) HBr
b) CO2
c) H2S
d) CH2Cl2
e) PH3
f) SF6
g) BCl3
h) CS2
i) PF5
Chlorine, oxygen and nitrogen are all gases at room temperatures which exist
as linear diatomic molecules. Describe the nature of the intermolecular bonding forces
between their molecules when these elements are in the liquid or solid states.
Draw valence structures for the following polyatomic molecules:
a) C2H6
b) C3H8
d) HCN
e) N2H4
f) H2O2
g) C2H2
h) SO3
Water (H2O) and hydrogen sulfide (H2S) are both V-shaped molecules with
identical valence structures. Whilst water has a boiling point of 100 0C hydrogen
sulfide vaporises at a temperature of just -62 0C. In terms of the intermolecular
bonding forces present explain the significant difference in boiling points between
these two substances.
Explain the variation observed in the melting points of the Noble Gases as
shown in the Table below by referring to the intermolecular bonding forces present.
Table 6.3 - Boiling points of the Noble Gases
Boiling Point (oC)
Carbon dioxide is a gas at room temperatures with a sublimation temperature
of -78 0C. Silicon dioxide is a solid with a melting temperature of 1700 0C. As both
carbon and silicon are in Group IV of the Periodic Table and so each has four valence
electrons we might reasonably expect their physical properties to be similar. Explain
this substantial difference in melting points.
Q10. The boiling point of phosphine (PH3) is -85 0C while that of ammonia (NH3),
also in Group V of the Periodic Table, is -33 0C. Explain this difference in boiling
points in terms of the intermolecular bonding forces present.
Q11. Classify each of the following molecules according to the type of
intermolecular bonding (dispersion forces, dipole-dipole attractions, hydrogen
bonding) which they would exhibit in the liquid state:
a) HCl
b) S2
c) NH3
d) CHCl3
e) OCl2
f) SiH4
g) N2
h) HF
i) NBr3
Q12. Solid carbon dioxide is widely known as ‘dry ice’. Explain the reasoning
behind this description. Use diagrams to assist in describing the intermolecular
bonding forces present in this substance.
Q13. Each of the elements of Period 2 (Li  Ne) is capable of forming a hydride [a
compound of the element with hydrogen] except neon.
a) Why does neon not form a compound with hydrogen?
b) With respect to the other seven hydrides which are formed, write the
formulae of the compounds formed.
c) Classify each of these hydrides as an ionic compound or a covalently
bonded molecule.
d) Draw valence structures for each of the hydride molecules and determine
the shape of the molecule formed.
e) Classify these molecules as either polar or nonpolar.
Q14. Explain why the molecule Mg2 does not exist under normal conditions of
temperature and pressure but the molecule O2 does.
Q15. By referring to the relevant structural formulae, give clear and concise
explanations for each of the following:
a) diamond is extremely hard whereas graphite is soft and greasy
b) graphite is an excellent lubricant which can be employed in high
performance racing car engines where the very high temperatures present would
vaporise conventional oil-based lubricants
c) diamond is widely used as an abrasive on the ends of drill bits
d) graphite is a good conductor of electricity whereas diamond is an excellent
Q16. From the following list choose a molecule which would meet the criteria as
listed below:
a) a substance which contains permanent dipoles
b) a substance containing linear molecules with pure covalent bonds
c) a covalent network lattice structure which is capable of conducting
d) a nonpolar molecule which nevertheless contains polar bonds
e) a substance which exhibits hydrogen bonding
f) a substance with an exceedingly low boiling point
carbon dioxide
hydrogen sulfide
hydrogen gas
liquid bromine
Q17. List the following molecules from lowest boiling point to highest. Explain the
reasoning behind your answer in terms of the relevant intermolecular bonding forces
present in each substance.
carbon dioxide, diamond, ammonia, oxygen
Q18. The following graph shows the boiling points of the Group VII hydrides as we
move down the Group. In terms of the intermolecular bonding forces present explain
the trend observed. Why does the boiling point of hydrogen fluoride not follow the
trend of the other hydrides listed?
[Insert graph to depict data as shown below. Sample 5 should not be there; I made a
mistake. Note that temperatures are measured in 0C. The X axis should indicate the
actual formulae , not merely numbers. It is not necessary to show actual tabulated
boiling points; it is the nature of the graph which is of relevance here.]
Boiing points ( C)
Group VII hydrides
Q19. Write the following compounds in order of their boiling points from lowest to
highest. Explain your reasoning in terms of the intermolecular bonding forces present.
butane (C4H10), octane (C8H18), ethane (C2H6), heptane (C7H16)
Q20. Ammonia (NH3) is highly soluble in water whereas octane (C8H18), the major
component of petrol, is effectively insoluble. Explain the difference in solubility of
these two solutes.
Ignore the box below. May be used elsewhere in the text.
Ionic bonding involves the donation of electrons from a metal atom to a nonmetal
atom. As the metal loses electrons it carries a positive charge and the nonmetal in
gaining electrons gains a negative charge. It is most important that we recognise that
these classifications of ionic and covalent bonding are somewhat arbitrary in nature
and indeed they should really be considered as part of a spectrum: