BONDING
Name: ………………………………………….
A chemical bond is a method of holding atoms or ions together. You should already
be familiar with the three types of chemical bond: ionic, covalent and metallic.
1. Ionic Bonding




Metals, since they are on the left hand side of the Periodic Table, tend to
lose electrons to gain a stable full outer shell.
The tendency for an atom to lose electrons is greatest at the bottom left of
the Periodic Table.
Non-metals, at the opposite end of the Periodic Table, tend to gain
electrons.
The tendency for a non-metal to gain electrons is greatest at the top right
of the Periodic Table.
If a metal and a non-metal (e.g. sodium and chlorine) come into contact, there is a
tendency for electrons to be transferred from the metal to the non-metal.
The metal atom loses electrons and becomes a positively charged ion (cation).
The non-metal gains electrons and becomes a negatively charged ion (an anion).
These opposite charges attract and the ions are held together by strong
electrostatic forces of attraction. The formation of an ionic bond is represented
by a dot and cross diagram:
TOPIC 12.3: BONDING & PERIODICITY
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Exercise 1: Draw diagrams to show the ionic bonding in
(a)
lithium fluoride
(b)
magnesium chloride
(c)
calcium oxide
(d)
sodium oxide
TOPIC 12.3: BONDING & PERIODICITY
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The Structure of Ionic Solids
A description of the bonding in a compound describes the way two atoms or ions
are held together, as above.
The structure of a solid describes how the many particles that make up that solid
are arranged and held together.
A solid with a regular shape which contains particles arranged in a regular
structure is called a crystal.
The physical properties of that solid crystal are determined by the bonds and
forces that hold the crystal together.
In an ionic solid, the individual particles are positive and negative ions (cations and
anions), arranged in a regular array (a lattice) e.g. sodium chloride, NaCl:
Na+ ion
Cl- ion
The ions in a sodium chloride crystal are arranged in a 3-dimensional giant ionic
lattice. Each sodium ion has a coordination number of 6 (i.e. each Na+ is surrounded
by 6 Cl-) and each chloride ion has a coordination number of 6. The sodium chloride
giant ionic lattice can thus be described as a 6:6 lattice or a face centred cubic
lattice (because the resulting crystals are cubes).
TOPIC 12.3: BONDING & PERIODICITY
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Each sodium ion is in contact with six chloride ions; each chloride ion is in contact with
six sodium ions.
In reality, the ions are touching, as shown in the following diagrams.
single layer
Other ionic compounds could have different lattice structures. The structure
depends on the relative size of the anions and cations. For example, caesium
chloride has a body centred cubic lattice and is described as an 8:8 lattice.
TOPIC 12.3: BONDING & PERIODICITY
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Explaining the physical properties of ionic solids
You should be familiar with the properties of ionic solids. These properties can be
explained by reference to the structure of the compounds.
Exercise 2: Explain the following properties of ionic solids in terms of their
structure.
1. Ionic compounds have high melting and boiling points
2. Ionic compounds are hard
3. Ionic compounds are brittle
4. Ionic compounds are non-conductors of electricity when solid, but will
conduct when molten (in the liquid state) or in solution
TOPIC 12.3: BONDING & PERIODICITY
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Polarisation of Ions
(Ionic Compounds with Covalent Character)
An ionic compound forms when electron transfer takes place between a metal and
a non-metal to produce a cation and an anion, which are then held together by
strong electrostatic forces.
If the electron is completely transferred, and the ions formed are perfectly
spherical, the bonds are perfectly ionic. This is more likely to happen between
metals at the bottom left of the periodic Table, and non-metals at the top right
e.g. caesium fluoride
Cs+
F-
However, if the positive ion (cation) formed is very small and/or highly charged,
and the negative ion (anion) is large and/or not highly charged the electron cloud
around the anion is distorted. It is no longer spherical. Because the electron
density of the anion is now partially localised between the two nuclei, the resulting
ionic compound is said to have covalent character e.g. lithium iodide:
Li+
I-
The type of bonding in lithium iodide is said to be ‘ionic with covalent character’.
A cation which is small and/or highly charged is said to be polarising. It polarises
the anion, as shown by Li+ above. The greater the charge density (i.e. a high charge
on a small cation) the greater the polarising power of the cation.
An anion which is large is said to be polarisable i.e. it can be polarised, as in the
case of I- above. The larger anions are more easily polarised by a cation.
TOPIC 12.3: BONDING & PERIODICITY
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Which ionic compounds are likely to display covalent character?
The Metal Cation
Cation size increases down a Group, so the most polarising cations are likely to be
at the top of a Group.
Cation size decreases across a Period (from Group I to Group III), so the
polarising power increases from left to right on the Periodic Table.
Thus, metal cations from the top of Groups, and furthest across the Period on the
Periodic Table are likely to be the most polarising e.g. Be2+ and Al3+, and are more
likely to form ionic compounds with covalent character.
The Non-Metal Anion
Anion size increases down a Group, so the most polarisable (the most easily
distorted) anions are likely to be near the bottom of a Group.
Anion size decreases across a Period (from Group V to Group VII), so the most
polarisable anions are P3-, then S2- and then Cl-.
Thus, non-metal anions from the bottom of Groups, and furthest across the Period
on the Periodic Table are likely to be the most polarisable (Their spherical shape is
most easily distorted) e.g. I-, and are more likely to form ionic compounds with
covalent character.
The effect of covalent character on physical properties
An ionic compound with covalent character is likely to show properties which
resemble those typical of covalent compounds.
For example, an ionic compound which displays covalent character is likely to have a
lower melting point than expected. It is likely to be softer, less brittle and show
reduced electrical conductivity when molten or in solution.
TOPIC 12.3: BONDING & PERIODICITY
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Exercise 3: Circle the ionic compound in each of the pairs below that exhibits the
most covalent character. Give a reason for your choice.
Ionic Compounds
LiCl
and
Reason
NaCl
LiCl
and
LiI
NaCl
and
MgCl2
CaO
and
CaS
MgCl2
and
AlCl3
KBr
and
KF
TOPIC 12.3: BONDING & PERIODICITY
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Special cases: Aluminium Chloride and Beryllium Chloride
The Al3+ ion has such a small size and high charge that it is highly polarising. It has
a very high charge density.
When aluminium forms a compound with oxygen, Al2O3, the ionic bonding has only
very slight covalent character because, despite the high charge density of the Al3+,
the O2- ion is not very polarisable.
When aluminium bonds with chlorine, however, the resulting chloride ions are
polarised to such an extent, and the bond has such a large amount of covalent
character, that any ionic description of the bonding is useless.
The bonding in AlCl3 is COVALENT, and AlCl3 is a molecule, despite the compound
being formed from a metal and a non-metal. It can be drawn as follows:
AlCl3 molecule:
The same argument and reasoning would apply to BeCl2, which is also a covalent
compound made of a metal and a non-metal.
Two AlCl3 molecules form an Al2Cl6 dimer – see notes on covalent bonding.
Across Period 3 from Na to Si, the type of bonding in the chlorides changes as
follows:
NaCl
ionic
MgCl2
ionic with a little
covalent character
TOPIC 12.3: BONDING & PERIODICITY
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AlCl3
covalent
SiCl4
covalent
Melting an Ionic Solid
In an ionic crystal, e.g. sodium chloride, the ions, formed by electron transfer, are
arranged in a regular array and are all held together by strong electrostatic
forces. The ions vibrate around a fixed point. This description of a solid should be
familiar from your GCSE course.
If the system is heated (i.e. it takes in energy), the ions vibrate faster and faster
until the melting point is reached.
At the melting point, the particles have taken in enough energy to partially
overcome the strong electrostatic forces holding them in a fixed position. They
can now move around each other. The electrostatic forces still act, however, to
keep the particles close together in the liquid (molten) state, but the ions are free
to move around. Again, the motion of the ions in the molten state should be familiar
from your GCSE description of the motion of particles in a liquid.
Because the melting process requires (takes in) energy, it is described as
ENDOTHERMIC.
The large amount of energy required to partially overcome the many strong
electrostatic forces is the reason that ionic solids have high melting points.
If the liquid (molten) ionic compound continues to be heated, the ions take in more
energy and move faster and faster until the boiling point is reached. At this point,
the ions possess enough energy to almost completely overcome all electrostatic
forces, and the ions move away from each other in random, chaotic motion. It is
now a gas.
TOPIC 12.3: BONDING & PERIODICITY
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2. Covalent Bonding




Non-metal atoms need to gain electrons to achieve a full outer shell.
When two non-metals bond, neither one will give up an electron, so the atoms
share a pair of electrons.
This means that each atom has a full outer shell, despite some of the
electrons in the outer shell being shared with another atom.
A covalent bond is defined as a shared pair of electrons.
A covalent bond can be represented by a dot and cross diagram. Only the outer,
bonding electrons need to be shown.
Hydrogen fluoride
The dot and cross diagram can be simplified, and the covalent bond shown as a
single line between two atoms. Every time you see a diagram with a line of this
sort, it means the atoms are covalently bonded together.
Hydrogen fluoride
TOPIC 12.3: BONDING & PERIODICITY
Chlorine
11
Exercise 4: Draw similar, simplified diagrams for the following molecules. You may
find it easier to first sketch the dot and cross diagrams. On your finished
diagrams, indicate the number of bonding and non-bonding pairs (known as lone
pairs) of electrons.
(a) methane (CH4)
(b) ammonia (NH3)
(c) water (H2O)
(d) hydrogen sulphide (H2S)
TOPIC 12.3: BONDING & PERIODICITY
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Dative Covalent (Coordinate) Bonding
The covalent bonds studied so far have been formed between two atoms, each of
which donates a single electron to make a shared electron pair.
However, both electrons may originate from the same atom, and be shared with
another atom.
A dative covalent bond (also known as a coordinate bond) is defined as a shared
pair of electrons where both electrons have originated from one atom.
For this to happen, one species must have a lone pair of electrons which it can
donate to another atom in order for the electrons to be shared. This atom is the
electron pair donor. The other atom acts as an electron pair acceptor.
A dative covalent bond is identical to a normal covalent bond except in the way it
forms.
The formation of a dative covalent bond can be represented as shown below. The
lines represent a shared pair of electrons, and the two crosses, xx, represent the
lone pairs.
Boron trifluoride gas and ammonia gas react together readily to give a white solid.
The boron atom acts as an electron pair acceptor (so BF3 is known as a Lewis acid –
an electron pair acceptor) because it is electron deficient.
The nitrogen atom in ammonia acts as an electron pair donor (so NH3 is known as a
Lewis base – an electron pair donor).
TOPIC 12.3: BONDING & PERIODICITY
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Exercise 5: When ammonia and hydrochloric acid react they produce a salt called
ammonium chloride. Draw a diagram to show how ammonia can react with the H + ion
in the acid.
Exercise 6: Two AlCl3 molecules form a dimer of Al2Cl6. Draw a diagram to show
how this dimer forms.
Exercise 7: When copper(II)sulphate solution is a blue colour due to the presence
of the [Cu(H2O)6]2+ complex ion. This forms between Cu2+ ions and water molecules.
Draw a diagram to show how this complex ion forms.
TOPIC 12.3: BONDING & PERIODICITY
14
Electronegativity and Bond Polarity



In a covalent bond, a pair of electrons is shared between two atoms.
Some atoms attract the bonding pair of electrons in a covalent bond more
than others.
The ability of an atom to attract the bonding pair of electrons in a covalent
bond is known as the electronegativity of the atom.
What factors do you think might affect how electronegative an atom is?
Linus Pauling assigned electronegativity values for each element. The scale was
arbitrary, and ranged from 0 (not electronegative, i.e. does not attract the
electron pair in a covalent bond) to 4 (the most electronegative – attracts the
electron pair in a covalent bond very strongly). This is shown for selected elements
below.
2.1
H
1.0
Li
0.9
Na
0.8
K
0.8
Rb
0.7
Cs
1.5
Be
1.2
Mg
1.0
Ca
1.0
Sr
0.9
Ba
2.0
B
1.5
Al
How does electronegativity vary across a Period? Why?
How does electronegativity vary down a Group? Why?
Which is the most electronegative atom? Why?
Why are the Noble Gases not included in this table?
TOPIC 12.3: BONDING & PERIODICITY
15
2.5
C
1.8
Si
3.0
N
2.1
P
3.5
O
2.5
S
4.0
F
3.0
Cl
2.8
Br
2.5
I
If a covalent bond forms between atoms with DIFFERENT electronegativities, the
electrons are not shared equally. The shared pair is held closer to the more
electronegative atom. This type of covalent bond is known as a polar covalent
bond, and it results from the differing electronegativities of the two atoms
involved.
The atom which is the most electronegative of the two will have the greater share
of the bonding pair of electrons. Because electrons are negatively charged, the
more electronegative atom gains a very slight negative charge and is said to be
electron rich. This is shown using a small - symbol next to the atom.
Similarly, the atom which has the lowest share of the electron pair is said to be
electron deficient and has a slight positive charge. This is shown using a small +
symbol next to the atom.
Bond polarisation is caused by the unequal sharing of electrons due to the
differing electronegativities of the atoms involved. As a result of bond
polarisation, the covalent compound may have some degree of ionic character,
because the + and - charges represent slight charges.
Exercise 8: Use the electronegativity values to deduce which atoms are electron
deficient and which are electron rich. Indicate, using + and - symbols on the
appropriate atoms, the bond polarisation of any one of the bonds in the molecules
shown below. If the bond is not polar, write ‘non-polar’ underneath the structure.
(a)
(b)
O
H
H
N
H
(c)
H
H
(d)
Br
Cl
H
Cl
Al
Cl
(d)
N
Cl
Cl
Cl
TOPIC 12.3: BONDING & PERIODICITY
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Calculating Percentage Ionic Character
Linus Pauling estimated the amount of partial ionic character in a covalent bond in a
binary compound (a compound consisting of 2 atoms). He calculated the
electronegativity difference of the two atoms involved in the bond, and thus
estimated the percentage ionic character of the resulting molecule.
Electronegativity
Difference
% Ionic
Character
Electronegativity
Difference
% Ionic
Character
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
0.9
1.0
1.1
1.2
1.3
1.4
1.5
1.6
0.5
1
2
4
6
9
12
15
19
22
26
30
34
39
43
47
1.7
1.8
1.9
2.0
2.1
2.2
2.3
2.4
2.5
2.6
2.7
2.8
2.9
3.0
3.1
3.2
51
55
59
63
67
70
74
76
79
82
84
86
88
89
91
92
Exercise 9: Use the electronegativity values to deduce the percentage ionic
character in the following binary compounds:
Compound
% Ionic Character
Chlorine, Cl2
Hydrogen chloride, HCl
Sodium chloride, NaCl
Magnesium oxide, MgO
Caesium fluoride, CsF
TOPIC 12.3: BONDING & PERIODICITY
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Symmetry and bond polarity
In symmetrical molecules with polar covalent bonds, the bond polarities can ‘cancel
out’. The results is a molecule which, overall, is non-polar but which contains polar
bonds.
-
+
-
O=C=O
 
Despite the two polar C=O bonds, CO2 is a non-polar molecule.
Exercise 10: Indicate if the following molecules are overall polar or non-polar.
(a)
HCl
(b)
NH3
(c)
H2C=O
(d)
H 2O
(e)
BeH2
(f)
AlF3
TOPIC 12.3: BONDING & PERIODICITY
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Water is a polar molecule:
Exercise 11: The structure of chloromethane and tetrachloromethane are shown
below. For each carbon-chlorine bond, show the bond polarity by using the symbols
- and + to indicate the electron rich and electron deficient atoms. Draw a circle
around the molecule that does not possess an overall dipole.
Cl
H
Cl
H
H
chloromethane
Cl
Cl
Cl
tetrachloromethane
Both molecules have 4 bonds arranged in a tetrahedral fashion around the carbon
atom.
Exercise 12: Carbon dioxide (CO2) contains polar C=O bonds but it is a
symmetrical molecule so it shows no overall dipole. Sulphur dioxide (SO 2) does have
an overall dipole. Draw the structure of an SO2 molecule, which could be consistent
with the information given. Hint: carbon has 4 outer electrons and sulphur has 6.
TOPIC 12.3: BONDING & PERIODICITY
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The Structure of Covalent Solids
There are two main types of covalent structures: Giant Covalent and Simple
Molecular. The two types have particles arranged and held very differently in the
solid state, and so have very different physical properties (e.g. melting point,
conduction of electricity, hardness, malleability).
(1) Giant Covalent Structures
Giant covalent structures have of a regular array (lattice) of atoms held by many
strong covalent bonds. Examples are diamond and graphite (see below).
The giant covalent structures are very strong. As a consequence the melting points
are very high because a large amount of energy is required to break the many
strong covalent bonds.
They are insoluble in water (a polar solvent) and organic solvents (non-polar
solvents) because the solvent cannot break down the lattice structure.
Giant covalent structures do not conduct electricity because there are no
free/mobile electrons (graphite is an exception and can conduct electricity – see
below).
Examples of giant covalent structures are diamond, graphite and silicon dioxide:
Diamond
structure
Silicon
dioxide
structure
Graphite
structure
TOPIC 12.3: BONDING & PERIODICITY
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Examples of Giant Covalent Structures:
Diamond, silicon and silicon dioxide (SiO2)
Graphite

3-dimensional tetrahedral lattice

hexagons of carbon atoms
arranged in layers (sheets)

Each carbon atom has a coordination
number of 4

Each carbon has a coordination
number of 3

All carbons joined to others by strong
covalent bonds

Each carbon is covalently bonded
to 3 others.

Diamond cannot conduct electricity
because there are no delocalised
(mobile) electrons

The remaining electron from the
carbon is delocalised between the
layers (sheets). These delocalised
electrons mean that graphite can
conduct electricity.

All the atoms are held by 4 very
strong covalent bonds. This makes
diamond very hard.

The layers (sheets) are held
together by weak Van der Waals’
forces (see later). This means
that the structure is easily
broken along the layers (the
layers can slide over each other).
TOPIC 12.3: BONDING & PERIODICITY
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Exercise 11: Explain the following observations
(a) diamond and graphite have high melting points.
(b) graphite will conduct electricity but diamond will not.
(c) graphite is soft, but diamond and silicon are very hard.
TOPIC 12.3: BONDING & PERIODICITY
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(2) Simple Molecular Structures
Simple molecular structures are made up of molecules held together by weak
intermolecular forces.
A molecule is defined as one or more atoms held together by covalent bonds.
Examples are Cl2, H2O, HCl, P4 and S8 (shown below).
A P4 molecule
A S8 molecule
Intermolecular forces are forces which act between molecules to hold them
together. They are weak forces. There are several different types of
intermolecular forces that can act between molecules.
In order to understand the structure and properties of simple molecular crystals,
it is essential that the intermolecular forces that hold the molecules in the solid
are fully understood.
There are three types of intermolecular forces you need to consider:
1. van der Waals forces
2. permanent dipole-dipole interactions
3. hydrogen bonding (which is not a bond but an intermolecular force for A
level purposes!!)





All intermolecular forces are weak forces – nowhere near as strong as covalent,
ionic or metallic bonds.
The three types, however, have different strengths. They are listed above in
order of increasing strength, so hydrogen bonding is the strongest
intermolecular force, then permanent dipole-dipole interactions. The weakest
intermolecular forces are van der Waals forces.
All species, even the atoms of the Noble Gases, are attracted to each other by
van der Waals forces.
Only polar molecules contain permanent dipole-dipole forces in addition to van
der Waals.
The only molecules that contain hydrogen bonding are those that have a
hydrogen atom covalently bonded to a N, O or F atom within the molecule.
TOPIC 12.3: BONDING & PERIODICITY
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Intermolecular forces
Van der Waals forces
Forces of attraction exists between all molecules. These weakest of these forces
are known as van der Waals forces.
They originate when two molecules approach each other. The electron cloud within
the molecule can become repelled at one end by repulsion from the electron cloud
of the second molecule. This causes a temporary dipole, which in turn causes
(induces) an opposite dipole on the adjacent molecule. The leads to very weak
electrostatic attractions between molecules.
The greater the size of the molecule, the greater the number of electrons and so
the greater the probability that electrons will be temporarily localised at one end
of the molecule. Larger molecules also have a greater surface area for the van der
Waals forces to act over.
Thus, the larger the molecule, the stronger the van der Waals forces between the
molecules. Stronger forces between molecules mean that the melting and boiling
point of the solid is higher.
The more branched an organic molecule is, the lower the surface contact area and
thus the weaker the van der Waals forces and the lower the boiling point.
Van der Waals forces are the only intermolecular forces that act between nonpolar molecules, and are the forces that hold molecules together in a solid and a
TOPIC 12.3: BONDING & PERIODICITY
24
liquid. There are even very weak van der Waals forces between gas atoms or
molecules.
TOPIC 12.3: BONDING & PERIODICITY
25
Permanent dipole-dipole attractions
We have already studied some molecules that are polar. Bond polarities arise when
the electronegativities of the two atoms that form the covalent bond are
different. Polar molecules are those which are non-symmetrical and which contain
polar covalent bonds. Examples of polar molecule are HCl and propanone (shown
below).
When two polar molecules arrange themselves in a solid, opposite charges align to
give small, very weak electrostatic attractions between the molecules. This is
known as a permanent dipole-dipole attraction, and it helps to hold the simple
molecular structures of polar molecules together.
Van der Waals forces would also act between polar molecules, but they are weaker
than permanent dipole-dipole attractions. This means that simple molecular solids
containing molecules held together by both van der Waals and permanent dipoledipole attractions are likely to have higher melting and boiling points than those
with a similar mass but only van der Waals forces acting between molecules.
e.g.
butane, C4H10, with an MR of 58, has a boiling point of 0oC
CH2
H3C
CH3
CH2
propanone, C3H6O, also with an MR of 58, has a boiling point of 56oC
O
C
H3C
CH3
TOPIC 12.3: BONDING & PERIODICITY
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Hydrogen bonding (it’s not a bond, it’s an intermolecular force!)
A hydrogen bond can be thought of as a very extreme example of a permanent
dipole-dipole attraction.
A hydrogen bond is the force of attraction between a very electron deficient
hydrogen atom and the lone pair from an atom on a neighbouring molecule.
In order that the hydrogen atom is sufficiently electron deficient, it must be
directly covalently bonded to a small electronegative atom. The only three atoms
which are both small enough and electronegative enough to cause this are
NITROGEN, OXYGEN and FLUORINE. The diagram below shows the hydrogen
bonding between two water molecules.
x
x
+
-
The reason hydrogen is a special case is that it has no inner shells of electrons.
When hydrogen is directly bonded to N, O or F, the electron density is drawn away
from the hydrogen, leaving the nucleus exposed. The nucleus of the poorly shielded
hydrogen atom attracts the lone pair of electrons from an atom on the next
molecule. There is some resemblance to a covalent bond – although a hydrogen bond
is nowhere near as strong as a covalent bond (approximately 5-10% the strength).
This explains why the hydrogen bond is stronger than a permanent dipole-dipole
attraction.
Hydrogen bonding explains the unusually high melting and boiling points of
molecules such as water. With an MR of only 18, water should really be a gas at
room temperature. If it wasn’t for the hydrogen bonding increasing the melting and
boiling point so that it is a liquid at room temperature, life on earth would not exist
as we know it today!
TOPIC 12.3: BONDING & PERIODICITY
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Exercise 12: Explain why the boiling points of the halogens increase from fluorine
(F2) to iodine (I2).
Exercise 13: Predict, with a brief reason, which organic compound in each pair
would have the higher boiling point. Circle the compound with the highest boiling
points (NB All hydrocarbons are considered to be non-polar molecules).
Organic Compound
CH2
H3C
CH3
and
Reason
H3C
CH3
CH3
H3C
CH
CH2
CH3
H3C
CH2
and
CH2
H3C
CH3
CH2
OH
H3C
Cl
H3C
CH2
and
Cl
H3C
C
H3C
CH2
and
CH2
H3C
H3C
CH3
CH2
C
and H3C
TOPIC 12.3: BONDING & PERIODICITY
28
O
Exercise 14: State the name of the strongest intermolecular force present
between molecules of each of the species below. (Lone pairs have not been included
to simplify the diagrams).
Br
O
Br
H
H
Br
H
CH3
S
H
H
H
F
H
H3C
C
C
CH3
H3C
O
Cl
H
O
O
H3C
H
C
OH
H
H
C
H
C
H
TOPIC 12.3: BONDING & PERIODICITY
N
H
H
N
H
Cl
Cl
29
Cl
Exercise 15: Sketch 2 graphs on the same axes to show the boiling points of the
hydrides of (a) Group VI and (b) Group VII using the data below:
Group VI Hydride
H 2O
H 2S
H2Se
H2Te
Boiling Point/K
373
212
232
271
Group VII Hydride
HF
HCl
HBr
HI
Boiling Point /K
293
188
206
238
Boiling Point /K
Hydride
Explain why there is a regular increase in the boiling points of the hydrides from
the 2nd to the 4th member of the group.
Explain why the first hydride in each group shows an exceptionally high boiling
point when compared to the rest of the compounds.
Draw a similar graph below and draw the predicted shapes of the graphs for the
group V hydrides (NH3 to SbH3) and group IV hydrides (CH4 and SnH4).
TOPIC 12.3: BONDING & PERIODICITY
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Exercise 16: Draw diagrams to show the formation of one hydrogen bond between
(a) molecules of hydrogen fluoride and (b) molecules of ammonia. Show all the
electron deficient and electron rich atoms (using + and - symbols) and include all
lone pairs.
(a) Hydrogen fluoride, HF
(b) Ammonia, NH3
TOPIC 12.3: BONDING & PERIODICITY
31
The Structure of Ice
Now that we have an understanding of the forces that can hold molecules together
in simple molecular structures, we can look at some examples.
Water has extensive hydrogen bonding between molecules and it can form a three
dimensional lattice. It is possible for 4 hydrogen bonds to form between each
water molecule – two from each lone pair, and two to each hydrogen atom per
molecule.
The diagram below shows how it is possible for each water molecule to connect to 4
other molecules through hydrogen bonds (note the loose similarity to a diamond
type structure – although in the case of water the structure is held together by
hydrogen bonding which is much weaker than the covalent bonds that hold the
giant covalent diamond structure together).
The structure above explains why ice has a lower density than water (i.e. why it
expands on freezing, and why ice floats on water) at 0oC.
The hydrogen bonds hold ice in an open structure. When ice melts, the lattice
breaks up, the water molecules pack closer together and the maximum density
occurs at 4oC.
As the temperature increases, the hydrogen bonds are overcome, more spaces
develop between the molecules so the density decreases.
TOPIC 12.3: BONDING & PERIODICITY
32
The Structure of Iodine
In iodine, the I2 molecules (consisting of two iodine atoms held together by a
covalent bond) are held together in a three dimensional lattice by weak van der
Waals forces. They have a ‘herringbone’ pattern.
The solid is only held be very weak intermolecular forces, and so on gentle heating
it will sublime.
Molecular lattices do not conduct electricity as there are no charged particles
(either electrons or ions) free to move through the lattice.
Many molecular solids (e.g. iodine, phosphorus and sulphur) are insoluble in water.
Can you think why this would be? What type of solvents would dissolve these types
of solids?
What type of substances would dissolve in water?
TOPIC 12.3: BONDING & PERIODICITY
33
Shapes of Molecules
Molecules are different shapes because electron pairs will maximise their distance
apart in order to minimise repulsion between them.
Basic Shapes of Molecules
These basic shapes should all be learned. You should be able to draw accurate
diagrams, name the shapes and label the bond angles.
No. of bonding
electrons
Name of
shape
Bond Angles
Example
2
linear
180o
BeCl2
3
trigonal
planar
120o
BCl3
4
tetrahedral
109.5o
CH4
5
trigonal
bipyramidal
120o and 90o
PF5
6
octahedral
90o
SF6
TOPIC 12.3: BONDING & PERIODICITY
34
Diagram
Exercise 17: Practice your diagrams below for CH4, PF5 and SF6. State the name
of the structure underneath, and the bond angles.
Origin of names of shapes of molecules
Molecules are named after the three dimensional shapes they would form if their
outer atoms were positioned at the vertices e.g.
Methane, CH4, is tetrahedral:
PF5 is trigonal bipyramidal:
SF6 is octahedral:
TOPIC 12.3: BONDING & PERIODICITY
35
The presence of lone pairs
If not all the pairs of electrons around the central atom are bonding electrons, the
shape can be affected.
Each electron pair occupies a volume of space, which can be regarded as a cloud of
electron density. A lone pair of electrons repels more than a bonding pair of
electrons because it takes up more space at the surface of the atoms, and is under
the influence of only one nucleus.
The order of repulsion is:
lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
Ammonia, NH3, and water, H2O, have 4 electron pairs around the central atom, but
some are lone pairs. Ammonia has 3 bonding pairs and 1 lone pair around the
nitrogen. Water has 2 bonding pairs and 2 lone pairs around the oxygen.
The shape of both molecules is thus based on tetrahedral (because they have 4
electron pairs around the central atom), but lone pairs occupy some of the
positions around the atom. The name of the shape is determined by the position of
the atoms and not the lone pairs, thus the name changes.
The bond angle is also affected, since lone pairs repel more than bonding pairs.
Molecule
Diagram
Shape
Bond
Angle
tetrahedral
109.5o
pyramidal
107o
bent planar
104.5o
H
Methane,
CH4
C
H
H
H
..
Ammonia,
NH3
N
H
H
H
..
Water, H2O
O
H
H
..
..
TOPIC 12.3: BONDING & PERIODICITY
36
Determining the Shape of Molecules
The arrangement in space of single covalent bonds around the central atom (and
thus the shape of the molecule) can be predicted by considering the number of
electrons in the outer shell of the central atom, and the number of bonds that it
forms.
For a given molecule, the shape can be predicted using the rules below:
1. Count the total number of electrons in the outer shell of the central atom.
2. If the species is an ion, add 1 for each negative charge and subtract 1 for
each positive charge.
3. Add 1 for each atom bonded to the central atom.
4. Total the number of electrons.
5. Divide by 2 to determine the total number of electron pairs. This tells you
the shape your molecule is based on.
6. Count the number of bonded atoms. This tells you the number of bonding
pairs of electrons. Any remaining bonds are lone pairs, and their effects must
be considered.
7. State the name of the shape, draw it and show any bond angles.
Example: Draw and name the shape of PCl3, indicating any bond angles on your
diagram. Following the steps as outlined above:
1.
2.
3.
4.
5.
6.
7.
5 electrons around P (Group 5)
0 to add or subtract – it’s not an ion
3 bonding Cl atoms
Total number of electrons = 5 + 0 + 3 = 8 electrons
8 ÷ 2 = 4 PAIRS OF ELECTRONS
(so at this stage you know that the shape is based on tetrahedral –
like methane)
3 bonded atoms so 3 bonding pairs. The remaining pair is a lone pair.
(so I can draw a basic tetrahedral shape, put three Cl atoms at three
of the positions, and show a lone pair in the remaining position.)
Sketch the shape, state the name of the shape and draw on the bond
angles (remembering to adjust the shape name and bond angle due to
the presence of the lone pair):
TOPIC 12.3: BONDING & PERIODICITY
37
Exercise 18: Predicting the shapes of molecules and ions. Draw and name the
shapes of the following molecules and ions. Include all the lone pairs, where
appropriate, and label the bond angles. Do your working on a separate piece of
paper.
BeCl2
NH3
CCl4
NH2-
H 3O +
AlF63-
XeF4
SiH4
AlH4-
H 2S
PF5
NH4+
I 3-
SF4
BrF3
AlI3
(hint: use one I as the
central atom)
TOPIC 12.3: BONDING & PERIODICITY
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3. Metallic Bonding
Metallic bonding, as the name suggests, occurs between metal atoms. Metals tend
to want to lose electrons, and they have the ability to lose their outermost
electrons to form cations.
A metal can be thought of as a giant lattice of close-packed metal cations
surrounded by a sea of delocalised electrons. These delocalised electrons come
from the outer shell of the metal ion, and are free to move through the lattice.
A metallic bond is the electrostatic force of attraction that two neighbouring
cations have for the delocalised electrons between them.
+
+
+
+
+
+
+
+
+
+
+
free electron
+
+
+
+
metal ion
Diagram to show metallic bonding
Close packed metal ions in the Giant Metallic Solid:
Metals have high melting points due to the strong electrostatic attraction between
the metal cations and the sea of delocalised electrons.
Conduction of electricity is possible due to the delocalised (and therefore mobile)
electrons, which can move freely through the structure.
The metallic bond strength varies between metals. This can be explained by the
charge on the cation, the size of the cation and the number of delocalised
electrons holding the cations together. The highest melting point will occur when
the cation size is small, the cation charge is large and the number of delocalised
electrons is large.
Exercise 19: State then explain the trend in melting points
(a)
across Period 3 (sodium to aluminium)
(b)
down Group 1 (sodium to rubidium)
TOPIC 12.3: BONDING & PERIODICITY
39
STATES OF MATTER
SOLIDS
In a solid, particles (which may be atoms, ions or
molecules) are arranged in a definite, ordered regular
pattern in three dimensions, and the particles are very
tightly packed. This arrangement manifests itself on a
macroscopic scale in the crystalline structure of
solids. At absolute zero, the particles are stationary,
but as the solid warms up, the particles vibrate about
a mean position.
As the temperature increases, the vibrations increase in amplitude and eventually
become so great that the attractive forces between the particles in the solid are no
longer sufficient to hold the structure together. At this point melting occurs.
A solid has a definite shape and a definite volume.
The different types of solid structure: ionic, simple molecular, macromolecular and
metallic, together with their associated properties have been discussed in detail in an
earlier section.
LIQUIDS
The liquid state is the least understood of the three states of matter.
In a liquid, the tendency of particles to stick together
because of their mutual attraction outweighs the
tendency to remain apart because of their thermal
energy. The thermal energy is however too great to
allow the particles to occupy fixed positions.
Thus, the particles in a liquid are still fairly tightly
packed together but lack the highly ordered
arrangement of a solid. Liquids are often described as
having short range order and long range disorder.
This means that over a small region (1 to
2 nm) there is order comparable to that in a crystal. On a larger scale there is much
disorder caused by the presence of ‘holes’ in the structure. These holes allow individual
particles to have translational movement, which means that particles can move past
each other. Structures similar to those in a solid, but only partially complete, exist
temporarily, constantly breaking and reforming in a different way as particles move
through the liquid.
The structure of a liquid manifests itself on a macroscopic scale by it having a definite
volume but no definite shape.
The particles in a liquid have a range of kinetic energies. Some of the faster moving
particles are able to overcome the attractive forces in the liquid and escape into the
space above the liquid, forming a vapour. This vapour exerts a pressure known as the
vapour pressure.
TOPIC 12.3: BONDING & PERIODICITY
40
As the temperature of a liquid
increases, the mean kinetic
energy of its particles increases,
and so more particles are able to
escape to form a vapour.
Therefore, as the temperature of a
liquid increases, its vapour
pressure increases.
When the vapour pressure of the
liquid becomes equal to the
atmospheric pressure above it, the
liquid boils.
Vapour
pressure
Temperature
That a liquid more closely resembles the solid state than the gaseous state is indicated
by the following observations:
 Density
There is only a small difference between the density of a solid and that of the
liquid formed from it.

Energy Differences
Much more energy is needed to convert a liquid to a gas than a solid to a liquid.
e.g. H2O(l)
H2O(g)
H = +40.6 kJ.mol-1
H2O(s)
H2O(l)
H = +6.0 kJ.mol-1
GASES
Gases are devoid of structure. Since the
energy of the particles is sufficient to overcome
the force of attraction between them, the
particles in a gas are disordered and widely
separated. The much greater separation of
particles in a gas, compared to solids and
liquids, can be seen when 1 mole of water
(approximately 18cm3 at room temperature)
gives 1 mole of steam (approx. 33000cm 3 at
100oC and 1 atm.)
The particles in a gas are in constant and rapid motion in straight lines. They are
constantly colliding with each other and with the walls of the container.
The pressure exerted by a gas depends on the number and the energy of the collisions
which the particles make with the container walls per unit area per unit time. As the
temperature of a gas increases, the mean kinetic energy of its particles increases.
Collisions with the container walls are more energetic and occur more frequently,
therefore the pressure increases.
A gas is confined only by its container. Therefore, it has no definite shape and no
definite volume, taking up instead the shape and volume of its container.
TOPIC 12.3: BONDING & PERIODICITY
41
Glossary of Terms
Word / Term
Definition / Description
IONIC BONDING
ionic bond
the attraction between oppositely charged ions formed by
electron transfer from a metal to a non-metal
polarising
a description given to a cation of high charge density (small
and highly charged) e.g. Al3+ that can distort the spherical
nature of an anion, causing an ionic bond to have covalent
character
polarisable
a description given to a large anion e.g. I- that has low control
over its electron cloud due to high shielding of the nuclear
charge, and which can be distorted away from its spherical
shape by a small, highly charged cation
cation
a positively charged ion
anion
a negatively charged ion
COVALENT BONDING
molecule
a group of two or more atoms held together by covalent bonds
e.g. Cl2, H2O, NH3, CH4, I2, BeCl2, BF3, PF5, SF6, AlCl3
covalent bond
a shared electron pair between two atoms
dative
covalent a shared electron pair between two atoms where both
bond
electrons have originated from one atom (also known as a
coordinate bond)
coordinate bond
a shared electron pair between two atoms where both
electrons have originated from one atom (also known as a
dative covalent bond).
electronegativity
the relative tendency of an atom to attract the electrons in a
covalent bond
polar bond
a covalent bond where the electron pair is shared unequally
due to the difference in electronegativities of the atoms in
the bond.
Lewis acid
an electron pair acceptor
Lewis base
an electron pair donor
lone pair
a pair of non-bonding electrons
dimer
a species formed between two small molecules e.g. Al2Cl6
electron rich
an atom which has a greater share of electron density (shown
by a - sign)
electron deficient an atom in a polar covalent bond which has the smallest share
of the electron pair (shown by a + sign)
ionic character
A covalent compound which shows some properties
characteristic of an ionic compound. Caused by the unequal
sharing of electrons in a covalent bond. Shown by polar
TOPIC 12.3: BONDING & PERIODICITY
42
covalent compounds.
polar molecule
a polar molecule is a non-symmetrical molecule which contains
polar bonds. If the molecule contains polar bonds but is
symmetrical, the overall molecule is non-polar.
STRUCTURES OF SOLIDS CONTAINING COVALENT BONDS
giant covalent
a type of structure. A lattice of atoms held together by many
strong covalent bonds. E.g. diamond, graphite, silicon, silicon
dioxide.
simple molecular
a type of structure. A lattice of molecules held together by
weak intermolecular forces.
intermolecular
the weak forces between molecules in a simple molecular
forces
solid.
van der Waals the weakest intermolecular forces. A temporary, induced
forces
dipole on all molecules caused by temporary repulsions of the
electron cloud in the molecules, causing the molecules to be
weakly attracted to each other. The only type of
intermolecular force present in solid iodine.
permanent dipole- an intermolecular force. Weaker than hydrogen bonding, but
dipole attractions stronger than van der Waals forces. The weak electrostatic
attraction between the permanent dipoles on polar molecules.
hydrogen bonding the strongest intermolecular force. The attraction between a
very electron deficient hydrogen atom (which is covalently
bonded to N, O or F) and the lone pair of electrons on a
neighbouring molecule. Responsible for the high melting and
boiling points of water.
ice
solid water! A simple molecular solid held together by
hydrogen bonding intermolecular forces.
SHAPES OF MOLECULES
linear
the basic shape of a molecule with two electron pairs around
the central atom. Bond angle 180o. E.g. BeCl2
trigonal planar
the basic shape of a molecule with three electron pairs around
the central atom. Bond angles 120o. E.g. BCl3
tetrahedral
the basic shape of a molecule with four electron pairs around
the central atom. Bond angles 109.5o. E.g. CH4
trigonal
the basic shape of a molecule with five electron pairs around
bipyramidal
the central atom. Bond angles 120o and 90o. E.g. PCl5
octahedral
the basic shape of a molecule with six electron pairs around
the central atom. Bond angles 90o. E.g. SF6
pyramidal
the shape of a molecule with 3 bonding pairs and one lone pair
e.g. ammonia. The bond angle is reduced from the tetrahedral
angle of 109.5o to 107o due to the greater repulsion of the
lone pair of electrons.
TOPIC 12.3: BONDING & PERIODICITY
43
bent planar
the shape of a molecule with 2 bonding pairs and 2 lone pairs
e.g. water. The bond angle is reduced from the tetrahedral
angle of 109.5o to 104.5o due to the greater repulsion of the 2
lone pairs of electrons.
METALLIC BONDING
metallic bond
the electrostatic force of attraction between the regular
array of metal cations and the sea of delocalised outer
electrons.
MISCELLANEOUS TERMS
crystal
a solid made up of a regular array (lattice) of particles
lattice
a regular pattern or 3-dimensional array of inter-linked
particles. A lattice can be used to describe the regular
arrangement in any of the four main structure types. Lattices
can be ionic e.g. an alternating face centred cubic
arrangement of Na+ and Cl- ions in NaCl held by electrostatic
forces, giant covalent e.g. the regular tetrahedral
arrangement of carbon atom in diamond held together by
many strong covalent bonds, simple molecular e.g. the lattice
of I2 molecules held together by weak van der Waals forces
or metallic.
solid
the lowest energy state of matter. Particles vibrate about a
fixed position. On heating, the particles vibrate more
vigorously until they have enough energy to move away from
the fixed position and move around. Incompressible due to
small spaces between particles.
liquid
particles move randomly throughout the bulk of the liquid.
There are still some forces of attraction acting between
particles. Slightly compressible because particles have larger
spaces between them.
gas
particles move with rapid random motion. Easily compressible
due to the large spaces between particles.
melting
The change of state from a solid to a liquid.
freezing
The change of state from a liquid to a solid.
boiling
The change of state from a liquid to a gas.
condensing
The change of state from a gas to a liquid.
sublimation
The change of state from a solid straight to a gas (missing out
the liquid state). Can also be used to describe the change
from a gas to a solid e.g. I2 (g) = I2 (s)
evaporation
The change in state from a liquid to a gas below the boiling
point of the substance.
TOPIC 12.3: BONDING & PERIODICITY
44
PERIODICITY
CLASSIFICATION OF ELEMENTS
1s
2p
s- 3d
block
p-block
d-block
4f
f-block
PERIODIC TRENDS IN PERIOD 3 (Na-Ar)
1. Atomic Radius
0.16 x
0.14
x
Atomic radius
/nm
x
0.12
x
x
x
x
0.10
Na
Mg
Al
Si
P
S
Cl
Ar
Atomic radius decreases across the period.
 the nuclear charge is increasing
 electrons are entering the same shell, so the shielding is constant

the outer electrons are more strongly attracted and are drawn closer to the
nucleus
TOPIC 12.3: BONDING & PERIODICITY
45
2. First Ionisation Energy
1600
x
x
1200
x
First ionisation
energy /kJ.mol-1
800
x
x
x
x
x
400
Na
Mg
Al
Si
P
S
Cl
Ar
This trend has already been discussed in STRUCTURE AND BONDING.
3. Melting & Boiling Points
3000
x
x
2500
x
Tb
x
Tm
2000
Temperature
1500
/K
x
x
x
1000
x
x
500 x
x
x
x
x
P
S
xx
x
Cl
Ar
0
Na
Mg
Al
Si
The melting and boiling points reflect the structure and bond strength of the elements.
Sodium, magnesium and aluminium are metals and have metallic bonding. From Na to
Al, the size of the ion decreases due to the increasing nuclear charge, and the number
of delocalised electrons increases. Therefore, the strength of the metallic bond
increases. This results in an increase in both the melting point and the boiling point.
Silicon has a macromolecular structure in which there are strong covalent bonds in
three dimensions, like diamond. Since a large amount of energy is required to break
these bonds, the melting and boiling points are high.
TOPIC 12.3: BONDING & PERIODICITY
46
Phosphorus (P4), sulphur (S8) and chlorine (Cl2) are simple molecular substances with
weak van der Waals’ forces holding the molecules together. Since van der Waals’
forces are weak, all three elements have fairly low melting and boiling points. The
strength of the van der Waals’ forces increases as the size of the molecule increases.
Therefore, sulphur, which has the largest and most polarisable molecule, has the
highest melting and boiling points. This is followed by phosphorus, then by chlorine.
Argon is monatomic and has very weak van der Waals’ forces between its atoms. Its
melting point (84K) and boiling point (87K) are both very low.
TOPIC 12.3: BONDING & PERIODICITY
47