CHEMICAL BONDING
In order to understand how atoms bond we need to make use of various scientific models. A model
would be a representation of a real phenomenon which enables us to grasp or explain a physical
reality. We will be using the atomic model to develop and understanding of how substances combine
to form compounds.
TYPES OF BONDING
The forces of attraction that hold atoms, ions and molecules together are called chemical bonds. The
formation of a stable bond must result in a lower energy state. There are three major kinds of bonds:
1.
2.
3.
covalent bonds,
ionic bonds, and
metallic bonds.
THE COVALENT BOND
Consider the simplest molecule, hydrogen. If two hydrogen atoms approach each other from infinity,
two kinds of coulombic forces come into play:
forces of repulsion between two positively charged nuclei and two negatively charged electron
clouds, and
forces of attraction between oppositely charged nuclei and electron clouds.
The following discussion refers to the accompanying diagram.
(a)
At infinity the total energy of the system is
taken as zero.
(b)
As the two atoms approach each other, the
force of attraction between the positively
charged nucleus of one atom and the
negatively charged electron cloud of the
other atom increase therefore the potential
energy of the system decreases. This
effect is illustrated by the short-dashed
line in the figure.
Potential Energy (kJ.mol-1 )
BOND ENERGY DIAGRAM FOR HYDROGEN
At large values of inter-nuclear distance, d, the
total energy of the system is equivalent to the
combined energy of the two isolated hydrogen
atoms and is arbitrarily taken as zero. An
increase in potential energy is indicated by a
positive sign and a decrease by a negative sign.
+
0
d1
d
d1 = bond length
HH
H
H
H
H
Change i n potenti al energy as two atoms
approach each other
(c)
Simultaneously the forces of repulsion between the two electron clouds and the two positively
charged nuclei increases causing an increase in potential energy. This is illustrated by the longdashed line.
(d)
The overall change in potential energy resulting from theses two opposing interactions is
illustrated by the continuous line. At the internuclear distance, d1 (bond length), the total energy
of the system is lowest and the maximum stability attained. At this distance the orbitals overlap
to form a covalent bond.
Bond Energy: Energy required to completely separate two covalently bonded atoms.
Bond length: Distance between the nuclei of two covalently bonded atoms.
DIATOMIC MOLECULES OF DIFFERENT SIZE.
A single covalent bond is a bond that is formed by the sharing of a pair of electrons of opposite spin.
Double or triple bonds (multiple bonds) are formed when more than one pair of electrons are shared.
Octet rule
It has been experimentally observed that many stable molecules are composed of covalently bonded
atoms that acquire the noble gas configuration of the nearest noble gas. This is often a configuration
with 8 electrons in the outer shell. The tendency of atoms to react in such a way as to attain this
configuration is known as the octet rule.
Two helium atoms would not bond to form a diatomic molecule as hydrogen did above because helium
atoms have a full outer shell of electrons.
The octet rule says that atoms tend to gain, lose or share electrons so as to have eight electrons in
their outer electron shell.
Fluorine oxide (OF2)
Boron Hydride (BH3)
(Not all molecules adhere to the octet rule.)
Lewis and Couper Structures:
It has been experimentally observed that many stable molecules are composed of covalently bonded
atoms that acquire the noble gas configuration of eight electrons when bonding. The tendency of
atoms to react in such a way as to attain this configuration is known as the octet rule. Two helium
atoms would not bond to form a diatomic molecule as hydrogen did above because helium atoms have
a full outer shell of electrons.
The octet rule is used in writing Lewis structures. The valence electrons in the outer energy level of
the atom are shown as dots or crosses arranged in pairs around the symbol of the element. Electrons in
atoms having up to four valence electrons are often represented by single dots, but the dots may also
be paired in the way the electrons occur in the orbitals, for example, the Lewis structure of carbon can
be written as: :C: or C .
Electrons from different atoms are represented by dots and crosses respectively. The dots
x x and crosses
x
are arranged so as to give each atom a stable octet configuration, e.g. chlorine gas: Cl xC1
x and
xx
hydrogen chloride:
H x Cl
Couper structures:
If all bond pairs are represented by dashes and the lone pairs of electrons are omitted, the resulting
structure is known as a Couper structure, for example, HC1: H—C1.
Other examples of Lewis and Couper structures:
Hx
Nx H
x
H
NH3
H2O
H
x
Ox H
O2
O
xx
x x
xOx
CO2
H
N H
O H
x x
xC x
O C
O O
H
O
O
O
H
CH4
Hx
H C
xH
x
H
x
N2
N
xx
x
xN
x
H
H
C
H2S
H
x
S xH
HF
Hx F
H
H
N
N
S
H
H
F
H
Exercise: Draw Lewis and Couper structures for the following moleculesHCN
PCl3
BeF2
SBr2
HOCl
Dative covalent (co-ordinate) bonds:
A dative covalent bond is a bond that is formed by the overlapping of orbitals and the sharing of an
electron pair both of which belong to one atom. The atom that contributes the electron pair is called
the donor atom. Once a dative covalent bond is formed, there is no difference between it and an
ordinary covalent bond. The electron pair involved is usually represented by an arrow.
Examples:
ammonium ion (NH4+)
(a)
Hx
Hx N
x
H
+
+
H

+
Hx
H
Hx N
x
H
OR
+
Hx
Hx N
H
x
H
The hydrogen ion has no electrons in the outer shell whereas the ammonia molecule has a lone
pair. By bonding to the ammonia molecule, the hydrogen ion gains a share of two electrons.
The positive charge resides over the whole polyatomic ion.
oxonium or hydronium ion (H3O+)
(b)
H
x
Hx O
+ H+

OR
Shapes of molecules
The shapes of molecules are determined using valence shell electron pair repulsion theory. (VSEPR)
In short this theory looks at the effects of repulsion between electrons and electron pairs and can be
crudely summarised into a few major principles:



Molecular shape is determined by minimising repulsion between negative electron pairs.
Bonding pairs (BP) of electrons repel more than lone pairs (LP) of electrons:
BPBP>BPLP>LPLP
Electron pairs will always occupy the geometrical shape that corresponds to the lowest energy
configuration.
Examples of common molecular shapes are summarised below;
Linear, Trigonal, Tetrahedral, Trigonal pyramidal, Angular, Trigonal bipyramidal, Octahedral
1.
LINEAR
e.g. carbon dioxide
2.
CO2
TRIGONAL
e.g. boron trihydride BH3
3.
TETRAHEDRAL
e.g. methane CH4
4.
TRIGONAL PYRAMIDAL
e.g. ammonia NH3
5.
ANGULAR
e.g. water
H2 O
Bond angle 105o
6.
TRIGONAL BIPYRAMIDAL
e.g.phosphorus pentachloride PCl5
7.
OCTAHEDRAL
e.g. sulphur hexafluoride SF6
Polar molecules
Electronegativity is the tendency of an atom in a molecule to attract a shared pair of electrons.
When two atoms having the same electronegativity combine, the bond pair is shared equally between
the atoms, for example, H2. The electron cloud around the molecule is symmetrical and the zone of
highest probability density lies between the two nuclei. This is a NON-POLAR molecule.
If there is an appreciable difference in electronegativity between the atoms, for example, HCl, the
electron cloud will be displaced towards the more electronegative atom and there will be an
asymmetrical distribution of electrons.
The + indicates a small positive charge. Such a bond is said to be POLAR COVALENT.
A polar covalent bond will be formed if the difference in electronegativity (Eneg) lies in the range of
1,1 to 2,0. If the atoms of a diatomic molecule are identical the molecule is non-polar (Eneg = 0), and
if there is a difference in electronegativity greater than 2,1 the molecules is polar.
Electronegativity
difference (Eneg)
Pure covalent
(Non polar)
Covalent
(weakly polar)
Polar covalent
Ionic
0
0< Eneg <1
1< Eneg  2.1
2.1< Eneg
When a molecule consists of more than two atoms, both the difference in electronegativity between the
atoms and the shape of the molecule must be considered. A polar molecule must have a resultant
dipole moment to be polar.
A dipole is formed when equal and opposite charges are at a distance from one another and a dipole
moment is a measure of the polarity of a bond. A dipole moment is a vector quantity because it has
both magnitude and direction. An arrow is used to represent a dipole moment and indicate the
direction of displacement of the electron cloud.
The overall polarity of a molecule is equal to the vector sum of the dipole moments (polarities) of the
individual bonds. A molecule that has polar bonds can therefore be non-polar if the resultant of the
dipole moments is zero.
Examples:
HCl:
Electronegativity of C1:
3,0
Electronegativity of H:
2,1
Difference in electronegativity: 0,9
Shape of HCl molecule: linear
H Cl
Vector diagram:
BF3
The HCl molecule will be polar because it is a diatomic linear molecule with a difference
in electronegativity of 0,9 and a resultant dipole moment.
F
Electronegativity of F
4,0
Shape of molecule:
Electronegativity of H:
2,1
planar trigonal
120 o 120 o
B
Difference in electronegativity: 1,9
o
120
Orientation of dipole moments:
Vector diagram:
F
F
60o
120 o
120 o
120 o
60o
60o
The resultant is zero therefore the molecule is non-polar.
It is important to distinguish between a polar bond and a polar molecule. As can be seen
from the above example, a molecule having polar bonds can be non-polar because the
resultant of the dipole moments of the polar bonds is zero.
CO2
Electronegativity of C
2,5
Electronegativity of O:
3,5
Difference in electronegativity: 1,0
Shape of molecule: linear
0=C=0
Orientation of dipole moments:
Vector diagram:
Resultant of the dipole moments of the polar bonds is zero therefore the molecule is nonpolar.
H2O
Electronegativity of O
3,5
Electronegativity of H:
2,1
Difference in electronegativity: 1,4
Shape of molecule: angular
Orientation of dipoles
Vector Diagram:
O
104,5 o
H
104,5 o
R
CF4
Electronegativity of F:
4,0
Electronegativity of C:
2,5
Difference in electronegativity: 1,5
75,5o
Shape of molecule: tetrahedral
Note that the resultant of the dipole moments of a symmetrical tetrahedral molecule
composed of two elements is always zero. It is not possible to illustrate this as a
tetrahedron is three-dimensional.
The resultant of the dipole moments is zero therefore the molecule is non-polar.
H
THE IONIC BOND (Grade 10 Revision)
If the difference in electronegativity between two bonding atoms is greater that 2,0 an ionic bond is
formed. As a result of the large difference in electronegativity, the bond pair is attracted so strongly by
the more electronegative atom that the less electronegative atom loses control over its electron. The
electron from the less electronegative atom is transferred completely to the strongly electronegative
atom, and two oppositely charged ions are formed. The atom that loses an electron forms a positive
ion while the atom that gains an electron forms a negative ion. An ion may be regarded as a spherical
particle the charge of which is concentrated at the centre. The positive and negative ions are held
together in a crystal lattice by electrostatic forces.
An ionic bond is a bond that is formed by electrostatic attraction between positive and negative ions in
a crystal lattice.
Ionic bonding will take place when:
(a) the difference in electronegativity between the two atoms forming the bond is greater than 2,0;
(b) atoms having high electron affinity react with atoms having low ionisation energy;
(c) the reaction is exothermic.
Example: NaC1
Ions forming an ionic bond do not bond in pairs but are part of a crystal
lattice in which each ion attracts all the surrounding oppositely charged
ions as shown in the accompanying sketch. In the sodium chloride
crystal lattice, each positive ion is surrounded by six negative ions, and
visa versa. Since the numbers of positively and negatively charged ions
in a crystal lattice are the equal, the crystal lattice is overall neutral.
Sodium chloride can therefore be written as the formula Na+Cl- or
NaCl.
In the case of ionic compounds, the formula therefore indicates the ratio
in which the two ions occur and does not represent a molecule.
Representation of ionic bonding using Lewis structures
Formation of sodium chloride (NaCl):
.
Na  [Na]+ + e.. ..
:Cl. + e-  [ :Cl.
.. ]
..
Sodium loses an electron to form a sodium ion:
Chlorine gains an electron to form a chloride ion:
The positive sodium ion attracts the negative
chloride ion to form an ionic bond:
..
..
+
[Na]+ + [ :Cl.
]

[Na]
[
:Cl.
..
.. ]
Formation of magnesium chloride (MgCl2):
Magnesium loses two electrons to form a magnesium ion:
Two chlorine atoms each gain an electron each to form
two chloride ions:
One positive magnesium ion attracts two negative
chloride ions to form ionic bonds:
Mg:  [Mg]2+ + 2e..
..
2 :Cl.
+
2e

2[
:Cl.
.. ]
..
..
..
2+
[Mg]2+ + 2[ :Cl.
]

[Mg]
[
:Cl.
.. ] 2
..
Properties of ionic substances
1.
Ionic substances have relatively high boiling and melting points because of the strong
electrostatic forces. The smaller the ions and the larger their charges, the higher their boiling
and melting points.
2.
Solid ionic substances do not appreciably conduct an electric current at room temperature. The
ions are held firmly in the crystal lattice and electrons are strongly attracted to their ions.
3.
Molten ionic substances conduct an electric current. As the temperature of solid ionic
substances increases, their ionic vibrations become more vigorous until a stage is reached when
the crystal lattice forces are overcome and the ions move apart. In the liquid phase the ions are
able to move freely in an electric field and conduct a current.
4.
Ionic crystals are brittle and easily broken. When a stress is placed on them, the crystal lattice is
distorted. Similarly charged ions are forced closer together with a resultant increase in repulsive
forces. If this repulsion becomes sufficiently strong the crystal becomes unstable and breaks.
THE METALLIC BOND (G10 Revision)
Atoms of metals have two common characteristics: low
+ - + +- + -+ +- + - +
ionisation energies and a number of vacant orbitals. The
- - -- - - - - + + + + +-+ + +
atoms of a metal are arranged closely together in a crystal
- + - +- + + - + - + -+ -+
lattice. The space around each atom experiences the
same positive nuclear charge and forms a region of
uniformly low potential energy. As the ionisation energy
of metals is low, the electrons move away from the atoms
Mobile electrons
Positive atomic
and become delocalised, leaving behind positively
residues (ions)
charged ions. The positive ions are embedded in a sea of
delocalised electrons. The coulombic forces between the positive ions and negative electrons hold the
crystal lattice together and are called metallic bonds. The electrons can move freely throughout the
crystal and are not associated with any particular metal ion. The overall charge of the metal is neutral.
The nuclear charge and number of outer electrons determine the strength of a metallic bond. The
strength of the bond increases with an increase in the number of electrons. An increase in electrons is
accompanied by an increase of the nuclear charge and hence stronger forces of attraction between the
ions and delocalised electrons. Melting and boiling points increase as the strength of
the bond increases.
Properties of metals:
1.
Lustre. Metallic lustre is caused by the delocalised electrons that are made to vibrate with the
same frequency as the incident light. Vibration of a charged particle involves acceleration that
produces electromagnetic radiation having the same frequency as the incident light. The light
emitted therefore appears to be a reflection of the original light.
2.
Workability. When a stress is placed on a metal the metal ions are able to slide over each other
because metallic bonds are non-directional.
3.
Electric conductivity. The electrons are delocalised and are therefore mobile in an electric field.
If a potential difference is applied across the ends of the conductor, an electric field is set up in
the conductor. Negatively charged electrons are attracted to the positive terminal while the metal
ions remain stationary because they are held in a crystal lattice. The movement of electrons is
impeded by collisions with the positive metallic ions causing resistance. Every electron that
leaves a conductor at the positive terminal is replaced by another from the negative terminal.
The overall charge of a conductor is therefore neutral. . The resistance of a conductor increases
with increasing temperature because the increased vibration of atomic kernels at higher
temperatures interferes with the movement of electrons.
4.
Conduction of heat. When heated, electrons gain kinetic energy. The mobility of the electrons
enables them to move rapidly from regions of high temperature to cooler regions, transferring
some of their kinetic energy to the crystal lattice in the process.
5.
Electron emission. When a metal is exposed to heat or light its electrons gain kinetic energy. At
a certain stage their kinetic energy will be sufficient to enable them to overcome the attractive
forces of the crystal and break away. This emission of electrons is known as the thermionic
effect if caused by heat, or photo-electric effect of caused by light.
Extention Material: (Not for Exams)
A single covalent bond is a bond that is formed by the overlapping of orbitals and the sharing of a pair
of electrons of opposite spin.
In order for covalent bonding to occur, each atom must have an unpaired electron of opposite spin.
During bonding these electrons pair up to form a bond pair. The electron density is higher where the
orbitals overlap and the electrons may be regarded as occupying the valence orbitals of both atoms.
They are simultaneously attracted to both nuclei. Sharing is possible only if the difference in
electronegativity of the participating atoms is less that 2,0.
Overlapping may take place between two s orbitals, one s and one p orbital, or two p orbitals.
Overlapping between other orbitals is not required by the syllabus.
(a)
s-s overlap, for
 example H2
-bond

E
1
H
-bond

H
H
1s
1
+
H
H
H
1s
1s
1s
Formation of a hydrogen molecule
Electron configurations
The overlapping of the two s orbitals forms a sigma bond.
(b)
 HC1
s-p overlap, for example

-bond


3p



2p
-bond
y
1H

3s
E

x
3p
y
x


2s

1s
z
x
H
Cl
z
1s
lone pairs
s-p overl apping in the hydrogen
chlori de molecule
17Cl
Electron configurations
If the unpaired 1s electron of the hydrogen atom and the 3p electron of the chlorine atom have
opposite spins, they will overlap to form a sigma () bond. Orbitals containing paired electrons
do not participate in bonding. These electrons are called lone pairs.
(c)
When p orbitals overlap two possibilities arise: either end-to-end overlapping or sideways
overlapping can occur.
When end-to-end overlapping occurs a sigma covalent bond is formed, for example C12. The
half-filled orbitals of the two approaching chlorine atoms will overlap if the electrons in the
overlapping orbitals have opposite spin.


-bond


3p




2p


3s

3s
E

3p



2p

2s


2s
1s
17Cl
1s
-bond y
3p orbital
3p orbital y

x
x
x
Cl
xx
z
Cl
z z
x x

lone pairs
17Cl
End-to-end p-p overlapping in the
chlorine molecule to form a -bond
Electron configurations in the
chlorine molecule


-bond


-bond
If two p orbitals overlap sideways a covalent pi ( ) bond will be formed, for example O2.


2p

2p orbital


2s
E

2p
1s
O
8
x
x

2s

1s
8O

-bond
x

2p orbital
-bond
x
O
O
xx
z z
y
y
z
lone pairs
Sideways p-p overl apping in the oxygen
molecule to form a -bond
Electron configurations in the
oxygen molecule
Pi-bonds are weaker than sigma bonds because the overlapping is not as great.
The double bond involves the end-to-end overlap of two p orbitals and the sideways overlap
between an additional two p orbitals. A double bond therefore consists of one -bond and one
-bond.
A triple bond is formed
when three p orbitals from each of two atoms are involved in bonding.

The triple bond consists
of one -bond and two -bonds, for example, in N2.

-bond

-bond

-bond

E

2p



2p

2s


2s

1s
1s
N
7
7
N
Electron configurations in the N2 molecule

2p orbital
x
N
x
z
2p orbital
-bond
-bond
x
N
x
-bond
z
z
x
y
y
Sideways p-p overl apping in the N2 molecule to
form two -bonds
Note that the sideways overlapping of the two lobes of a p orbital on either side of the -bond
represents only one -bond.
The Theory of Hybridisation
The theory of hybridisation is not examinable but is helpful when trying to understand the shapes of
some molecules and why the carbon atom, which has only two unpaired electrons, has a valency of
four. It is therefore included in these notes as enrichment but discussion will be confined to the
methane molecule.
Methane (CH4)
The shape of the methane molecule can be explained in terms of sp3 hybridisation, i.e. the mixing of
one s and three p orbitals to form four sp3 orbitals. Hybridisation takes place in the following steps:
(a)
The approach of hydrogen atoms causes a rearrangement of electrons and energy changes to
take place in the carbon atom. The energy released as the hydrogen atoms approach the carbon
atom promotes one of the 2s electrons into the empty 2p orbital.

E


2p

2s
Promotion of
an electron

2p

2s
E


1s
6C (ground state)
(b)

1s
6C (excited state)
*
One 2s and three 2p orbitals then hybridise (mix) to form
new sp3 hybrid orbitals which have equivalent energy and
shape. The shape of a hybrid orbital is illustrated in the
adjacent sketch.

The new electron configuration is:


four
Shape of hybrid orbital

3
sp
E

1s
6C (hybridised state)
(c)
The four sp3 orbitals repel each other so that the hybrid orbitals
are tetrahedrally orientated. Four hydrogen atoms, each having
one electron, bond covalently to the sp3 orbitals. The bond
angle for a tetrahedral molecule is 109,5o.
The shapes of the ammonia and water molecules can also be
explained by the theory of hybridisation. The bond angles of
ammonia and water are 107,3o and 104,5o
respectively


-bond
1s orbital
H
sp3 hybrid
orbital
x
109,5o
C
C
x
H
H
x
x
H
The methane molecule
because ammonia has one lone pair of electrons and water has two lone pairs that distort the
shapes of the molecules.
Ionic Bonding
When sodium reacts with chlorine to form sodium chloride the reaction appears to take place in one
simple step although it is probably more complex. The actual steps are not known but can be
visualised as a series of five hypothetical steps. All but one of these steps involve transfers of energy
that have been experimentally measured.
Step 1: Conversion of sodium from the solid to the vapour phase. Such a phase change directly from
a solid into a vapour without going through the liquid phase is known as sublimation, and is
an endothermic process.
Na(s) + energy  Na(g)
Step 2: Conversion of gaseous chlorine molecules into gaseous chlorine atoms. This process is called
dissociation and is endothermic.
½Cl2(g) + energy  C1(g)
Step 3: Conversion of gaseous sodium atoms into ions in the gaseous phase, i.e.
ionisation. This process is endothermic.
Na(g) + energy  Na+(g) + eStep 4: Conversion of chlorine atoms into chlorine ions in the vapour phase.
This process is exothermic and the energy given out is called electron affinity.
C1(g) + e-  C1-(g) + energy
Step 5: Attraction of positive sodium ions and negative chlorine ions in the vapour phase. This
process is exothermic and the energy given out is known as the crystal lattice energy.
Na+(g) + Cl-(g)  Na+Cl-(s) + energy
Crystal lattice energy is the energy liberated when one mole of an ionic crystal is formed from
gaseous ions.
4
3
E
2
0
1
5
6
1. sublimation energy
2. dissociation energy
3. ionisation energy
4. electron affinity
5. lattice energy
6. energy of formation
Factors affecting bond strength
The force of attraction between oppositely charge ions depends on the size of the charge on the
ions and the distance between them. As the charge on an ion may be regarded as being situated
at the centre, the force of attraction may be determined using Coulomb's Law. The larger the
ionic radii, the greater the distance between the ions and the weaker is the force of attraction.
This accounts for the fact that sodium chloride has a stronger ionic bond than CsC1. The Cs ion
has a larger radius than the sodium ion. The ions in CsC1 are therefore further apart than in
sodium chloride, and the ionic bond is weaker.