Download (iii) Formation of Hydrogen chloride molecule

CHEMICAL BONDING
INTRODUCTION
Why do atoms combine to form compounds?



Elements that do not have stable
electron configuration of noble-gas
attain such a configuration through
chemical reactions.
During the process their outer shell
either holds eight electrons (octet) or
two (doublet) electrons.
The process of attaining stable electron
configuration is called chemical
bonding.
Chemical bonding
Chemical bonding is the chemical process in
which atoms of elements combine to form
molecules or compounds.
Types of Chemical Bonding
There are three types of bonding, namely:
(i)
Ionic or Electrovalent Bonding.
(ii)
Covalent Bonding and.
(iii) Metallic bonding.
1.
Ionic (Electrovalent) Bonding
Ionic bonding is a type of bonding where
electrons are completely transferred
from one atom to another atom or group.


The strong force of attraction
between the oppositely charged
positive and negative ions results
in an ionic bond or electrovalent
bond.
In an ionic structure, the ions are
arranged in a regular repeating
pattern called lattice.
Facts about ionic bonding

The electron donor loses electron(s)
equal to its valence.

The electron(s) move to the outer shell
of the electron accepter (nonmetallic atom). Both atoms acquire the
electronic structure of a noble gas.

In ionic bonding, oppositely charged
ions cations and anions are formed.

The ions attract each other by
electromagnetic force and form ionic
bonds.


Dot-and-cross diagrams
We use dots and crosses to show the
electronic configuration of the metallic
and non-metallic ions so as to help us
keep track of where the electrons
come from and where they move to.
A dot-and-cross diagram shows:
(i) the outer electron shell only
(ii) that the charge of the ion is
spread evenly, by using square
brackets
(iii) the charge on each ion, written at
the top right-hand corner of the
square brackets.
NB:

When drawing a dot-and cross diagram
for an ionic compound, we draw the
outer electron shell of the metal ion
without any electrons.
Some examples of dot-and-cross diagrams
(i) Sodium chloride

When sodium reacts with chlorine
to form sodium chloride the one
electron in the outer shell of each
sodium atom are transferred to
the incompletely filled orbitals of
a chlorine atom.

By losing one electron, each
sodium atom achieves the electron
configuration [2,8] of neon.

By gaining one electron, each
chlorine atom achieves the
electron configuration [2,8,8] of
argon.

The sodium ion (Na+) and the chlorine
ion (Cl-) attract to form sodium chloride.
+
Na
Cl
Na
2,8,1
2,8,7
[2,8]+
-
Cl
[2,8,8]-
(ii)
Calcium oxide

When magnesium reacts with
oxygen to form magnesium oxide
the two electrons in the outer shell
of each magnesium atom are
transferred to the incompletely
filled orbitals of an oxygen atom.

By losing and gaining two
electrons, both magnesium atom
and oxygen atom respectively
achieve the electron configuration
[2,8] , noble-gas configuration of
neon.

The magnesium ion (Mg2+) and the
oxygen ion (O2-) attract each other to
form magnesium oxide.
2-
2+
Mg
O
2,8,2
2,6
Mg
[2,8]2+
O
[2,8]2+
(iii) Calcium chloride

Each calcium atom has two
electrons in its outer shell. The two
electrons are transferred to two
chlorine atoms.

By losing two electrons, each
calcium atom achieves the electron
configuration [2,8,8] noble-gas
configuration of argon.

The two chlorine atoms each gains
one electron and achieves the
electron configuration [2,8,8],
noble-gas configuration of argon.

The calcium ion (Ca2+) and the two
chlorine ions (Cl-) combine to form
calcium chloride.
-
Cl
2+
Ca
2,8,7
Cl
2,8,8,2
2,8,7
Ca
Cl
[2,8,8]-
[2,8,8]2+
Cl
[2,8,8]-
2.
Covalent bonding
Covalent bonding is the type of bonding
in which one or more pairs of electrons
are shared by the two atoms.

A covalent bonding occurs
between non-metallic atoms.
Dot-and-cross diagrams for covalent bonding
When drawing the arrangement of electrons
in a molecule, we
(i) Use ‘dot’ for electrons from one of the
atoms and ‘cross’ for the electrons
from the other atom.
(ii) If there are more than two types of
atom we can use additional symbols
such as a small circle or a small
triangle.
(iii) Draw the outer electrons in pairs, to
emphasise the number of bond pairs
and the number of lone pairs.
(iv) Put the pair of shared in the region of
the overlapping shell.
Types of covalent bonds
There are three types of covalent bonds,
namely;
(i)
Single covalent bond.
(ii)
Multiple covalent bonds.
-
Double covalent bond
-
Double-double covalent bond
Co-ordinate or dative covalent
bond
(a)
Single covalent Bonds

A single covalent bond or a bond
pair shared pair of electrons.

A single covalent bond is
represented by a single line
between the atoms.
Examples of single covalent bonds
among others are seen in;
(i) Hydrogen molecule,
(ii) Chlorine molecule,
(iii) Water molecule
(iv) Hydrogen chloride
(v) Methane
(i)
Formation of hydrogen molecule (H2)

A hydrogen molecule is formed
by combination of two hydrogen
atoms.

The two atoms approach one
another and their outer shells
overlap.

Each atom contributes one
electron to form a single covalent
bond.
H
(1)
+
H
(1)
H
H
(2) (2)
H-H
(ii)
Formation of Chlorine molecule

A chlorine molecule is formed
when two chlorine atoms combine.

The two chlorine atoms approach
one another and their outer shells
overlap.

Each atom contributes one electron
to form a bond pair.
Cl
(2,8,7)
+
Cl
(2,8,7)
Cl
Cl
(2,8,8) (2,8,8)
Cl-Cl
(iii) Formation of Hydrogen chloride
molecule

A hydrogen chloride molecule is
formed when one hydrogen atom
combines with one chlorine atom.

The two atoms approach one
another and their outer shells
overlap.

Each atom contributes one electron
to form a bond pair.
H
(1)
+
Cl
(2,8,7)
H
Cl
(2) (2,8,8)
H-Cl
(iv) Formation of water molecule

Water molecule is formed when
one oxygen atom combines with
two hydrogen atoms.

The three atoms approach one
another and their outer shells
overlap.

The oxygen atom contributes two
electrons and each hydrogen atom
contributes one electron two single
covalent bonds.
(v)
Formation of Methane molecule

Methane molecule is formed when
one carbon atom combines with
four hydrogen atoms.

The four hydrogen atoms approach
the carbon one. The outer shells
overlap.

Each atom contributes one electron
resulting to four single covalent
bonds.
(2,8)
2
H
(1)
+
O
(2,6)
H
(2)
O
H
(2)
H
4
H
(1)
+
C
(2,4)
H
O
H
H
H
H
C
H
Each hydrogen atom shares
two electrons with carbon
H
(b)
Multiple Covalent Bonds

Atoms can form covalent bonds by
sharing more than one pair of
electrons between them.

The types of multiple bonds are;
(i)
Double bond.
E.g Oxygen (O2), Ethene (C2H4)
(ii)
Double-double bond
E.g. Carbon dioxide (CO2).
(iii) Triple bond.
E.g. Nitrogen (N2)
(i)
Formation of oxygen molecule

Oxygen molecule is formed by
combination of two oxygen atoms.

The two atoms approach one
another and their outer shells
overlap.

Each atom contributes two
electrons resulting to two double
bonds.
O
(2,6)
+
O
(2,6)
O
O
(2,8 (2,8)
O=O
(ii)
Formation of Ethene molecule

Ethene molecule is formed by
combination of two carbon atoms
and four hydrogen atoms.

The two carbon atoms approach
one another and their outer shells
overlap.

Each carbon atom contributes two
electrons resulting to two double
bonds between the carbon atoms.

Of the four hydrogen atoms, two
atoms overlap with one carbon
atom to form four bond pairs.
H
4
H
(1)
+ 2
C
(2,4)
H
O
H
O
H
(iii) Formation of Carbon dioxide molecule
(CO2)

Carbon dioxide molecule is formed
by combination of one carbon
atom and two oxygen atoms.

The two oxygen atoms approach
the carbon atom on either side and
their outer shells overlap.

The carbon atom contributes all
the four valence electrons and
each oxygen atom contributes a
pair of electrons resulting to
double, double bonds.
(2,6)
O
+
O
(2,6)
C
(2,4)
O
O
O=C=O
(iii) Formation of Nitrogen molecule

Nitrogen is formed by combination
of two nitrogen atoms.

The two nitrogen atoms approach
one another and their outer shells
overlap.

Each atom contributes three
electrons to complete the octet
stable structure resulting to a
triple bond.

A triple bond is represented by
three short lines (≡) between the
nitrogen atoms.
N
(2,5)
+
N
(2,5)
N
N
N≡N
(c)
Co-ordinate bonding (dative covalent
bonding)
A co-ordinate bond is a type of covalent
bond where the pair of shared electrons
is donated by one atom.
For a dative covalent bonding to occur,
(i)
The donor atom must contain
a lone pair (a pair of electrons
that are not involved in bonding).
(ii)
The acceptor atom must have an
unfilled orbital.

A co-ordinate bond is represented by an
arrow pointing from the donor to the
recipient atom.
Examples of co-ordinate covalent
bonding are;
-
Ammonium ion.
-
Aluminium chloride molecule.
(i)
Formation of ammonium ion (NH4+)

Ammonium ion is formed by
combination of ammonia molecule
and a hydrogen ion.

The nitrogen atom in ammonia has
a lone pair of electrons.

It donates the lone pair of electrons
to the hydrogen ion.
H
H
H
+
C
+
+
H
H
NH3 (g)
H
O
H
+
H+ (aq)  NH4+ (aq)
H
(ii)
The formation of Aluminium chloride
molecule

Aluminium chloride is a covalent
molecule. At high temperatures it
exists as molecules with the
formula AlCl3.

At lower temperatures two
chlorine atoms each on a separate
molecule donates a lone pair of
electrons to an aluminium atom on
separate molecule to form a
molecule Al2Cl6 .

The molecule Al2Cl6 has two
co-ordinate covalent bonds.
Cl
Cl
Cl
Al
Cl
Al
+
Cl
Cl
Cl
Al
Al
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Al
Cl
Cl
Al
Cl
Cl
Shapes of Simple Molecules

The shape of a refers to the
three-dimensional arrangement of the
atoms that constitute a molecule.

It determines several properties of a
substance including; reactivity, polarity,
phase of matter, colour, magnetism and
biological activity.

The 'Electron-pair' Repulsion Theory is
used to predict shapes and bond angles
of simple molecules.
Electron-pair Repulsion Theory
The main points of electron-pair repulsion
theory are:
(i)
(ii)
In a molecule there are two types of
active set of electron pairs
surrounding the central atom;
-
the bond pair and
-
lone pair.
The electron pairs repel one another.
(iii) The repulsion forces the pairs of
electrons apart until the repulsive
forces are minimised.
(iv) The force of repulsion between lone
pairs and bond pairs is not the same.
The order of repulsion
The order of repulsion is as follows:
lone pair-lone pair > lone pair-bond pair > bond
pair-bond pair.
Factors which determine the shape of a
molecule

The shape of a molecule depends upon
total number of bond pairs and lone
pairs of electrons surrounding the
central atom.
How to Determine a Molecular Structure
1.
Draw the Lewis Structure.
2.
Count the bonds and lone pairs about
the central atom.
NB: Count the multiple bonds as one region
of electron density.
3.
Determine the basic arrangement of all
electron density regions (bonds +
lone pairs) about the central atom.
4.
Determine the molecular geometry
based on resulting positions of atoms.
How to predict the shapes of molecules
using bond pairs only
(i)
If the central atom has two bond pairs
and no lone pair or has two double
bond, the shape is linear with bond
angles of 180o.
E.g. Carbon dioxide (CO2)
(ii)
If the central atom has three bond
pairs, and no lone pair, the shape is
trigonal with bond angles of 120o.
E.g. Boron triflouride (BF3)
(iii) If the central atom has four bond pairs
and no lone pair, the shape is
tetrahedral with bond angles of 109.5o
E.g. Methane (CH4)
(iii) If the central atom has six bond pairs
and no lone pair, the shape is
octahedral with bond angles of 90o
E.g. Sulfur hexafluoride (SF6)
How to predict the shapes of molecules
using bond pairs and lone pairs
(i)
If the central atom has two bond pairs,
and two lone pairs, the shape is bent or
non-linear with bond angles of 104.5o.
E.g. Water (H2O)
(ii)
If the central atom has three bond
pairs and one lone pair, the shape is
pyramidal with bond angles of 107o.
E.g. Ammonia (NH3)
Shapes of molecules without lone pair
Examples of molecules without lone pairs
are; BF3, CO2, CH4 and SF6.
(i)
The shape of boron triflouride (BF3)

The central atom, B has three bond
pairs in its outer shell.

Minimizing the repulsion causes
the boron triflouride molecule to
have a trigonal planar shape with
a bond angle of 120o.
F
B
F
F
(ii)
The shape of Carbon dioxide (CO2)

The central atom, C atom has
four bond pairs in its outer shell.

The molecule has two double
bonds.

Minimizing the repulsion causes
the carbon dioxide molecule to
have a linear shape with a bond
angle is 180o.
O
O
(iii) The shape of Methane (CH4)

The central atom, C atom has
four bond pairs in its outer shell.

Minimizing the repulsion causes
the bond pairs arrange themselves
about the C atom resulting to a
tetrahedral shape with a bond
angle of 109.5o.
H
H
O
H
H
H
H
C
H
H
(iv) The shape of Sulfur hexaflouride (SF6)

The central atom, S atom has six
bond pairs in its outer shell.

Minimizing the repulsion causes
the bond pairs to arrange
themselves about the S atom
resulting to a octahedral shape
with a bond angle of 90o.

The 6 F atoms are located at the
corners of the octahedron.
F
F
F
S
F
F
F
Shapes of molecules with lone pair
Examples of molecules without lone pairs
are; H2O and NH3.
(i)
The shape of water (H2O)

The central atom, O has two bond
pairs and two lone pairs in its
outer shell.

Minimizing the repulsion causes
the water molecule to have a bent
or non-linear shape with a bond
angle is 104.5o.
(ii)
The shape of Ammonia (NH3)

The central atom, N atom has three
bond pairs and a lone pair in its
outer shell.

The lone pair repels the bond pairs
and then arrange themselves about
the N atom, thus resulting to a
trigonal pyramid or pyramidal
shape with a bond angle of 107o.
H
N
H
H
Summary of shapes of molecules
Molecule Bonding
pair
Lone
pairs
Bond
angle
Shape
CO2
2
0
180o
Linear
BF3
3
0
120o
Trigonal
CH4
4
0
109.5o
NH3
3
1
107o
H 2O
2
2
104.5o
0
90o
SF6
6
Tetrahedral
Pyramidal
Bent
Octahedral
 bonds and ) bonds
(i)
(ii)
 bond (Sigma bond)
A sigma bond is a bond formed a
hybridised p orbital (modified p orbital)
overlaps linearly (end-on) with an s
orbital or another modified p orbital.
 bond
A  bond is a bond formed when
p orbitals overlap sideways.

A single  bond is drawn as two
electron clouds one arising from
each lobe of the p orbitals.
Shapes of some organic molecules
Ethane (CH4) and Ethene (C2H4)

Shapes of molecules are explained in
terms of the patterns of electron
density found in
(i)
 bonds and  bonds.
Shape of Ethane (CH4)
Ethane has a displayed formula as;


All the bonds in ethane are  bonds.
All the areas of electron density repel
each other equally. This makes the
H – C – H bond angles all the same
(109.5o.
(ii)
Shape of Ethene

The displayed formula of ethene is

Each carbon atom in ethene uses
three of its four outer electrons to
form  bonds.

Two  bonds are formed with the
other carbon atom.

The fourth electron from each carbon
atom occupies a p orbital, which
overlaps sideways with a similar
p orbital on the other carbon atom to
form a
 bond.
(iii) The shape of benzene

The molecular formula of benzene
is C6H6.

The shape of benzene is a planar
regular hexagon, with bond angles
of 120°.
The symbol for benzene

All the bonds are identical. This means
that there aren't alternating double and
single bonds. The electrons are
delocalised.

Therefore, the symbol of benzene is;
NB:


The hexagon shows the ring of 6 carbon atoms,
each of which has one hydrogen attached.
The circle represents the delocalised electrons.
Formation of Benzene structure

In benzene, each of the 6 carbon atoms
undergoes sp2 hybridisation.

Each atom uses three hybrid orbitals to
form three sigma bonds, two with two
carbon atoms and one with a hydrogen
atom.

The p orbitals in the same plane in each
carbon atom overlap sideways to form
3

 bonds
.
The electrons in the
delocalised.
 bonds are
H
H
H
H
H
H
H
H
H
H
H
H
Bond energies, Bond lengths and Bond
polarities
(i)
Bond length
Bond length or bond distance is the
average distance between nuclei of two
bonded atoms in a molecule.

In general double bonds are shorter
that single bonds.
Reason:

Double bonds have a greater
quantity of negative charge
between the two atomic nuclei.

The greater force of attraction
between the electrons and the
nuclei pulls the atoms closer
together.
(ii) Bond energy
Bond energy is the energy needed to
break one of a given bond in a gaseous
molecule.

It is usually expressed in kJ mol-1.
Table showing some values of bond lengths
and bond energies
Bond
Bond Energy/(kJmol-1)
Bond length/nm
C-C
350
0.154
C=C
610
0.134
C-O
360
0.143
C=O
740
0.116
Relationship between Bond length and Bond
energy

For non-metallic atoms in the same
group bonded to the same element, the
bond energy decreases down the group
while the bond lengths increase down
the group.
Example
The table below shows the bond lengths
and bond energies of some.
Hydrogen halide
Bond
length/nm
Bond Energy/
(kJmol-1)
H - Cl
0.127
431
H - Br
0.141
366
H-I
0.161
299
(a)
What is the relationship between the
bond lengths and the bond energy for
these hydrogen halides?
(b)
Suggest why the bond energy values
decrease in the order HCl > HBr > HI.
(c)
Suggest a value for the bond length in
hydrogen fluoride, HF.
Electronegativity and Bond polarity
(i)
Electronegativity
Electronegativity is the ability of a
particular atom, which is covalently
bonded to another atom, to attract the
bond pair of electrons towards itself.

The greater the value of the
electronetavity, the greater is the
power of an atom to attract the
electrons in a covalent bond
towards itself.
Solution
(a)
The bond energies of the halides
decrease down the group since the
bond lengths of the halides increase
down the group.
(b)
The atomic size of the halogens
increase down the group. Therefore,
the reactivity of the halogens decrease
down the group.
(c)
0.122 nm
The pattern of Electronegativity in the
Periodic Table
Electronegativity;
(i) General increases across a period from
Group I to Group VII.
Reason:

Due to increasing nuclear charge
and decreasing atomic radius.
(ii)
Decreases down or increases up each
group.
Reason:

Due to increasing atomic radius
and shielding effect.
For the most electronegative elements, the
order of electronegativity is;
Increasing electronegativity
F > O > N > Cl > Br > C > I > H
Electronegativity values for some elements
Element
Electronegativity
F
4.0
Cl
3.0
Br
2.8
S
2.5
C
2.5
H
2.1
(ii)
Bond polarity

The polarity of a covalent bond
depends on the difference of the
electronegativity values of the
atoms forming the bond.

The difference can make the
covalent bond to be either non-polar
or polar.
How to predict the type of bond using
electronegativity differences
The difference in electronegativities
between the two elements is determined.

(i)
If the difference is
ionic.
1.7, the bond is
(ii)
If the difference is < 1.7 but > 0.5, the
bond is polar.
(iii) If the difference is < 0.5, the bond is
non-polar.

Non-polar Bonds

When the electronegativity values
of the two atoms forming a
covalent bond are the same, the
pair of electrons is equally shared.

The electron distribution is said to
be symmetric and the bond is said
to be non-polar.
For example, the molecules such as;
H2, Cl2, F2, Br2 and O2 are non-polar
molecules.

Polar (dipole) bonds

When a covalent bond is formed
between two atoms having
different electronegativity values,
the more electronegative atom
attracts the pair of electrons in the
bond towards itself.

The bond becomes polar or dipole.
Facts about polar bonds
(i)
The centre of the positive charge does
not coincide with the centre of the
negative charge.
(ii)
The electron distribution is asymmetric.
(iii) The two atoms are partially charged.
-
The less electronegative atom is
shown with the partial charge
(delta positive).
-
+
The more electronegative atom is
shown with the partial (delta negative).
(iv) The degree of polarity of a molecule is
measured as a dipole moment.
The direction of the dipole is
shown by the sign
-
The arrow points to the partially
negatively charged end of the
dipole.
For example, hydrogen chloride (HCl)
molecule is a polar molecule and is
shown as;
+
H
Cl
3.
Metallic Bonding

A metal consists of a regular array
of closely packed atoms.

The atoms lose their outer shell
electrons into the spaces between
the atoms and become positive
ions.

The delocalised electrons distribute
themselves throughout the entire piece
of the metal resulting to electrostatic
attraction between the positive ions and
the free moving electrons.

The electrostatic force of attraction
constitute metallic bond.

Metallic bond is strong.
Reason:

The ions are held together by the
strong electrostatic attraction between
the positive ions and the negative
charges of the delocalised electrons.
Factors which determine the strength of
metallic bonding
The strength metallic bonding increases
with;
(i)
Increasing positive charge on the metal
ions in the lattice.
(ii)
Decreasing size of the metal ions in the
lattice.
(iii) Increasing number of mobile electrons
per atom.
Bond polarity in molecules containing more
than two atoms
We take into account;
(i) The polarity of each bond.
(ii) The arrangement of the bonds in
the molecule.

Consider the trichloromethane (CHCl3)
& tetrachloromethane (CCl4) molecules.

Each has four bond pairs surrounding
the central carbon atom.

The trichloromethane (CHCl3) molecule
is a polar while the tetrachloromethane
(CCl4) is non-polar.
Why trichloromethane (CHCl3) molecule is
polar

The three C-Cl bonds are dipoles due to
the difference in electronegativity.

The three C-Cl dipoles point in a similar
direction and the combined effect of the
three dipoles is not cancelled out by the
polarity of C-H bond. Since the C-H
bond is virtually non-polar.

The electron distribution in the molecule
is asymmetric. Therefore, the molecule
is polar, with the negative end towards
the chlorine atoms.
Why tetrachloromethane (CCl4) molecule is
non-polar

Tetrachloromethane has four polar C-Cl
bonds pointing towards the four corners
of a tetrahedron.

The dipoles in each bond cancel each
other. So, tetrachloromethane is
non-polar.
(a)
H
C
Cl
Cl
(b)
Cl
+
Cl -
Trichloromethane,
a polar molecule
C
Cl
Cl
Cl
Tetrachloromethane,
a non-polar molecule
Polarity and Chemical reactivity

Bond polarity influences chemical
reactivity.
For example, both nitrogen (N≡N) and
carbon monoxide (C≡O) have triple
bonds requiring a similar amount of
energy to break them. But carbon
monoxide is more reactive than nitrogen
molecule.
Explanation:

Nitrogen molecule


Nitrogen molecule id non-polar
molecule and is fairly unreactive.
Carbon monoxide molecule

Carbon monoxide molecule is a
polar molecule.

This explains its reactivity with
oxygen and its use as a reducing
agent.
How bond polarity helps in starting a
chemical reaction

Many chemical reactions are started by
a reagent attacking one of the
electrically charged ends of a polar
molecules.
For example, Chloroethane (C2H5Cl) is
far more reactive than ethane (C2H6).
Explanation:

The Chloroethane molecule is polar.
H
H
H
C
+
Cl
C
H
H

Therefore, reagents such as OH- ions
can attack the delta positive carbon
atom of the polarised C-Cl bond.

Such attack is not possible with ethane
because the C-H bond is virtually
non-polar.
Intermolecular forces
Intermolecular forces are forces between
molecules.
Types of intermolecular forces
There are three types of intermolecular
forces, namely;
(i)
van der Waals’ forces
(ii)
Permanent dipole-dipole forces
(iii) Hydrogen bonding
The intermolecular forces are weaker than
forces within molecules due to covalent
and ionic bonding.
A table showing the relative strength of
intermolecular forces and other bonds
Bond type
Bond strength/
kJmol-1
Ionic bonding in NaCl
760
O-H covalent in H2O
464
Hydrogen bonding
20 - 50
Permanent dipole-dipole
force
5 - 20
Van der Waals’ forces
1 - 20
(i)
Van der Waals’ forces
van der Waals’ forces are very weak forces
between atoms and molecules.
Facts about van der Waals’ forces
(i)
They are due to dipoles set up between
molecules. The dipoles are as a result of
the electron charge clouds in non-polar
molecule having more of the charge
cloud on one side than the other. Such
that one dipole induces a dipole on the
neigbhouring molecules.
(ii)
Van der Waals’ forces are also called
temporary dipole-induced dipole forces
and dispersion forces .
(iii) Van der Waals’ forces increase with;
-
Increasing number of electrons
and protons in the molecule.
-
Increasing number of contact
points between molecules.
NB:

Contact points are places where the
molecules come close together.

They affect the boiling points of liquids.
Effect of increasing number of electrons on Enthalpy
of vaporisation and boiling points

Enthalpy of Vaporization is the quantity of
heat that must be absorbed if a certain
quantity of liquid is vaporized at a constant
temperature.

Both enthalpy of vaporisation and boiling
points of molecules increase with increase in
the number of electrons as shown in graphs
below.
Enthalpy of vaporisation of noble gases
against number of electrons
Enthalpy of vaporisation/
kJmol-1
20
Xe
10
Ar
Kr
Ne
He
0
0
20
40
60
Number of electrons
Boiling points of noble gases against number
of electrons
Boiling/oC
-100
Xe
Kr
Ar
-200
Ne
-270
He
0
20
40
60
Number of electrons
Effect of increasing number of contact on
boiling points
Consider the two isomers pentane,
CH3CH2CH2CH2CH3 and 2,2methylpropane,
(CH3)4C.


The two molecules have the same
number of electrons but different
boiling points (b.pt of pentane = 36oC
and that of 2,2-methylpropane = 10oC.
Explanation:

The molecules in pentane can link up
besides each other so there are a large
number of contact points. The van der
Waals’ forces are higher, so the boiling
point of pentane is higher.

The molecules in 2,2-dimethylpropane
are more compact and have smaller
number of contact points. The van der
Waals’ forces are relatively lower, so
the boiling point is lower.
This can be illustrated as shown below.
CH2
CH2
CH3
CH2
CH2
CH3
CH2
CH3
CH2
Contact point
CH3
Pentane, Boiling point 36oC
CH3
CH3
C
CH3
CH3
Contact point
CH3
Contact point
CH3
CH3
C
CH3
2,2-dimethylpropane, Boiling point 10oC
(ii)
Permanent dipole-dipole forces
Permanent dipole-dipole forces are
forces of attraction between two
molecules having permanent dipoles.
Facts about dipole-dipole forces

The molecules always have
negatively and positively charged
ends.

The attractive force between the
+ charge on one molecule and the
- on a neigbhouring molecule
causes a weak attractive force
between the molecules.

For small molecules with the same
number of electrons, permanent
dipole-dipole forces are often stronger
than van der Waals’ forces.
Effect of permanent dipole-dipole forces on
the boiling points of compounds

Consider the boiling points of propanone
(CH3COCH3, Mr = 58 and butane
(CH3CH2CH2CH3, Mr = 58).

The two molecules have the same
relative molecular mass but different
boiling points (b.pt of propanone = 56oC
and that of butane = 0oC.
CH2
CH3
CH2
CH3
Butane, boiling point 0oC
CH3
+ C=O
-
CH3
Propanone, boiling point 56oC
Explanation:

The permanent dipole-dipole forces
between propanone molecules are
strong enough to make this substance
a liquid at room temperature and has a
higher boiling point.

There are only van der Waals’ forces
between butane molecules. These
forces are comparatively weak, so
butane is a gas at room temperature
and has a lower boiling point.
(iii) Hydrogen bonding
A hydrogen bonding is the attractive force
between a hydrogen attached to an
electronegative atom of one molecule and an
electronegative atom of a different molecule.
Conditions for hydrogen bonding to occur
For hydrogen bonding to occur between two
molecules we need;
(i)
One molecule having hydrogen atom
covalently bonded to F, O or N (the three
most electronegative atoms).
(ii)
A second molecule having F, O or N
atom with a lone pair of electrons.
Therefore, the average number of hydrogen
bonds formed per molecule depends on;
(i)
The number hydrogen atoms attached
to F, O or N in the molecule.
(ii)
The number of lone pairs present on
the F, O or N atom in the molecule.
Examples of compounds in which there is
hydrogen bonding are;
Water and
Ammonia
(i)
Hydrogen bonding in water

Water has two hydrogen atoms
and two lone pairs per molecule.

So water is extensively hydrogen
bonded with other water
molecules.

It has an average of two hydrogen
bonds per molecules as shown in
the diagram below.
+

H
H
O
H
+
H
-
H
+
O
O
-
-
H
NB:

The hydrogen bonding repeats itself
continuously.
(ii)
Hydrogen bonding in Ammonia

Ammonia has three hydrogen atoms
and one lone pair per molecule.

So ammonia is less extensively
hydrogen bonded than water.

It can form on average, only one
hydrogen bond per molecules as
shown in the diagram below.
H
H
H
-
N
+
H
H
H
N
CHECK UP
Draw diagrams to show hydrogen bonding
between the following molecules;
(a)
Ethanol, C2H5OH.
(b)
Two hydrogen fluoride molecules.
Effect of hydrogen bonding on boiling point
of compounds

Hydrogen bonding makes some
compounds to have higher boiling points
than expected.
For example the graph of boiling points
of halides, HF, HCl, HBr and HI, plotted
against the position of the halogens in
the Periodic Table, shows an irregular
pattern.
The boiling points of halides against their
position in the Periodic Table
Boiling point/oC
-50
0
-50
-100
HF
HCl
HBr
HI
Explaining the shape of the graph

The high boiling point of fluorine is due
to the stronger intermolecular forces
of hydrogen bonding between the HF
molecules as a result of high
electronegativity of fluorine.

The rise in boiling point from HCl to HI
is due to the increasing number of
electrons in the halogen atoms down
the group.

This leads to increased van der Waals’
forces as the molecules get bigger.
The graph of boiling points against relative
molecular mass of other hydrides
CHEP UP
The table below shows the boiling points
of some Group V hydrides.
Hydride
(a)
(b)
Boiling point/oC
Ammonia, NH3
-33
Phosphine, PH3
-88
Arsine, AsH3
-55
Stibine, SbH3
-17
Explain the trend in the boiling points
from phosphine to stibine.
Explain why the boiling point of
ammonia does not follow this trend.

1.
Effect of hydrogen bonding on water
The intensive Hydrogen bonding in
water causes water to have peculiar
properties.
High enthalpy change of vaporisation
and boiling point.

The enthalpy change of
vaporisation and the boiling point
of water are much higher than
predicted by the trend in boiling
points for the other Group VI
hydrides. This is because water is
intensively hydrogen bonded.
Enthalpy of vaporisation/kJmol-1
The boiling points of halides against their
position in the Periodic Table
40
20
H 2O
H2Te
H 2S
H2Se
10
0
Relative molecular mass (not to scale)
2.
Surface tension and Viscosity
(i) Surface tension

Water has high surface tension
compared to other liquids.
Reason:

The hydrogen bonds in water
exert a significant downward
force at the surface of the
liquid causing the surface
tension of water to be high.
(ii)
Viscosity

Water has high viscosity
compared to other liquids.
Reason:

The hydrogen bonding reduces
the ability of water molecules
to slide over each other, so the
viscosity of water is high.
3.
Ice is less dense than water
Reason:

There is a three-dimensional
hydrogen bonded network of water
molecules. This produces a rigid
lattice in which each oxygen atom
is surrounded by a tetrahedron of
hydrogen atoms.

This ‘more open’ arrangement
allows the water molecules to be
slightly further apart than in the
liquid state. So the density of ice is
less than the density of water.
Bonding and Physical Properties

The type of bonding between atoms,
ions or molecules influences the
physical properties of a substance.
(a)
Physical state at room temperature and
pressure, Boiling point, melting point and
enthalpy change of vaporisation.
(i)
Ionic compounds

Are solid at room temperature
and pressure.
Reason:



There are strong electrostatic forces
(ionic bonds) holding the positive and
negative ions together.
The ions are regularly arranged in a
lattice, with the oppositely charged
ions close to each other.
Ionic compounds have high melting
point, high boiling point and high
enthalpy change of vaporisation.
Reason:
 It takes a lot of energy to
overcome the strong electrostatic
attractive forces.
(ii)
Metals

Metals, apart from mercury, are
solids.

Most metals have high melting
points and boiling points and high
enthalpy changes of vaporisation.
Reason:

It takes a lot of energy to
overcome the strong attractive
forces between the positive ions
and the ‘sea’ of delocalised
electrons.
(iii)
Covalent compounds

Covalently bonded substances with simple
molecular structure, for example water
and ammonia, are usually liquids or gases.

Some are solids at room temperature
and pressure.
E.g. Iodine, paraffin wax, polyethene etc.

They have low melting point, boiling point
and low enthalpy changes of vaporisation
compared with ionic compounds.
Reason:

The forces between the molecules
are weak.
(b)
Solubility
(i) Ionic compounds

Most ionic compounds are
soluble in water.
Reason:

Water molecules are polar
and they are attracted to the
ions on the surface of the
ionic solid.

These attractions, ion-dipole,
replace the electrostatic
forces between the ions and
the ions go into solution.
(ii)
Metals

Metals do not dissolve in water.
However, some metals react with
water.
E.g. Sodium, potassium & calcium.
(iii) Covalent compounds

Covalently bonded substances with
a simple molecular structure fall
into two groups.

Those that are insoluble in
water.
Most covalently bonded
molecule are non-polar.

Water molecules are not attracted to
them so they are insoluble.

Those that are soluble in water.
Small molecules that can form hydrogen
bonds with water are generally soluble.
E.g. Ethanol, C2H5OH.
Some covalently bonded substances
react with water rather than dissolving
in it.
E.g. Hydrogen chloride (HCl) reacts
with water to form hydrogen ions
and chloride ions.
(c)
Electrical conductivity
(i)
Ionic compounds

Ionic compounds do not conduct
electricity when in the solid
state.
Reason:

The ions are held strongly by the
electrostatic forces of attraction
and are not free to move.

They contact electricity in
molten and aqueous.
Reason:

The ions are free to move.
(ii)
Metals

Metals conduct electricity both
when solid and when molten.
Reason:

There are free to moving electrons.
(iii) Covalent compounds

Covalently bonded substances with
a simple molecular structure do
not conduct electricity.
Reason:

They have neither ions nor
electrons which are free to move.
END OF CHAPTER