NCERT Textbook - States of Matter : Gases & Liquids Class 11 Notes | EduRev

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Class 11 : NCERT Textbook - States of Matter : Gases & Liquids Class 11 Notes | EduRev

 Page 1


132        C:\ChemistryXI\Unit-5\Unit-5(4)-Lay-2.pmd  14.1.6 (Final), 17.1.6, 24.1.6
132 CHEMISTRY
UNIT 5
After studying this unit you will be
able to
•
explain the existence of different
states of matter in terms of
balance between intermolecular
forces and thermal energy of
particles;
•
explain the laws governing
behaviour of ideal gases;
•
apply gas laws in various real life
situations;
•
explain the behaviour of real
gases;
•
describe the conditions required
for liquifaction of gases;
•
realise that there is continuity in
gaseous and liquid state;
•
differentiate between gaseous
state and vapours;
•
explain properties of liquids in
terms of intermolecular
attractions.
STATES OF MATTER
INTRODUCTION
In previous units we have learnt about the properties
related to single particle of matter, such as atomic size,
ionization enthalpy, electronic charge density, molecular
shape and polarity, etc. Most of the observable
characteristics of chemical systems with which we are
familiar represent bulk properties of matter, i.e., the
properties associated with a collection of a large number
of atoms, ions or molecules. For example, an individual
molecule of a liquid does not boil but the bulk boils.
Collection of water molecules have wetting properties;
individual molecules do not wet. Water can exist as ice,
which is a solid; it can exist as liquid; or it can exist in
the gaseous state as water vapour or steam. Physical
properties of ice, water and steam are very different. In
all the three states of water chemical composition of water
remains the same i.e., H
2
O. Characteristics of the three
states of water depend on the energies of molecules and
on the manner in which water molecules aggregate. Same
is true for other substances also.
Chemical properties of a substance do not change with
the change of its physical state; but rate of chemical
reactions do depend upon the physical state. Many times
in calculations while dealing with data of experiments we
require knowledge of the state of matter. Therefore, it
becomes necessary for a chemist to know the physical
The snowflake falls, yet lays not long
Its feath’ry grasp on Mother Earth
Ere Sun returns it to the vapors Whence it came,
Or to waters tumbling down the rocky slope.
Rod O’ Connor
© NCERT
not to be republished
Page 2


132        C:\ChemistryXI\Unit-5\Unit-5(4)-Lay-2.pmd  14.1.6 (Final), 17.1.6, 24.1.6
132 CHEMISTRY
UNIT 5
After studying this unit you will be
able to
•
explain the existence of different
states of matter in terms of
balance between intermolecular
forces and thermal energy of
particles;
•
explain the laws governing
behaviour of ideal gases;
•
apply gas laws in various real life
situations;
•
explain the behaviour of real
gases;
•
describe the conditions required
for liquifaction of gases;
•
realise that there is continuity in
gaseous and liquid state;
•
differentiate between gaseous
state and vapours;
•
explain properties of liquids in
terms of intermolecular
attractions.
STATES OF MATTER
INTRODUCTION
In previous units we have learnt about the properties
related to single particle of matter, such as atomic size,
ionization enthalpy, electronic charge density, molecular
shape and polarity, etc. Most of the observable
characteristics of chemical systems with which we are
familiar represent bulk properties of matter, i.e., the
properties associated with a collection of a large number
of atoms, ions or molecules. For example, an individual
molecule of a liquid does not boil but the bulk boils.
Collection of water molecules have wetting properties;
individual molecules do not wet. Water can exist as ice,
which is a solid; it can exist as liquid; or it can exist in
the gaseous state as water vapour or steam. Physical
properties of ice, water and steam are very different. In
all the three states of water chemical composition of water
remains the same i.e., H
2
O. Characteristics of the three
states of water depend on the energies of molecules and
on the manner in which water molecules aggregate. Same
is true for other substances also.
Chemical properties of a substance do not change with
the change of its physical state; but rate of chemical
reactions do depend upon the physical state. Many times
in calculations while dealing with data of experiments we
require knowledge of the state of matter. Therefore, it
becomes necessary for a chemist to know the physical
The snowflake falls, yet lays not long
Its feath’ry grasp on Mother Earth
Ere Sun returns it to the vapors Whence it came,
Or to waters tumbling down the rocky slope.
Rod O’ Connor
© NCERT
not to be republished
   133    C:\ChemistryXI\Unit-5\Unit-5 (4)-Lay-2.pmd   14.1.6 (Final), 17.1.6, 24.1.6
133 STATES OF MATTER
laws which govern the behaviour of matter in
different states. In this unit, we will learn
more about these three physical states of
matter particularly liquid and gaseous states.
To begin with, it is necessary to understand
the nature of intermolecular forces, molecular
interactions and effect of thermal energy on
the  motion of particles because a balance
between  these determines the state of a
substance.
5.1 INTERMOLECULAR FORCES
Intermolecular forces are the forces of
attraction and repulsion between interacting
particles (atoms and molecules). This term
does not include the electrostatic forces that
exist between the two oppositely charged ions
and the forces that hold atoms of a molecule
together i.e., covalent bonds.
Attractive intermolecular forces are known
as van der Waals forces, in honour of Dutch
scientist Johannes van der Waals (1837-
1923), who explained the deviation of real
gases from the ideal behaviour through these
forces. We will learn about this later in this
unit. van der Waals forces vary considerably
in magnitude and include dispersion forces
or London forces, dipole-dipole forces, and
dipole-induced dipole forces. A particularly
strong type of dipole-dipole interaction is
hydrogen bonding. Only a few elements can
participate in hydrogen bond formation,
therefore it is treated as a separate
category. We have already learnt about this
interaction in Unit 4.
At this point, it is important to note that
attractive forces between an ion and a dipole
are known as ion-dipole forces and these are
not van der Waals forces. We will now learn
about different types of van der Waals forces.
5.1.1 Dispersion Forces or London Forces
Atoms and nonpolar molecules are electrically
symmetrical and have no dipole moment
because their electronic charge cloud is
symmetrically distributed. But a dipole may
develop momentarily even in such atoms and
molecules. This can be understood as follows.
Suppose we have two atoms ‘A’ and ‘B’  in the
close vicinity of each other (Fig. 5.1a). It may
so happen that momentarily electronic charge
distribution in one of the atoms, say ‘A’,
becomes unsymmetrical i.e., the charge cloud
is more on one side than the other (Fig. 5.1 b
and c). This results in the development of
instantaneous dipole on the atom ‘A’ for a very
short time. This instantaneous or transient
dipole distorts the electron density of the
other atom ‘B’, which is close to it and as a
consequence a dipole is induced in the
atom ‘B’.
The temporary dipoles of atom ‘A’ and ‘B’
attract each other. Similarly temporary dipoles
are induced in molecules also. This force of
attraction was first proposed by the German
physicist Fritz London, and for this reason
force of attraction between two temporary
Fig. 5.1 Dispersion forces or London forces
between atoms.
© NCERT
not to be republished
Page 3


132        C:\ChemistryXI\Unit-5\Unit-5(4)-Lay-2.pmd  14.1.6 (Final), 17.1.6, 24.1.6
132 CHEMISTRY
UNIT 5
After studying this unit you will be
able to
•
explain the existence of different
states of matter in terms of
balance between intermolecular
forces and thermal energy of
particles;
•
explain the laws governing
behaviour of ideal gases;
•
apply gas laws in various real life
situations;
•
explain the behaviour of real
gases;
•
describe the conditions required
for liquifaction of gases;
•
realise that there is continuity in
gaseous and liquid state;
•
differentiate between gaseous
state and vapours;
•
explain properties of liquids in
terms of intermolecular
attractions.
STATES OF MATTER
INTRODUCTION
In previous units we have learnt about the properties
related to single particle of matter, such as atomic size,
ionization enthalpy, electronic charge density, molecular
shape and polarity, etc. Most of the observable
characteristics of chemical systems with which we are
familiar represent bulk properties of matter, i.e., the
properties associated with a collection of a large number
of atoms, ions or molecules. For example, an individual
molecule of a liquid does not boil but the bulk boils.
Collection of water molecules have wetting properties;
individual molecules do not wet. Water can exist as ice,
which is a solid; it can exist as liquid; or it can exist in
the gaseous state as water vapour or steam. Physical
properties of ice, water and steam are very different. In
all the three states of water chemical composition of water
remains the same i.e., H
2
O. Characteristics of the three
states of water depend on the energies of molecules and
on the manner in which water molecules aggregate. Same
is true for other substances also.
Chemical properties of a substance do not change with
the change of its physical state; but rate of chemical
reactions do depend upon the physical state. Many times
in calculations while dealing with data of experiments we
require knowledge of the state of matter. Therefore, it
becomes necessary for a chemist to know the physical
The snowflake falls, yet lays not long
Its feath’ry grasp on Mother Earth
Ere Sun returns it to the vapors Whence it came,
Or to waters tumbling down the rocky slope.
Rod O’ Connor
© NCERT
not to be republished
   133    C:\ChemistryXI\Unit-5\Unit-5 (4)-Lay-2.pmd   14.1.6 (Final), 17.1.6, 24.1.6
133 STATES OF MATTER
laws which govern the behaviour of matter in
different states. In this unit, we will learn
more about these three physical states of
matter particularly liquid and gaseous states.
To begin with, it is necessary to understand
the nature of intermolecular forces, molecular
interactions and effect of thermal energy on
the  motion of particles because a balance
between  these determines the state of a
substance.
5.1 INTERMOLECULAR FORCES
Intermolecular forces are the forces of
attraction and repulsion between interacting
particles (atoms and molecules). This term
does not include the electrostatic forces that
exist between the two oppositely charged ions
and the forces that hold atoms of a molecule
together i.e., covalent bonds.
Attractive intermolecular forces are known
as van der Waals forces, in honour of Dutch
scientist Johannes van der Waals (1837-
1923), who explained the deviation of real
gases from the ideal behaviour through these
forces. We will learn about this later in this
unit. van der Waals forces vary considerably
in magnitude and include dispersion forces
or London forces, dipole-dipole forces, and
dipole-induced dipole forces. A particularly
strong type of dipole-dipole interaction is
hydrogen bonding. Only a few elements can
participate in hydrogen bond formation,
therefore it is treated as a separate
category. We have already learnt about this
interaction in Unit 4.
At this point, it is important to note that
attractive forces between an ion and a dipole
are known as ion-dipole forces and these are
not van der Waals forces. We will now learn
about different types of van der Waals forces.
5.1.1 Dispersion Forces or London Forces
Atoms and nonpolar molecules are electrically
symmetrical and have no dipole moment
because their electronic charge cloud is
symmetrically distributed. But a dipole may
develop momentarily even in such atoms and
molecules. This can be understood as follows.
Suppose we have two atoms ‘A’ and ‘B’  in the
close vicinity of each other (Fig. 5.1a). It may
so happen that momentarily electronic charge
distribution in one of the atoms, say ‘A’,
becomes unsymmetrical i.e., the charge cloud
is more on one side than the other (Fig. 5.1 b
and c). This results in the development of
instantaneous dipole on the atom ‘A’ for a very
short time. This instantaneous or transient
dipole distorts the electron density of the
other atom ‘B’, which is close to it and as a
consequence a dipole is induced in the
atom ‘B’.
The temporary dipoles of atom ‘A’ and ‘B’
attract each other. Similarly temporary dipoles
are induced in molecules also. This force of
attraction was first proposed by the German
physicist Fritz London, and for this reason
force of attraction between two temporary
Fig. 5.1 Dispersion forces or London forces
between atoms.
© NCERT
not to be republished
134        C:\ChemistryXI\Unit-5\Unit-5(4)-Lay-2.pmd  14.1.6 (Final), 17.1.6, 24.1.6
134 CHEMISTRY
dipoles is known as London force. Another
name for this force is dispersion force. These
forces are always attractive and interaction
energy is inversely proportional to the sixth
power of the distance between two interacting
particles (i.e., 1/r
6
 where r is the distance
between two particles). These forces are
important only at short distances (~500 pm)
and their magnitude depends on the
polarisability of the particle.
5.1.2 Dipole - Dipole Forces
Dipole-dipole forces act between the molecules
possessing permanent dipole. Ends of the
dipoles possess “partial charges” and these
charges are shown by Greek letter delta (d).
Partial charges are always less than the unit
electronic charge (1.610
–19 
C). The polar
molecules interact with neighbouring
molecules. Fig 5.2 (a) shows electron cloud
distribution in the dipole of hydrogen chloride
and Fig. 5.2 (b) shows dipole-dipole interaction
between two HCl molecules. This interaction
is stronger than the London forces but is
weaker than ion-ion interaction because only
partial charges are involved. The attractive
force decreases with the increase of distance
between the dipoles. As in the above case here
also, the interaction energy is inversely
proportional to distance between polar
molecules. Dipole-dipole interaction energy
between stationary polar molecules (as in
solids) is proportional to 1/r
3
  and that
between rotating polar molecules is
proportional to 1/r 
6
, where r is the distance
between polar molecules. Besides dipole-
dipole interaction, polar molecules can
interact by London forces also. Thus
cumulative effect is that the total of
intermolecular forces in polar molecules
increase.
5.1.3 Dipole–Induced Dipole Forces
This type of attractive forces operate between
the polar molecules having permanent dipole
and the molecules lacking permanent dipole.
Permanent dipole of the polar molecule
induces dipole on the electrically neutral
molecule by deforming its electronic cloud
(Fig. 5.3). Thus an induced dipole is developed
in the other molecule. In this case also
interaction energy is proportional to 1/r
6
where r is the distance between two
molecules. Induced dipole moment depends
upon the dipole moment present in the
permanent dipole and the polarisability of the
electrically neutral molecule. We have already
learnt in Unit 4 that molecules of larger size
can be easily polarized. High polarisability
increases the strength of attractive
interactions.
Fig. 5.2 (a) Distribution of electron cloud in HCl –
a polar molecule, (b) Dipole-dipole
interaction between two HCl molecules
Fig. 5.3 Dipole - induced dipole interaction
between permanent dipole and induced
dipole
In this case also cumulative effect of
dispersion forces and dipole-induced dipole
interactions exists.
5.1.4  Hydrogen bond
As already mentioned in section (5.1); this is
special case of dipole-dipole interaction. We
have already learnt about this in Unit 4. This
© NCERT
not to be republished
Page 4


132        C:\ChemistryXI\Unit-5\Unit-5(4)-Lay-2.pmd  14.1.6 (Final), 17.1.6, 24.1.6
132 CHEMISTRY
UNIT 5
After studying this unit you will be
able to
•
explain the existence of different
states of matter in terms of
balance between intermolecular
forces and thermal energy of
particles;
•
explain the laws governing
behaviour of ideal gases;
•
apply gas laws in various real life
situations;
•
explain the behaviour of real
gases;
•
describe the conditions required
for liquifaction of gases;
•
realise that there is continuity in
gaseous and liquid state;
•
differentiate between gaseous
state and vapours;
•
explain properties of liquids in
terms of intermolecular
attractions.
STATES OF MATTER
INTRODUCTION
In previous units we have learnt about the properties
related to single particle of matter, such as atomic size,
ionization enthalpy, electronic charge density, molecular
shape and polarity, etc. Most of the observable
characteristics of chemical systems with which we are
familiar represent bulk properties of matter, i.e., the
properties associated with a collection of a large number
of atoms, ions or molecules. For example, an individual
molecule of a liquid does not boil but the bulk boils.
Collection of water molecules have wetting properties;
individual molecules do not wet. Water can exist as ice,
which is a solid; it can exist as liquid; or it can exist in
the gaseous state as water vapour or steam. Physical
properties of ice, water and steam are very different. In
all the three states of water chemical composition of water
remains the same i.e., H
2
O. Characteristics of the three
states of water depend on the energies of molecules and
on the manner in which water molecules aggregate. Same
is true for other substances also.
Chemical properties of a substance do not change with
the change of its physical state; but rate of chemical
reactions do depend upon the physical state. Many times
in calculations while dealing with data of experiments we
require knowledge of the state of matter. Therefore, it
becomes necessary for a chemist to know the physical
The snowflake falls, yet lays not long
Its feath’ry grasp on Mother Earth
Ere Sun returns it to the vapors Whence it came,
Or to waters tumbling down the rocky slope.
Rod O’ Connor
© NCERT
not to be republished
   133    C:\ChemistryXI\Unit-5\Unit-5 (4)-Lay-2.pmd   14.1.6 (Final), 17.1.6, 24.1.6
133 STATES OF MATTER
laws which govern the behaviour of matter in
different states. In this unit, we will learn
more about these three physical states of
matter particularly liquid and gaseous states.
To begin with, it is necessary to understand
the nature of intermolecular forces, molecular
interactions and effect of thermal energy on
the  motion of particles because a balance
between  these determines the state of a
substance.
5.1 INTERMOLECULAR FORCES
Intermolecular forces are the forces of
attraction and repulsion between interacting
particles (atoms and molecules). This term
does not include the electrostatic forces that
exist between the two oppositely charged ions
and the forces that hold atoms of a molecule
together i.e., covalent bonds.
Attractive intermolecular forces are known
as van der Waals forces, in honour of Dutch
scientist Johannes van der Waals (1837-
1923), who explained the deviation of real
gases from the ideal behaviour through these
forces. We will learn about this later in this
unit. van der Waals forces vary considerably
in magnitude and include dispersion forces
or London forces, dipole-dipole forces, and
dipole-induced dipole forces. A particularly
strong type of dipole-dipole interaction is
hydrogen bonding. Only a few elements can
participate in hydrogen bond formation,
therefore it is treated as a separate
category. We have already learnt about this
interaction in Unit 4.
At this point, it is important to note that
attractive forces between an ion and a dipole
are known as ion-dipole forces and these are
not van der Waals forces. We will now learn
about different types of van der Waals forces.
5.1.1 Dispersion Forces or London Forces
Atoms and nonpolar molecules are electrically
symmetrical and have no dipole moment
because their electronic charge cloud is
symmetrically distributed. But a dipole may
develop momentarily even in such atoms and
molecules. This can be understood as follows.
Suppose we have two atoms ‘A’ and ‘B’  in the
close vicinity of each other (Fig. 5.1a). It may
so happen that momentarily electronic charge
distribution in one of the atoms, say ‘A’,
becomes unsymmetrical i.e., the charge cloud
is more on one side than the other (Fig. 5.1 b
and c). This results in the development of
instantaneous dipole on the atom ‘A’ for a very
short time. This instantaneous or transient
dipole distorts the electron density of the
other atom ‘B’, which is close to it and as a
consequence a dipole is induced in the
atom ‘B’.
The temporary dipoles of atom ‘A’ and ‘B’
attract each other. Similarly temporary dipoles
are induced in molecules also. This force of
attraction was first proposed by the German
physicist Fritz London, and for this reason
force of attraction between two temporary
Fig. 5.1 Dispersion forces or London forces
between atoms.
© NCERT
not to be republished
134        C:\ChemistryXI\Unit-5\Unit-5(4)-Lay-2.pmd  14.1.6 (Final), 17.1.6, 24.1.6
134 CHEMISTRY
dipoles is known as London force. Another
name for this force is dispersion force. These
forces are always attractive and interaction
energy is inversely proportional to the sixth
power of the distance between two interacting
particles (i.e., 1/r
6
 where r is the distance
between two particles). These forces are
important only at short distances (~500 pm)
and their magnitude depends on the
polarisability of the particle.
5.1.2 Dipole - Dipole Forces
Dipole-dipole forces act between the molecules
possessing permanent dipole. Ends of the
dipoles possess “partial charges” and these
charges are shown by Greek letter delta (d).
Partial charges are always less than the unit
electronic charge (1.610
–19 
C). The polar
molecules interact with neighbouring
molecules. Fig 5.2 (a) shows electron cloud
distribution in the dipole of hydrogen chloride
and Fig. 5.2 (b) shows dipole-dipole interaction
between two HCl molecules. This interaction
is stronger than the London forces but is
weaker than ion-ion interaction because only
partial charges are involved. The attractive
force decreases with the increase of distance
between the dipoles. As in the above case here
also, the interaction energy is inversely
proportional to distance between polar
molecules. Dipole-dipole interaction energy
between stationary polar molecules (as in
solids) is proportional to 1/r
3
  and that
between rotating polar molecules is
proportional to 1/r 
6
, where r is the distance
between polar molecules. Besides dipole-
dipole interaction, polar molecules can
interact by London forces also. Thus
cumulative effect is that the total of
intermolecular forces in polar molecules
increase.
5.1.3 Dipole–Induced Dipole Forces
This type of attractive forces operate between
the polar molecules having permanent dipole
and the molecules lacking permanent dipole.
Permanent dipole of the polar molecule
induces dipole on the electrically neutral
molecule by deforming its electronic cloud
(Fig. 5.3). Thus an induced dipole is developed
in the other molecule. In this case also
interaction energy is proportional to 1/r
6
where r is the distance between two
molecules. Induced dipole moment depends
upon the dipole moment present in the
permanent dipole and the polarisability of the
electrically neutral molecule. We have already
learnt in Unit 4 that molecules of larger size
can be easily polarized. High polarisability
increases the strength of attractive
interactions.
Fig. 5.2 (a) Distribution of electron cloud in HCl –
a polar molecule, (b) Dipole-dipole
interaction between two HCl molecules
Fig. 5.3 Dipole - induced dipole interaction
between permanent dipole and induced
dipole
In this case also cumulative effect of
dispersion forces and dipole-induced dipole
interactions exists.
5.1.4  Hydrogen bond
As already mentioned in section (5.1); this is
special case of dipole-dipole interaction. We
have already learnt about this in Unit 4. This
© NCERT
not to be republished
   135    C:\ChemistryXI\Unit-5\Unit-5 (4)-Lay-2.pmd   14.1.6 (Final), 17.1.6, 24.1.6
135 STATES OF MATTER
is found in the molecules in which highly polar
N–H, O–H or H–F bonds are present.  Although
hydrogen bonding is regarded as being limited
to N, O and F; but species such as Cl
 
may
also participate in hydrogen bonding. Energy
of hydrogen bond varies between 10 to 100
kJ mol
–1
. This is quite a significant amount of
energy; therefore, hydrogen bonds are
powerful force in determining the structure and
properties of many compounds, for example
proteins and nucleic acids. Strength of the
hydrogen bond is determined by the coulombic
interaction between the lone-pair electrons of
the electronegative atom of one molecule and
the hydrogen atom of other molecule.
Following diagram shows the formation of
hydrogen bond.
HF HF
d + d- d+ d -
-··· -
Intermolecular forces discussed so far are
all attractive. Molecules also exert repulsive
forces on one another. When two molecules
are brought into close contact with each other,
the repulsion between the electron clouds and
that between the nuclei of two molecules comes
into play. Magnitude of the repulsion rises very
rapidly as the distance separating the
molecules decreases. This is the reason that
liquids and solids are hard to compress. In
these states molecules are already in close
contact; therefore they resist further
compression; as that would result in the
increase of repulsive interactions.
5.2 THERMAL ENERGY
Thermal energy is the energy of a body arising
from motion of its atoms or molecules. It is
directly proportional to the temperature of the
substance. It is the measure of average kinetic
energy of the particles of the matter and is
thus responsible for movement of particles.
This movement of particles is called thermal
motion.
5.3 INTERMOLECULAR FORCES vs
THERMAL INTERACTIONS
We have already learnt that intermolecular
forces tend to keep the molecules together but
thermal energy of the molecules tends to keep
them apart. Three states of matter are the result
of balance between intermolecular forces and
the thermal energy of the molecules.
When molecular interactions are very
weak, molecules do not cling together to make
liquid or solid unless thermal energy is
reduced by lowering the temperature. Gases
do not liquify on compression only, although
molecules come very close to each other and
intermolecular forces operate to the maximum.
However, when thermal energy of molecules
is reduced by lowering the temperature; the
gases can be very easily liquified.
Predominance of thermal energy and the
molecular interaction energy of a substance
in three states is depicted as follows :
We have already learnt the cause for the
existence of the three states of matter. Now
we will learn more about gaseous and liquid
states and the laws which govern the
behaviour of matter in these states. We shall
deal with the solid state in class XII.
5.4 THE GASEOUS STATE
This is the simplest state of matter.
Throughout our life we remain immersed in
the ocean of air which is a mixture of gases.
We spend our life in the lowermost layer of
the atmosphere  called troposphere, which is
held to the surface of the earth by gravitational
force. The thin layer of atmosphere is vital to
our life. It shields us from harmful radiations
and contains substances like dioxygen,
dinitrogen, carbon dioxide, water vapour, etc.
Let us now focus our attention on the
behaviour of substances which exist in the
gaseous state under normal conditions of
temperature and pressure. A look at the
periodic table shows that only eleven elements
© NCERT
not to be republished
Page 5


132        C:\ChemistryXI\Unit-5\Unit-5(4)-Lay-2.pmd  14.1.6 (Final), 17.1.6, 24.1.6
132 CHEMISTRY
UNIT 5
After studying this unit you will be
able to
•
explain the existence of different
states of matter in terms of
balance between intermolecular
forces and thermal energy of
particles;
•
explain the laws governing
behaviour of ideal gases;
•
apply gas laws in various real life
situations;
•
explain the behaviour of real
gases;
•
describe the conditions required
for liquifaction of gases;
•
realise that there is continuity in
gaseous and liquid state;
•
differentiate between gaseous
state and vapours;
•
explain properties of liquids in
terms of intermolecular
attractions.
STATES OF MATTER
INTRODUCTION
In previous units we have learnt about the properties
related to single particle of matter, such as atomic size,
ionization enthalpy, electronic charge density, molecular
shape and polarity, etc. Most of the observable
characteristics of chemical systems with which we are
familiar represent bulk properties of matter, i.e., the
properties associated with a collection of a large number
of atoms, ions or molecules. For example, an individual
molecule of a liquid does not boil but the bulk boils.
Collection of water molecules have wetting properties;
individual molecules do not wet. Water can exist as ice,
which is a solid; it can exist as liquid; or it can exist in
the gaseous state as water vapour or steam. Physical
properties of ice, water and steam are very different. In
all the three states of water chemical composition of water
remains the same i.e., H
2
O. Characteristics of the three
states of water depend on the energies of molecules and
on the manner in which water molecules aggregate. Same
is true for other substances also.
Chemical properties of a substance do not change with
the change of its physical state; but rate of chemical
reactions do depend upon the physical state. Many times
in calculations while dealing with data of experiments we
require knowledge of the state of matter. Therefore, it
becomes necessary for a chemist to know the physical
The snowflake falls, yet lays not long
Its feath’ry grasp on Mother Earth
Ere Sun returns it to the vapors Whence it came,
Or to waters tumbling down the rocky slope.
Rod O’ Connor
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133 STATES OF MATTER
laws which govern the behaviour of matter in
different states. In this unit, we will learn
more about these three physical states of
matter particularly liquid and gaseous states.
To begin with, it is necessary to understand
the nature of intermolecular forces, molecular
interactions and effect of thermal energy on
the  motion of particles because a balance
between  these determines the state of a
substance.
5.1 INTERMOLECULAR FORCES
Intermolecular forces are the forces of
attraction and repulsion between interacting
particles (atoms and molecules). This term
does not include the electrostatic forces that
exist between the two oppositely charged ions
and the forces that hold atoms of a molecule
together i.e., covalent bonds.
Attractive intermolecular forces are known
as van der Waals forces, in honour of Dutch
scientist Johannes van der Waals (1837-
1923), who explained the deviation of real
gases from the ideal behaviour through these
forces. We will learn about this later in this
unit. van der Waals forces vary considerably
in magnitude and include dispersion forces
or London forces, dipole-dipole forces, and
dipole-induced dipole forces. A particularly
strong type of dipole-dipole interaction is
hydrogen bonding. Only a few elements can
participate in hydrogen bond formation,
therefore it is treated as a separate
category. We have already learnt about this
interaction in Unit 4.
At this point, it is important to note that
attractive forces between an ion and a dipole
are known as ion-dipole forces and these are
not van der Waals forces. We will now learn
about different types of van der Waals forces.
5.1.1 Dispersion Forces or London Forces
Atoms and nonpolar molecules are electrically
symmetrical and have no dipole moment
because their electronic charge cloud is
symmetrically distributed. But a dipole may
develop momentarily even in such atoms and
molecules. This can be understood as follows.
Suppose we have two atoms ‘A’ and ‘B’  in the
close vicinity of each other (Fig. 5.1a). It may
so happen that momentarily electronic charge
distribution in one of the atoms, say ‘A’,
becomes unsymmetrical i.e., the charge cloud
is more on one side than the other (Fig. 5.1 b
and c). This results in the development of
instantaneous dipole on the atom ‘A’ for a very
short time. This instantaneous or transient
dipole distorts the electron density of the
other atom ‘B’, which is close to it and as a
consequence a dipole is induced in the
atom ‘B’.
The temporary dipoles of atom ‘A’ and ‘B’
attract each other. Similarly temporary dipoles
are induced in molecules also. This force of
attraction was first proposed by the German
physicist Fritz London, and for this reason
force of attraction between two temporary
Fig. 5.1 Dispersion forces or London forces
between atoms.
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134        C:\ChemistryXI\Unit-5\Unit-5(4)-Lay-2.pmd  14.1.6 (Final), 17.1.6, 24.1.6
134 CHEMISTRY
dipoles is known as London force. Another
name for this force is dispersion force. These
forces are always attractive and interaction
energy is inversely proportional to the sixth
power of the distance between two interacting
particles (i.e., 1/r
6
 where r is the distance
between two particles). These forces are
important only at short distances (~500 pm)
and their magnitude depends on the
polarisability of the particle.
5.1.2 Dipole - Dipole Forces
Dipole-dipole forces act between the molecules
possessing permanent dipole. Ends of the
dipoles possess “partial charges” and these
charges are shown by Greek letter delta (d).
Partial charges are always less than the unit
electronic charge (1.610
–19 
C). The polar
molecules interact with neighbouring
molecules. Fig 5.2 (a) shows electron cloud
distribution in the dipole of hydrogen chloride
and Fig. 5.2 (b) shows dipole-dipole interaction
between two HCl molecules. This interaction
is stronger than the London forces but is
weaker than ion-ion interaction because only
partial charges are involved. The attractive
force decreases with the increase of distance
between the dipoles. As in the above case here
also, the interaction energy is inversely
proportional to distance between polar
molecules. Dipole-dipole interaction energy
between stationary polar molecules (as in
solids) is proportional to 1/r
3
  and that
between rotating polar molecules is
proportional to 1/r 
6
, where r is the distance
between polar molecules. Besides dipole-
dipole interaction, polar molecules can
interact by London forces also. Thus
cumulative effect is that the total of
intermolecular forces in polar molecules
increase.
5.1.3 Dipole–Induced Dipole Forces
This type of attractive forces operate between
the polar molecules having permanent dipole
and the molecules lacking permanent dipole.
Permanent dipole of the polar molecule
induces dipole on the electrically neutral
molecule by deforming its electronic cloud
(Fig. 5.3). Thus an induced dipole is developed
in the other molecule. In this case also
interaction energy is proportional to 1/r
6
where r is the distance between two
molecules. Induced dipole moment depends
upon the dipole moment present in the
permanent dipole and the polarisability of the
electrically neutral molecule. We have already
learnt in Unit 4 that molecules of larger size
can be easily polarized. High polarisability
increases the strength of attractive
interactions.
Fig. 5.2 (a) Distribution of electron cloud in HCl –
a polar molecule, (b) Dipole-dipole
interaction between two HCl molecules
Fig. 5.3 Dipole - induced dipole interaction
between permanent dipole and induced
dipole
In this case also cumulative effect of
dispersion forces and dipole-induced dipole
interactions exists.
5.1.4  Hydrogen bond
As already mentioned in section (5.1); this is
special case of dipole-dipole interaction. We
have already learnt about this in Unit 4. This
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   135    C:\ChemistryXI\Unit-5\Unit-5 (4)-Lay-2.pmd   14.1.6 (Final), 17.1.6, 24.1.6
135 STATES OF MATTER
is found in the molecules in which highly polar
N–H, O–H or H–F bonds are present.  Although
hydrogen bonding is regarded as being limited
to N, O and F; but species such as Cl
 
may
also participate in hydrogen bonding. Energy
of hydrogen bond varies between 10 to 100
kJ mol
–1
. This is quite a significant amount of
energy; therefore, hydrogen bonds are
powerful force in determining the structure and
properties of many compounds, for example
proteins and nucleic acids. Strength of the
hydrogen bond is determined by the coulombic
interaction between the lone-pair electrons of
the electronegative atom of one molecule and
the hydrogen atom of other molecule.
Following diagram shows the formation of
hydrogen bond.
HF HF
d + d- d+ d -
-··· -
Intermolecular forces discussed so far are
all attractive. Molecules also exert repulsive
forces on one another. When two molecules
are brought into close contact with each other,
the repulsion between the electron clouds and
that between the nuclei of two molecules comes
into play. Magnitude of the repulsion rises very
rapidly as the distance separating the
molecules decreases. This is the reason that
liquids and solids are hard to compress. In
these states molecules are already in close
contact; therefore they resist further
compression; as that would result in the
increase of repulsive interactions.
5.2 THERMAL ENERGY
Thermal energy is the energy of a body arising
from motion of its atoms or molecules. It is
directly proportional to the temperature of the
substance. It is the measure of average kinetic
energy of the particles of the matter and is
thus responsible for movement of particles.
This movement of particles is called thermal
motion.
5.3 INTERMOLECULAR FORCES vs
THERMAL INTERACTIONS
We have already learnt that intermolecular
forces tend to keep the molecules together but
thermal energy of the molecules tends to keep
them apart. Three states of matter are the result
of balance between intermolecular forces and
the thermal energy of the molecules.
When molecular interactions are very
weak, molecules do not cling together to make
liquid or solid unless thermal energy is
reduced by lowering the temperature. Gases
do not liquify on compression only, although
molecules come very close to each other and
intermolecular forces operate to the maximum.
However, when thermal energy of molecules
is reduced by lowering the temperature; the
gases can be very easily liquified.
Predominance of thermal energy and the
molecular interaction energy of a substance
in three states is depicted as follows :
We have already learnt the cause for the
existence of the three states of matter. Now
we will learn more about gaseous and liquid
states and the laws which govern the
behaviour of matter in these states. We shall
deal with the solid state in class XII.
5.4 THE GASEOUS STATE
This is the simplest state of matter.
Throughout our life we remain immersed in
the ocean of air which is a mixture of gases.
We spend our life in the lowermost layer of
the atmosphere  called troposphere, which is
held to the surface of the earth by gravitational
force. The thin layer of atmosphere is vital to
our life. It shields us from harmful radiations
and contains substances like dioxygen,
dinitrogen, carbon dioxide, water vapour, etc.
Let us now focus our attention on the
behaviour of substances which exist in the
gaseous state under normal conditions of
temperature and pressure. A look at the
periodic table shows that only eleven elements
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136        C:\ChemistryXI\Unit-5\Unit-5(4)-Lay-2.pmd  14.1.6 (Final), 17.1.6, 24.1.6
136 CHEMISTRY
exist as gases under normal conditions
(Fig 5.4).
The gaseous state is characterized by the
following physical properties.
• Gases are highly compressible.
• Gases exert pressure equally in all
directions.
• Gases have much lower density than the
solids and liquids.
• The volume and the shape of  gases are
not fixed. These assume volume and shape
of the container.
• Gases mix evenly and completely in all
proportions without any mechanical aid.
Simplicity of gases is due to the fact that
the forces of interaction between their
molecules  are negligible. Their behaviour is
governed by same general laws, which were
discovered as a result of their experimental
studies. These laws are relationships between
measurable properties of gases. Some of  these
properties like pressure, volume, temperature
and mass are very important because
relationships between these variables describe
state of the gas. Interdependence of these
variables leads to the formulation of gas laws.
In the next section we will learn about gas
laws.
5.5  THE GAS LAWS
The gas laws which we will study now are the
result of research carried on for several
centuries on the physical properties of gases.
The first reliable measurement on properties
of gases was made by Anglo-Irish scientist
Robert Boyle in 1662. The law which he
formulated is known as Boyle’s Law. Later
on attempts to fly in air with the help of hot
air balloons motivated Jaccques Charles and
Joseph Lewis Gay Lussac to discover
additional gas laws. Contribution from
Avogadro and others provided lot of
information about gaseous state.
5.5.1 Boyle’s Law (Pressure - Volume
Relationship)
On the basis of his experiments, Robert Boyle
reached to the conclusion that at constant
temperature, the pressure of a fixed
amount  (i.e., number of moles n) of gas
varies inversely with its volume. This is
known as Boyle’s law. Mathematically, it can
be written as
1
  p
V
?
  ( at constant T and n)             (5.1)
1
1
 =k ? p
V
                 (5.2)
where k
1
 is the proportionality constant. The
value of constant k
1
 depends upon the
amount of the gas, temperature of the gas
and the units in which p and V are expressed.
On rearranging equation (5.2) we obtain
pV = k
1
   (5.3)
It means that at constant temperature,
product of pressure and volume of a fixed
amount of gas is constant.
If a fixed amount of gas at constant
temperature T occupying volume V
1
 at
pressure p
1
 undergoes expansion, so that
volume becomes V
2
 and pressure becomes p
2
,
then according to Boyle’s law :
 p
1
V
1
 = p
2
V
2
= constant   (5.4)
12
21
= ?
p V
p V
  (5.5)
Fig. 5.4  Eleven elements that exist as gases
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not to be republished
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