CHAPTER 5 - STATES OF MATTER

Vaporisation Process

• Vaporization is a process of changing the state from liquid
to vapor.

• molecules in a liquid moves freely, collide with each other
and posses different magnitudes of kinetic energy

• some molecules have a relatively high kinetic energies

• molecules at the surface possess sufficient enough kinetic
energy to overcome the intermolecular attractive forces that
bind them

• when the kinetic energy is sufficient enough to overcome
the intermolecular forces acting on them, the molecules
will escape as vapour.

101

Factors affecting the rate of vaporisation :

i) Surface area

• the larger the surface area, the higher the chances for the
molecules to escape from the surface

 surface area increases, evaporation rate increases evaporation –

the change state of a liquid into a vapour at a temperature below bp of the liquid, occurs at the
surface of liquid

ii) Temperature

• as temperature is increased, the total number of molecules
with high kinetic energy is increased

• more molecules have enough energy to escape from the
surface of the liquid

• thus evaporation rate increases 102

iii) Intermolecular attractive forces
• the weaker the intermolecular attractive forces, the faster

the evaporation rate

103

Condensation Process
• a process whereby vapour molecules turn to liquid
• some of the vapour molecules may lose their kinetic energy

during the collision
• they do not have enough energy to remain as vapour

molecules
• they reached the surface of the liquid and trapped by the

attractive forces
• if they cannot overcome the attractive forces, these vapour

molecules return as liquid molecules
• the process is known as condensation

104

Vapour Pressure

Definition:
Vapour pressure is the pressure exerted by a vapour in
equilibrium with its liquid phase.

• Molecules can escape from the surface of liquid at any
temperature by evaporation

• in an open system ,vapour molecules which evaporate off
will diffuse away

105

• In a closed system, vapour molecules which leaves the
surface are trapped in the close container

• These vapour molecules are in constant random motion

• The molecules strike the wall of container and exert some
pressure

• As the quantity of molecules in the vapour phase increase,
some molecules may lose energy and condense

• Eventually, the rate of evaporation = the rate of
condensation (The system achieved dynamic equilibrium)

• At equilibrium, the number of vapour molecules above
liquid are constant

• The pressure exerted by those molecules is called vapour

pressure (or maximum vapour pressure) 106

Boiling – the process

• Increasing the temperature will increase in the vapour
pressure.

• As heat is applied, the vapour pressure of a system will
increase until it reaches a point whereby the vapour
pressure of the liquid system is equal to the atmospheric
pressure.

• Boiling occurs and the temperature taken at this point is
known as the boiling point.

• At this point, the change of state from liquid to gas occurs
not only at the surface of the liquid but also in the inner part
of the liquid.

• Bubbles form within the liquid. 107

Definition:

Boiling Point : the temperature at which the vapour
pressure of a liquid is equal to the external
atmospheric pressure.

Normal Boiling Point: the temperature at which liquid boils
when the external pressure is 1 atm.

(that is the vapour pressure is 760 mmHg)

108

The Relationship Between Intermolecular
Forces and Vapour Pressure

• Molecules with weak intermolecular forces need to have
less kinetic energy to escape from the liquid surface.

• Less energy is needed to overcome the intermolecular
forces between liquid molecules.

• Thus, a liquid can easily vapourise.
• More vapour molecules will be present and exert higher

pressure.
• Vapour pressure increases.

Intermolecular forces  , Vapour pressure

109

The Relationship Between
Vapour Pressure and Boiling Point

• A liquid boils when its vapour pressure equals to atmospheric
pressure.

• The temperature in which the vapour pressure of its liquid equal to
atmospheric pressure is called boiling point.

• Liquids with weaker intermolecular forces has higher vapour
pressure.

• Thus, a molecule with weak intermolecular forces can easily
vapourise and the system requires less heat to achieve
atmospheric pressure, therefore it boils at a lower temperature.

Intermolecular forces  , Vapour pressure, Boiling point 

110

evaporation boiling

 The changes  The changes of liquid into
vapour state, occurs at the inner
of a liquid into a part (body) of the liquid at a
certain temperature (when the
vapour state vapour pressure is equal to the
atmospheric pressure)
occuring at the
 The temperature is called the
surface of a boiling point

liquid at any 111

temperatures

and pressures











LEARNING OUTCOMES

At the end of the lesson, students should be able to :

• Explain the fixed shape of solids

• Apply the kinetic concept of the following process :
(i) freezing; (ii) melting (fusion); (iii) sublimation; (iv)
deposition

• Differentiate between amorphous and crystalline
solids

• State the following types of crystalline solids with

appropriate examples.

i. metallic iii. Molecular covalent

ii. Ionic iv. Giant covalent 117

5.3 Solid

Fixed shape of a solid

• Intermolecular forces between solid molecules are strong
enough to hold molecules together and lock them in place

• There is very little free space between the molecules
• Particles are closely arranged and regularly in order
• Rigid arrangement- particles can vibrate, rotate about fixed

position and cannot move freely without disrupting the
whole structure.

118

In principle, solid, liquid and gas states are
interconvertible

solid liquid gas
sublimation
deposition

5

119

Freezing (solidification)

• Occurs when a liquid changes to solid
• Freezing process: when the temperature of a liquid is

lowered:
- the kinetic energy of the liquid molecules decrease
- the molecules move slowly
- the intermolecular attractions overcome the motion of
the molecules
- when the intermolecular attraction are strong enough

to hold the particles together in a fixed and orderly
arrangement, the liquid freezes.
Freezing point is a temperature at which the liquid and
solid phases of a substances coexist at equilibrium. 120

Melting (fusion)

• A process at which a solid changes into a liquid
• The melting process: when a solid substance is heated

- kinetic energy of the molecules increase
- the molecules vibrate and rotate more rapidly
- at certain temperature, the kinetic energy is higher

enough to overcome the intermolecular forces of
attraction between solid particles.
- the particles are free to move and the solid starts to melt

Melting point is the temperature at which solid and liquid
coexist in equilibrium

121

Sublimation
• The process by which a substance changes directly from

solid to the gaseous state without passing through the
liquid state.
• Occurs on solid with weak intermolecular forces of
attraction.

Deposition
• The process where molecules from gaseous state change

to the solid state.
• The opposite process of sublimation

122

TYPES OF SOLID

CRYSTALLINE SOLID AMORPHOUS SOLID

• A solid that has highly • Solid that does not have a
ordered structure where regular three dimensional
atoms, ions or molecules arrangement of atoms or
show a regular repetition in molecules
three dimensional • Formed when a saturated
arrangement liquid is cooled rapidly
• Formed when a saturated • Example : glass,
liquid is cooled slowly.
• Its atoms, molecules or ions plastic material
occupy specific position.
• Example : ice, sugar, salt 123

Types of crystalline solid

i. Metallic crystal

• Metallic crystals are formed by metal

• These crystals are made up of atoms of the same element,
thus they can pack very closely in 3-dimensional structure

• Composed of atoms of the same metal linked together by
metallic bonds

• The physical properties of metal:

- high electrical and thermal conductivity

- lustre

- ductile and malleable

• Examples: all metallic elements – Na, Mg, Fe

124

ii. Ionic crystal

• Ionic crystals are composed of positive and negative ions
which are held together by ionic bonds

• Physical properties of ionic crystal:
- high melting point
- hard but brittle
- does not conduct electricity in the solid state but does
conduct electricity in molten or in aqueous state.

• Example : NaCl

125

iii. Molecular covalent crystals
• Composed of molecules held together by intermolecular

forces(van der Waals or hydrogen bonds)
• Example : iodin, I2

iv. Giant covalent solids
• Very large molecules / gigantic structure
• Composed of atoms linked together in large networks by

covalent bonds
• Example : diamond, graphite, SiO2

126

Diamond
• In diamond, each carbon atom bonded to four other

carbon atoms through strong covalent bonds in a
tetrahedral arrangement.
• This arrangement is repeated to give a three-dimensional
giant structure
• The physical properties of diamond:
- transparency
- great hardness
- non conductor of electricity
- high melting point (about 3500oC)

127

128

Graphite

• Graphite is an example of layered structure in the
hexagonal crystalline system.

• In the layers, each carbon atom is covalently bonded to
three other carbon atoms to form a hexagonal ring.

• The layers are held together by weak Van der Waals
forces.

• The physical properties of graphite:

- gray in colour

- metallic lustre

- soft

- conductor of electricity and heat 129

130







5.4 PHASE DIAGRAM

134

LEARNING OUTCOMES

At the end of this lesson students should be able to:

(a) Define phase, triple point and critical point

(b) Sketch and differentiate the phase diagram of H2O and
CO2

(c) Identify triple point and critical point on the phase diagram.

(d) Explain the anomalous behaviour of H2O
(e) Describe the changes in phase with respect to

i. temperature (at constant pressure)

ii. pressure (at constant temperature) 135

• Phase is a homogeneous part of a system in contact with
other parts of the system but separated from them by a
well-defined boundary

• Phase Diagram is a diagram that describes the stable
phases and phase changes of a substance as a function of
temperature and pressure

• Phase diagram is used to predict the phase that exist under
a certain temperature and pressure

136

Phase diagram of carbon dioxide, CO2
BC

T
A

137

Regions of the diagram
- The diagram has three main regions :

Solid, liquid and gas
- Each region corresponds to one phase of the substance.
- CO2 is a gas under normal conditions of temperature and

pressure ( 25oC and 1 atm)

Lines between regions
Any point along a line shows the pressure and temperature at

which the two phases exist in equilibrium.

138

TA curve:

• at the temperatures and pressures along TA curve, solid and
vapor are at equilibrium (sublimation or deposition)

CO2(s) ⇌ CO2(g)
• TA curve represents the vapor pressure of solid CO2 as it

sublimes at different temperature

• sublimation is a process of a solid changing directly into a gas

• Solid CO2 does not melt, but sublimes when heated below 5.2
atm

• At 1 atm, CO2 sublimes at -78oC
• CO2 never exists as liquid under normal atmospheric pressure
• reverse process of sublimation is deposition

139

TB curve:

• at the temperatures and pressures along TB curve, solid
and liquid are at equilibrium (melting or freezing)

CO2(s) ⇌ CO2(l)
• Known as the melting point or freezing point curve

• TB curve represents the change in melting point of CO2
with increasing pressure

• TB curve slant to the right because solid is denser than
liquid. This is normal characteristic of most substances

• Melting point of CO2 increases with pressure.
• Solid is harder to melt at higher pressure

higher melting point at higher pressure

140

TC Curve:
• at the temperatures and pressures along TC curve, liquid

and vapour are at equilibrium
CO2(l) ⇌ CO2(g)

• Known as the boiling point curve
• TC curve shows that the boiling point of this liquid

increases with increase in pressure

141

• The vapor pressure curve ends at critical point, C
• Critical point is the temperature and pressure at which

the liquid and its vapor become identical and
indistinguishable
• Critical point for CO2 is at 73 atm pressure and temperature
31oC
• At the temperature above 31oC the CO2 gas cannot be
condensed no matter how high the applied pressure is

142

• Point T is known as the triple point
• Triple point is the temperature and pressure at which

the vapour, liquid and solid states of a substance are in
equilibrium.

CO2(s) ⇌ CO2(l) ⇌ CO2(g)
• Triple point for carbon dioxide is -57°C and 5.2 atm.

143

Phase diagram of water C

B

T

A

144

• Water is a liquid under normal conditions of temperature
and pressure ( 25oC and 1 atm )

• At point T, the triple point; ice, water and steam are in
equilibrium
H2O(s) ⇌ H2O(l) ⇌ H2O(g)

• Ttermippleerpaotiunrtefo0r.0H12oOCis at 0.006 atm pressure and

• At the temperatures and pressures along the curves, 2
phases are in equilibrium.

145

TA curve

• at the temperatures and pressures along TA curve, ice and
steam are at equilibrium

H2O(s) ⇌ H2O(g)
TB curve

• at the temperatures and pressures along TB curve, ice and
water are at equilibrium

H2O(s) ⇌ H2O(l)

• TB curve represents the change in melting point or freezing
point of water with increasing pressure

• melting point of a substance is identical to its freezing point

• melting point of a liquid at 1 atm is the normal melting point

• Normal melting point for water is 0oC 146

Anomalous behaviour of water
• for H2O, the TB curve (melting point or freezing point curve)

slant to the left, that is, the melting point decreases with
pressure.
• This is because, water is denser than ice
• Ice is easier to melt at higher pressure
 lower melting point at higher pressure

147

TC Curve
• at the temperatures and pressures along TC curve, water

and steam are at equilibrium
H2O(l) ⇌ H2O(g)

• TC curve also represent the boiling point of water at
different pressure

• boiling point of a liquid when its vapour pressure is 1 atm is
the normal boiling point

• Normal boiling point for water is 100oC

148

• Critical point for water is at 218 atm and temperature 374oC
• beyond the critical point, water and steam are

indistinguishable
• At the temperature above 374oC steam cannot be

condensed no matter how high the applied pressure

149

Exercise:
1. Describe the phase change when carbon dioxide

undergoes isobaric heating at 5.2 atm pressure.
2. Describe the phase change when pressure is applied to

water isothermally at 0.01oC

150