The document discusses states of matter and changes between different states. It provides information on: - The three main states of matter - solid, liquid, gas - and how a change of state involves a change in physical form but not chemical identity. - Energy must be gained or lost for a substance to change states, altering temperature or motion of particles. - Specific changes include melting (solid to liquid), freezing (liquid to solid), evaporation/boiling (liquid to gas), condensation (gas to liquid), and sublimation (solid to gas). - Phase diagrams illustrate conditions where states coexist and phase boundaries. Eutectic mixtures have a lower melting point than their components.

1. STATES OF MATTER
CHANGES OF PHASES
 CHANGES IN THE STATES OF MATTER
 LATENT HEATS
 VAPOUR PRESSURE
 CRITICAL POINT
 EUTECTIC MIXTURES
Nabeela Moosakutty
Lecturer
Dept of Pharmaceutics
K.T.N College of Pharmacy

2. THE FACT OF THE MATTER
What happens when matter changes state?
• The three most familiar states of matter are solid, liquid, and gas.
• A change of state is the change of a substance from one
physical form of matter to another.
• When a substance undergoes a physical change, it doesnot change
its identity, just its appearance.
• change a substance from one state to another, energy must be
added or removed.
• When a substance gains or loses energy, its temperature
changes or its state changes.
• All matter is made of tiny particles that are in constant motion.
During a change of state, the motion of the particles changes.
• Particles can break away from each other and gain more freedom to
move, or they may attract each other more strongly and have less
freedom to move.
• During a change of state, a substance gains energy from or loses
energy to the environment, but the total amount of energy is
conserved.

4. Melting
• Change from a solid to a liquid
• The melting point of a substance is fixed in pure states.
• Impure substances have a variable melting point.
• When heat is applied, the particles are pushed apart.
• When a solid is warmed, its particles gain energy and speed up, and
the attraction between them decreases. Eventually they slide past
one another.
• The temperature at which a substance changes from a solid to a liquid is
called its
melting point.
• Example: When ice turns to water.
The temperature remains at 0 until all of the solid substance is melted.
CHANGING STATES
Melting, Freezing,
Evaporation & boiling,
Condensation and
Sublimation

5. Freezing
• The change from a liquid to a solid.
• When a liquid is cooled, its particles have less energy, they
slow down, and they lock into the fixed arrangement of a
solid.
• The temperature at which when a liquid freezes is called
the “freezing point”
• For most substances, the “freezing point” is the same as
the “melting point”.

6. EVAPORATION & BOILING
• Evaporation is the process in which particles of a liquid leave the
surface as a vapour.
• When a liquid is heated, the particles move faster and collide with each
other. When the particles acquire sufficient energy, they break free of the
surface and escape
The rate of evaporation depends on:
• The nature of the liquid
• The temperature
• The amount of exposed surface
• Boiling is the process by which a liquid is freely converted to a gas or
vapour.
• When a substance is heated, the temperature rises until it reaches the
boiling point.
• When boiling has started, the temperature remains steady.
• Evaporation takes place at the surface, but boiling takes place
throughout the liquid.

8. Condensation
• The change from a gas to a liquid.
• The particles lose energy and move
closer together as the gas cools therefore
increasing the attraction between them.
Sublimation
When a substance changes directly from a
solid to a gas without going through the
liquid state
examples of substances that sublime are
iodine and ammonium chloride
Deposition
It is the change in statefrom a
gas directly to a solid
Rate of Diffusion
• Solids -> Very Slow
• Liquids -> Slow
• Gases -> Rapid

9. CHANGING STATES

10. PHASE EQUILIBRIA & THE PHASE RULE

11.  A phase is defined as any homogeneous and physically
distinct part of a system which is separated from other parts
of the system by interfaces.
 A part of a system is homogeneous if it has identical physical
properties and chemical
composition throughout the part.
A phase may be gas, liquid or solid.
A gas or a gaseous mixture is a single phase.
Totally miscible liquids constitute a single phase.
In an immiscible liquid system, each layer is counted as a
separate phase.
Every solid constitutes a single phase except when a solid
solution is formed.
A solid solution is considered as a single phase.
Each polymorphic form constitutes a separate phase.
PHASE DEFINITION

12. Examples
The number of phases in a system is denoted P
1.Liquid water,pieces of ice and water vapour are present together.
The number of phases is 3 as each form is a separate phase. Ice in the
system is a single phase even if it is present as a number of pieces.
2.Calcium carbonate undergoes
thermal decomposition The chemical
reaction is:
CaCO3(s)  CaO(s) + CO2 (g)
Number of phases = 3 : This system consists of 2 solid phases, CaCO3 and
CaO and one gaseous phase, that of CO2.
3.Ammonium chloride undergoes thermal decomposition.
The chemical reaction is: NH4Cl(s) NH3 (g) + HCl (g) Number of phases =2
This system has 2 phases, one solid, NH4Cl and one gaseous, a mixture of NH3
and HCl
4.For a liquid system, according to the solubility to decide whether a system
consists of one phase or of two.
a solution of sodium chloride in water is a singlephase.
A pair of liquids that are partially miscible or immiscible is a two-phase
system(P=2)
5.A gas, or a gaseous mixture is a single phase. P=1

13. COMPONENTS
The number of components of a system at equilibrium
is the smallest number of independently varying
chemical constituents using which the composition of
each and every phase in the system can be expressed
Examples
Counting the number of components
1. The sulphur system is a one component system. All
the phases, monoclinic, rhombic, liquid and vapour
– can be expressed in terms of the single
constituent – sulphur.
2. A mixture of ethanol and water is an example of a
two component system. We need both ethanol and
water to express its composition.

14.  Phase Equilibrium:
A stable phase structure with lowest free-energy (internal
energy) of a system, and also randomness or disorder of the
atoms or molecules (entropy)
 Any change in Temperature, Composition and Pressure causes
an increase in free energy and away from Equilibrium thus
forcing a move to another‘state’
 Phase diagram or (Equilibrium Phase Diagram)
It summarizes the conditions at which a substance exists as a
solid, liquid, or gas. Or
It is a “map”of the information about the control of
phase structure of a particular material system.
 The relationships between temperature and the compositions
and the quantities of phases present at equilibrium are
represented.
PHASE EQUILIBRIUM & PHASE RULE

15. General Phase diagram

16. Key features
• The major features of a phase diagram are phase boundaries and the triple
point.
• Phase diagrams demonstrate the effects of changes in pressure and
temperature on the state of matter
• At phase boundaries, two phases of matter coexist (which two depends on
the phase transition taking place).
Phase boundary: The line in a phase diagram that indicates the conditions
under which two (transitioning) states of matter exist at equilibrium.
• Triple point is the point on the phase diagram at which three distinct
phases of matter coexist in equilibrium.
Key terms
• Critical point – the point in temperature and pressure on a phase diagram
where the liquid and gaseous phases of a substance merge together into a
single phase. Beyond the temperature of the critical point, the merged
single phase is known as a supercritical fluid
• Fusion(melting) (or freezing) curve – the curve on a phase diagram which
represents the transition between liquid and solid states
• Vaporization (or condensation) curve – the curve on a phase diagram
which represents the transition between gaseous and liquid states
• Sublimation (or deposition) curve – the curve on a phase diagram which
represents the transition between gaseous and solid states
• Critical temperature- it is the temperature above which the liquid state of a
substance no longer exist regardless of pressure
• Critical pressure- it is the pressure required to bring about liquefaction at
the critical temperature

17. PHASERULE
 J W Gibbs formulated the phase rule
It is a general relation between the variance F, the number of components
C, and the number of phases P, at equilibrium for a system of any
composition
 It’s a relation to determine the least number of intensive variable
(independent variable that do not depend on the volume or the size) that
can be changed without changing the equilibrium state of the system or
The least number required to define the state of the system,
which is called DEGREE OF FREEDOM F
F = C  P +
2
 For a system in equilibrium Phase rule
 The degrees of freedom or variance F of a system is defined as the
minimum number of variables such as
 temperature
 pressure
 concentration

18. EXAMPLES
1. A gaseous mixture of CO2 and N2.
Three variables: pressure, temperature and composition are
required to define this system. This is, hence, a trivariant
system.
2. A system having only liquid water has two degrees of
freedom or is bivariant. Both temperature and pressure
need to be mentioned in order to define the system.
3. If to the system containing liquid water, pieces of ice are
added and this system with 2 phases is allowed to come
to equilibrium, then it is an univariant system.
Only one variable, either temperature or
pressure need to be specified in order to define the system. If
the pressure on the system is maintained at 1 atm, then the
temperature of the system gets automatically fixed at 0oC, the
normal melting point of ice.

19. Degrees of
Freedom
= What you
can
control
What the
system
controls

F = C + 2 P
Can control
the no. of
components
added and P &
T
System decided
how many
phases to
produce given
the conditions
A WAY OF UNDERSTANDING THE GIBBS PHASE RULE:
THE DEGREES OF FREEDOM CAN BE THOUGHT OF AS THE DIFFERENCE
BETWEEN WHAT YOU (CAN) CONTROL AND WHAT THE SYSTEM
CONTROLS

20. ONE-COMPONENT
SYSTEMS
Phase diagram ofwater
Curve O -C
Sublimatio
nDeposition
Curve O-A
Vaporization
Condensation
Curve O -B
Melting
Freezin
g

21. A) AT THE TRIPLE POINT:
P = 3 (SOLID, LIQUID, AND GAS) C= 1
(WATER)
P + F = C + 2
F = 0 (NO DEGREE OF FREEDOM)
B) LIQUID-SOLID CURVE
P = 2
2+F = 1 + 2
F= 1
One variable (T or P) can be changed
For One component system (C=1) : 3 conditions may arise when P=1, P=2 and P=3
C) LIQUID
P =1
So F =2
Two variables (T and P) can be varied independently and the system will
remains a single phase

24. TWO COMPONENT EUTECTIC SYSTEMS
Figure shows the simplest of two component phase diagrams. The components are A
and B, and the possible phases are pure crystals of A, pure crystals of B, and liquid with
compositions ranging between pure A and pure B. Compositions are plotted across the
bottom of the diagram. Note that composition can be expressed as either a
percentage of A or a percentage of B, since the total percentage must add up to 100.
(Compositions might also be expressed as mole fraction of A or B, in which case the
total must add up to 1). Temperature or pressure is plotted on the vertical axis. For the
case shown, we consider pressure to be constant, and therefore have plotted
temperature on the vertical axis.
The curves separating the fields of A + Liquid from Liquid and B + Liquid from Liquid are
termed liquidus curves. The horizontal line separating the fields of A + Liquid and B +
Liquid from A + B all solid, is termed the solidus. The point, E, where the liquidus curves
and solidus intersect, is termed the eutectic point. At the eutectic point in this two
component system, all three phases, that is Liquid, crystals of A and crystals of B, all exist
in equilibrium. Note that the eutectic is the only point on the diagram where this is true.

25. Two component system
C=2 When P=1
Degree of freedom F = C – P + 2
F = 2 – 1 + 2 = 3
Where two component systems (solid and liquid phases only) are present
then, the effect of pressure can be neglected
Condensed system: solid / liquid System with the absent vapor phase (gas
phase) is called condensed system
Remaining variables temperature and concentration
Hence, experimental measurements of temperature an concentration in
condensed systems are carried out under atmospheric pressure, given that
the F is reduced by one
Then rule of the reduced phase
as F’ = C – P + 1
F = 2 – 1 + 1 = 2 bivariant
So reduced phase rule is more convenient to apply to the two component
solid/ liquid condensed system
When P=2
Then F’ = C – P + 1
F’ = 2 – 2 + 1 = 1 monovariant
When P=3 (at eutectic
point) F’ = C – P + 1
F’ = 2– 3 + 1 = 0 non
variant

26. A eutectic mixture is defined as a mixture of two or more components
which usually do not interact to form a new chemical compound but,
which at certain ratios, inhibit the crystallization process of one another
resulting in a system having a lower melting point than either of the
components.
Eutectic mixture formation is usually, governed by following factors:
(a)components must be miscible in liquid state and mostly immiscible in
solid state
(b)Intimate contact between eutectic forming materials is necessary for
contact induced melting point depression
(c)the components should have chemical groups that can interact to form
physical bonds such has intermolecular hydrogen bonding etc.
(d)the molecules which are in accordance to modified VantHoff’s equation
can form eutectic mixtures.
Eutectic mixtures, can be formed between Active Pharmaceutical
Ingredients (APIs), between APIs and excipients or between excipients;
thereby providing a vast scope for its applications in pharmaceutical
industry
• During pre-formulation studies- Testing foreutectic mixture formation
can help in anticipation of probable physical incompatibility between
drug and excipient molecules
• Eutectic mixtures are commonly used in drug designing and delivery
processes for various routes of administration
Eutectic mixture Pharmaceutical Application

27. During tablet compaction the heat produced in the punch and die
cavities may lead to fusion or melting of tablet powder compacts
leading to manufacturing defects. Thus knowledge of eutectic points
of powder components may help avoid these problems
During pharmaceutical analysis, understanding of eutectic mixtures
can help in the identification of compounds having similar melting
points. Compounds having similar melting points, as a rule will have
different eutectic point with a common other component. This
knowledge could be used to identify compounds like Ergotamine,
Allobarbital etc.
 EMLA® (lidocaine 2.5% and prilocaine 2.5%) Cream
 EMLA Cream (lidocaine 2.5% and prilocaine 2.5%) is an emulsion in
which the oil phase is a eutectic mixture of lidocaine and prilocaine in
a ratio of 1:1 by weight. This eutectic mixture has a melting point
below room temperature and therefore both local anaesthetics exist
as a liquid oil rather than as crystals

28. Tm: Melting point temperature e: Eutectic Mixture; IDR: Intrinsic Dissolution
Rate; AUC: Area Under Curve; PEG: Poly Ethylene Glycol
Sl
No
Eutectic
Component
s (a,b)
Ratio Tm °C
(a)
Tm °C
(b)
Tm °C
(e)
Challenges Findings
1 Curcumin,
Nicotinamid
e
1:2 181.4 128.3 110.5 Oral route-Low
solubility, poor
oral
bioavailbility
10-fold faster IDR and 6-
times higher AUC compared
to crystalline curcumin
2 Ibuprofen
, Thymol
2:3 76.0 52.0 32.0 Transdermal
Route
Limited
ability to
penetrate
the skin
A flux of 150 mg/ cm / h, 5.9
times the flux from a
saturated aqueous solution
with thymol pretreated skin
and 12.7 times the flux from
a saturated aqueous solution
across non-pretreated skin
3 Genistein,
PEG 460
1:24 305.0 2.0 0.2 Parentral
Route Low
aqueous
solubility, thus
formulation
difficulty.
Could help in solubilization
of genistein crystals for
injection development
APPLICATIONS OF EUTECTIC MIXTURES IN FORMULATION DEVELOPMENT

29. LATENT HEAT
HEAT REQUIRED TO PRODUCE A PHASE CHANGE IS
CALLED LATENT HEAT (LH )
Examples:
heat of freezing- amount of thermal energy released when liquid
freezes Heat of vapourization- amount of thermal energy added to
transform liquid into gas
Lh of vapourization of water : 540 cal /g
Lh of freezing of water : 80 cal /g
Heat added or subtracted for a phase change = latent heat x mass
Q = Lh x M
Where, Q =
heat
Lh = latent heat M =
mass

30. VAPOUR PRESSURE
What is Vapour pressure ?
• Vapour pressure is defined as the pressure exerted by a vapour in
thermodynamic equilibrium with its condensed phases (solid or
liquid) at a given temperature in a closed system
• Vapour pressure is nothing but the tendency of particles to escape
from the liquid (or a solid)
• At normal temperatures, substance with a high vapour pressure is
often referred to as volatile.
Characteristics of Vapour Pressure
 A pure liquid experiences a greater amount of vapour pressure as
against a liquid’s solution
 It is inversely proportional to the forces of attraction existing between
the molecules of a liquid
 It increases with a rise in the temperature. This is because the
molecules gain kinetic energy and thus, vapourise briskly
 The vapour pressure of a liquid depends on the nature of the liquid.
The low- boiling liquid exerts more vapour pressure at a given
temperature
 Liquids in which the intermolecular forces are weak shows high
vapour pressure

31. Thank You