1. States of Matter
Dr. Prashant L. Pingale
Associate Professor-Pharmaceutics
GES’s Sir Dr. M. S. Gosavi College of Pharm. Edu. & Research, Nashik

2. Content…
 State of matter,
 Changes in the state of matter,
 Latent heats,
 Vapour pressure,
 Sublimation critical point,
 Eutectic mixtures,
 Gases: Liquefaction of gases, aerosols– inhalers, relative humidity, liquid complexes, liquid crystals, glassy
states,
 Solid: Crystalline, amorphous
 Methods of crystal analysis: X-Ray Diffraction, Bragg’s equation
 Polymorphism (Definition, Different shapes of polymorphs, Example and its Pharmaceutical applications, Brief
introduction of Detection techniques).
 Physicochemical properties of drug molecules: Refractive index, optical rotation, dielectric constant,
dipole moment, dissociation constant, déterminations and applications.
2

3. States of Matter
 States of matter are the physical forms a substance can take.
 There are three common states of matter: solid, liquid, and gas.
 Each of these states is also called a phase
3

4. The three common states of matter
Most substances, like water, can exist in all three states
An iceberg is made of
water in solid form.
This glass contains
liquid water
A cloud is made of water
vapor, a type of gas.
4

5. States of matter: Solids
 Particles of solids are tightly packed, vibrating
about a fixed position.
 Solids have a definite shape and a definite volume.
 Solid particles vibrate in place but cannot move
from their position, which is why solids maintain
their rigid shape.
Heat
5

6. States of matter: Liquid
 Particles of liquids are tightly packed, but are
far enough apart to slide over one another.
 Liquids have an indefinite shape and a definite
volume.
 Liquid particles move slightly, which allows
liquids to flow and take the shape of the
container they are in.Heat
6

7. States of matter: Gas
 Particles of gases are very far apart and
move freely.
 Gases have an indefinite shape and an
indefinite volume.
 Gas particles move freely and will expand
to fill a container of any size or shape.
Heat
7

8. States of Matter
SOLID LIQUID GAS
Tightly packed, in a
regular pattern
Vibrate, but do not
move from place to
place
Close together
with no regular
arrangement.
Vibrate, move
about, and slide
past each other
Well separated
with no regular
arrangement.
Vibrate and move
freely at high
speeds
8

9. Three states of matter
Solid Liquid Gas
 At room temperature most substances exist in one of three physical states.
9

10. Changes of state
 Matter can change from one state to another.
 Even though the physical form of the matter changes, it remains the same
substance.
 Changes of state occur when thermal energy (heat energy) is absorbed or
released by a substance.
10

11. How does matter change state?
 As heat increases, a substance changes from a solid to a liquid, and finally to a gas.
 As heat decreases, a substance changes from a gas to a liquid, and finally to a solid.
11

12. Latent Heat
 The latent heat associated with melting a solid or freezing a liquid is called the heat of fusion.
 If it is associated with vaporizing a liquid or a solid or condensing a vapour is called the heat of
vaporization.
 The latent heat is normally expressed as the amount of heat (in units of joules or calories) per mole or
unit mass of the substance undergoing a change of state.
12
 The term was introduced around 1762 by British chemist Joseph
Black.
 Latent heat, energy absorbed or released by a substance during a
change in its physical state (phase) that occurs without changing
its temperature.
 Latent heat is defined as the heat or energy that is absorbed or
released during a phase change of a substance. It could either be
from a gas to a liquid or liquid to solid and vice versa. Latent
heat is related to a heat property called enthalpy.

13. Explanation of Latent Heat
 When the change of state is studied carefully, the temperature of a substance remains
constant during a change in the state.
 This is very extraordinary. As if the change in state opens up new portals or spaces where
the supplied energy hides, therefore it may be referred as hidden energy, the latent or the
hidden heat.
 Example:
13

14. Explanation of Latent Heat
 Suppose we have a block of ice we want to convert to water.
 We all know that ice turns to water and vice versa at 0°C.
 Now assume, we start heating ice at 0°C.
 You will observe that when we do so, the temperature of ice does not change.
 It starts converting to water but the temperature does not rise until the entire ice block has
been converted to water.
 But we are heating the ice block right? So, what happened.
 If a mass ‘m’ of any substance undergoes a change in state by absorbing an amount of heat, Q at
a constant Temperature T, then we have:
L = Q/m or Q = mL
 All the heat supplied to the ice at 0 0C is used by the ice to change its phase from solid to liquid.
Thus the heat supplied is not used up to raise the temperature of the substance. There are 2
kinds of Latent heat: Latent Heat of Fusion and Latent Heat of Vaporization
14

15. Latent Heat: Types
 Two common forms of latent heat are latent heat of fusion (melting) and latent heat of
vaporization (boiling).
 Latent Heat of Fusion: The heat energy supplied per unit mass of a substance at its
melting point to convert the state of the substance from solid to liquid is known as
Latent heat of Fusion. Latent heat of Fusion of water is 334 Joules/gram of water.
 Latent Heat of Vaporization: The heat that a substance absorbs per unit mass at its
boiling point to convert the phase of the substance from liquid to gas is the Latent heat of
Vaporization. Latent heat of Vaporization of water is 2230 Joules/gram of water. (Click
for video: https://youtu.be/hiGzXJ3emcY)
15

16. Vapour Pressure
 When the liquid is heated, the energy of the molecules rises; it becomes lighter and
occupies the surface of the liquid. This process is known as ‘evaporation’. The molecules
which can be seen on the liquid surface are called ‘vapor’.
 Vapour pressure is a measure of the tendency of a material to change into the gaseous
or vapour state, and it increases with temperature. The temperature at which the
vapour pressure at the surface of a liquid becomes equal to the pressure exerted by the
surroundings is called the boiling point of the liquid.
 When some molecules of the liquid in the vapor phase, strikes the walls of the
containers or the surface of the liquid, it may get converted back to the liquid phase.
This process is called condensation.
16

17. Vapour Pressure
 The evaporation continues at a constant rate the temperature of the
liquid is kept constant.
 As time passes, the number of molecules in the vapor phase
increases while the rate of condensation also increases. It reaches a
stage where the rate of evaporation is equal to the rate of
condensation. This phase is called the stage of equilibrium.
 As represented by the manometer, at this point the pressure exerted
by the molecules is called the vapor pressure of the liquid. Vapour
pressure is defined as the pressure exerted by the vapor present
above the liquid.
17

18. Factors affecting
Vapour Pressure
 A liquid’s vapor pressure is a vapor’s equilibrium pressure above its liquid (or solid); that
is, the vapor pressure resulting from a liquid (or solid) evaporation above a liquid (or
solid) sample in a closed container.
 Nature of the liquid: Liquids have weak intermolecular forces. Heating the molecules of
the liquid can help change them to the vapor phase and thus increase the vapor pressure
of the liquid. For example, Acetone and benzene have higher vapor pressure than water at
a particular temperature.
 Effect of temperature: The vapor pressure of the liquid increases with an increase in its
temperature. The molecules of the liquid have higher energy at higher temperatures.
18

19. Sublimation
 Sublimation is the transition from the solid phase to the gas phase without passing through an
intermediate liquid phase.
 This endothermic phase transition occurs at temperatures and pressures below the triple
point.
 The term "sublimation" only applies to physical changes of state and not to the transformation
of a solid into a gas during a chemical reaction.
 The opposite process of sublimation- where a gas undergoes a phase change into solid form-is
called deposition or desublimation.
19

20. Example of sublimation
 Dry ice is solid carbon dioxide. At room temperature and pressure, it sublimates into carbon
dioxide vapor. When dry ice gets exposed to air, dry ice directly changes its phase from solid-
state to gaseous state which is visible as fog. Frozen carbon dioxide in its gaseous state is
more stable than in its solid-state. (https://youtu.be/39Cb-qG8Ozk)
 Naphthalene is usually found in pesticides such as mothball. This organic compound sublimes
due to the presence of non-polar molecules that are held by Van Der Waals intermolecular
forces. At a temperature of 176F naphthalene sublimes to form vapors. It desublimates at cool
surfaces to form needle-like crystals (https://youtu.be/29x7vArV2NI)
20

21. Critical Point
 A critical point (or critical state) is the end point of a phase
equilibrium curve.
 The most prominent example is the liquid-vapor critical
point, the end point of the pressure-temperature curve that
designates conditions under which a liquid and its vapor
can coexist.
 A point at which, two phases of a substance initially become
indistinguishable from one another. The critical point is the
end point of a phase equilibrium curve, defined by a critical
pressure Tp and critical temperature Pc. At this point, there
is no phase boundary.
21

22. Critical Point vs Triple Point22

23. Critical Point vs Triple Point23

24. Critical Point: Example
 The liquid-vapor critical point is the most common example,
which is at the end point of the pressure-vapor temperature
curve distinguishing a substance's liquid and vapor.
 The meniscus between steam and water vanishes at
temperatures above 374°C and pressures above 217.6 atm,
forming what is known as a supercritical fluid.
 There is also a liquid-liquid critical point in mixtures, which
occurs at the critical solution temperature.
24
Tubes containing water at several
temperatures. Note that at or above 374oC
(the critical temperature for water), only
water vapor exists in the tube.

25. Eutectic mixtures
 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 mixtures, can be formed between Active Pharmaceutical Ingredients (APIs), between
APIs and excipient or between excipient; thereby providing a vast scope for its applications in
pharmaceutical industry.
25

26. Factors influencing eutectic mixture formation
 Eutectic mixture formation is usually, governed by following factors:
 The components must be miscible in liquid state and mostly immiscible in solid state,
 Intimate contact between eutectic forming materials is necessary for contact induced
melting point depression,
 The components should have chemical groups that can interact to form physical bonds such
has intermolecular hydrogen bonding etc.,
 The molecules which are in accordance to modified VantHoff’s equation can form eutectic
mixtures.
26

27. Applications of Eutectic Mixtures
 During pre formulation stage, compatibility studies between APIs and excipient play a crucial
role in excipient selection.
 Testing for eutectic 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.
 During manufacturing of pharmaceutical dosage form, it is extremely necessary to anticipate the
formation of eutectics and avoid manufacturing problems if any.
 During pharmaceutical analysis, understanding of eutectic mixtures can help in the
identification of compounds having similar melting points
27