This document is a chemistry investigatory project on colligative properties, completed by Chinmaya Jagadev Pattanayak under the guidance of Mr. S.K. Bagh. It details the definition of colligative properties, their types such as lowering vapor pressure, boiling point elevation, freezing point depression, and osmotic pressure, along with relevant equations and concepts. The project is certified as part of the CBSE curriculum for the session 2015-2016.

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CHEMISTRY
INVESTIGATORY
PROJECT
SESSION-2015-16
TOPIC: - COLLIGATIVE PROPERTY
MADE BY: - CHINMAYA JAGADEV
PATTANAYAK

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GUIDED BY: -MR. S.K.BAGH (PGT Chemistry)

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Certificate
This is to certify that Master Chinmaya
Jagadev Pattanayak bearing Roll No. __________ of
class-XII(science.) JAWAHAR NAVODAYA VIDYALAYA,
Sarang, Dhenkanal (ODISHA) has successfully completed
the project on the topic “COLLIGATIVE PROPERTY” under
the guidance of Mr.S.K. Bagh (P.G.T.Chemistry). This
project can be approved as a part of C.B.S.E in the
session 2015-2016.

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Signatureofstudent:
SignatureofGuide:- SignatureofPrincipal:
MR. S.K. BAGH MR. N.C. KAR

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ACKNOWLEDGEMENT
I wish to express my deep sense of gratitude to them
whose encouragement and cooperation has been a
source of inspiration. It’s an honor to thank our
principal Mr.N.C.Kar, Jawahar Navodaya Vidyalaya,
Sarang,Dhenkanal for providing me the opportunity
to makethis project andlearn through it. Iwould like
to thank Mr.S.K. Bagh (PGT Chemistry) for
providing us the right direction and for his
cooperationandhelpin successful completionof our
project.
I amindebtedtobothofthemfortheirhelp.
.

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CONTENTS:-
What is colligative property?
Types of colligative property
Lowering Vapour Pressure (∆P) of
solutions.
Boiling point elevation.
Freezing point depression
Osmotic pressure of the solution.

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WHAT IS COLLIGATIVE PROPERTY?
In chemistry, colligative properties are properties
of solutions that depend upon the ratio of the number of
solute particles to the number of solvent molecules in a
solution, and not on the type of chemical species
present. The number ratio can be related to the various units
for concentration of solutions. The independence of the
nature of solute particles is exact for ideal solutions, and
approximate for dilute real solutions.
Here we consider only those properties which result because
of the dissolution of nonvolatile solute in a volatile liquid
solvent. They are independent of the nature of the solute
because they are due essentially to the dilution of the solvent
by the solute. The word colligative is derived from the
Latin colligatus meaning bound together.
TYPES OF COLLIGATIVE PROPERTY:-
1. Lowering Vapour pressure of solution: Roult’s law.
2. Elevation in boiling point (ΔTB ).
3. Depression in freezing point (ΔTf ).

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4. Osmosis.

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1.Lowering Vapour Pressure (∆P) of solutions: Roult’s
law:-
When a non-volatile solute is added to a solvent, the
vapour pressure of the solution decreases.
According to Roult’s Law, the vapour pressure of a solvent
(P1) in a solution containing a non-volatile solute is given
by
According to Raoult's Law,
Vapour pressure of the pure solvent = P1°

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Vapour pressure of the solvent in solution = P1
P1 = x1P1°
ΔP1 = P1° - P1
= P1° - x1P1°
= P1° (1 - x1)
In a binary solution, 1 - x1 = x2
ΔP1 = P1° x2
ΔP1/P1° = (P1° - P1)/P1° = x2
The lowering of vapour pressure relative to the vapour
pressure of pure solvent is called relative lowering of
vapour pressure.
ΔP1/P1° → Relative lowering of Vapour pressure
Thus, the relative lowering in vapour pressure depends
only on the concentration of solute particles and is
independent of their identity.
If the solution contains more than one non-volatile solute,
then the relative lowering in vapour pressure of a solvent is

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equal to the sum of the mole fractions of all the
non-volatile solutes.
If n1 and n2 are respectively the number of moles of the
solvent and solute in a binary solution, then the relative
lowering in the vapour pressure of the solvent,
(P1° - P1)/P1° = x1 + x2 + x3 + ... + xn
if n1 and n2 are the number of moles of the solvent and
solute,
(P1° - P1)/P1° = n2/(n1+n2)
For dilute solutions n2 << n1
(P1° - P1)/P1° = n2/n1
n1 = W1/M1 , n2 = W2/M2
(P1° - P1)/P1° = (W2xM1)/(W1xM2)
W1 = Mass of solvent
W2 = Mass of solute
M1 = Molar mass of solvent
M2 = Molar mass of solute

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2. Boiling point elevation (ΔTB ):-
The exact relation between the boiling point of the solution
and the mole fraction of the solvent is rather complicated,
but for dilute solutions the elevation of the boiling point is
directly proportional to the molal concentration of the
solute:
OR ΔTb = Kb.Cm
Kb = Ebullioscopy constant, which is 0.512°C kg/mol
for the boiling point of water.
The vapour pressure of a liquid increases with an increase in
temperature. When vapour pressure of the liquid becomes
equal to the atmospheric pressure (or) external pressure,
then liquid starts boiling. The temperature at which the
vapour pressure of the liquid is equal to the external
pressure is known as its boiling point.

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At any temperature, the vapour pressure of a solution
containing a non-volatile solute is less than that of the pure
solvent.
The temperatures at which the vapour pressure of the
solvent and the solution become equal to the atmospheric
pressure are Tb0
and Tb.
Tb-Tb0
=∆Tb
Thus, it can be seen that the boiling point of a solution is
greater than the boiling point of the pure solvent.
The boiling point of a solvent changes as the concentration
of the solute in the solution changes, but it does not depend
on the identity of the solute particles.
The elevation of the boiling point depends upon the
concentration of the solute in the solution and is directly
proportional to molality (m) of the solute in the solution.
ΔTb = Tb - Tb°
Tb > Tb°
ΔTb ∝ Concentration of solute
ΔTb ∝ m (Molarity)
ΔTb = Kb m
Kb = Boiling point elevation constant or molal elevation
constant or ebullioscopic constant

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Molal elevation constant is defined as the elevation in the
boiling point when 1mole of a solute is dissolved in
1kilogram of a solvent.
If w2 grams of a solute with M2 molar mass is dissolved in
w1gram of a solvent, then molality (m) of the solution is,
m = (W2x1000)/(W1xM2)
ΔTb = Kb (W2x1000)/(W1xM2)
3. Freezing Point Depression (ΔTf ):-
The freezing point of a substance is defined as the
temperature at which its solid phase is in dynamic
equilibrium with its liquid phase. At the freezing point, the
vapour pressure of the substance in its liquid phase is the
same as the vapour pressure of the substance in its solid
phase.

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When a non-volatile solute is added to a solvent, the
freezing point of the solution gets lowered.
According to Roult’s law, the vapour pressure of a solution
containing a non-volatile solute is lower than that of the
pure solvent. Thus freezing point of a solvent decreases
when a non-volatile solute is added to it.
The depression in freezing point depends upon the
concentration of the solution. For dilute solutions,
depression in the freezing point is directly proportional to
molality (m).
Thus, ∆Tf =Kf m
Where
Kf =freezing point depression constant (or) molal
depression constant (or) cryoscopic constant.
Molal depression constant Kf can be defined as the
depression in freezing point when 1mole of solute dissolved
in 1kg of solvent. The unit for Kf is kelvin kilogram /mole.
As Kf depends upon the nature of the solvent, its value is
different for different solvents.
The values of Kf can be calculated from this expression
Kf = (R x M1 x Tf
2
)/(1000 x ΔfusH)
R = Gas constant

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M1= Molar mass of the solvent
Tf = Freezing point of the pure solvent
ΔfusH = Enthalpy for the fusion of the solvent
If w2 grams of a solute with molar mass M2 is dissolved in
w1 grams of a solvent, then molality m of the solution is
given by W2 multiplied by 1,000 divided by w1 multiplied
by M2
Substituting this value of molality in the freezing point
depression equation, we get depression in freezing point
Molarity , m = (W2 x 1000)/(W1xM2)
ΔTf = Kf m
ΔTf = (Kf x W2 x 1000)/(W1xM2)
M2 = = (Kf x W2 x 1000)/(W1xΔTf)
Thus, the molar mass of a non-ionic solute can be calculated
by using the depression in freezing point.
4. Osmosis:-
Membranes which allows only solvent particles but not
solute particles of solution is called semi-permeable
membranes (or) SPM. These membranes can be of natural
origin (or) synthetic origin.

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Vegetable membranes, membranes found under the shell
of an egg are examples of natural membranes and
cellophane is an example of synthetic membrane.
Thus 'Osmosis' can be defined as the spontaneous flow of
solvent through a semi-permeable membrane from a pure
solvent to a solution or from a dilute solution to a
concentrated solution.
It is important to note that osmosis drives solvent
molecules through a semi-permeable membrane from low
solute concentrations to high solute concentrations.
Osmosis ends when the solute concentration becomes
equal on either side of the membrane and equilibrium is
attained.
The flow of solvent molecules from low concentration to
high concentration can be stopped by applying some extra
pressure on the high concentration side. The minimum
pressure required to do so is known as the osmotic pressure
of the solution.

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Thus, osmotic pressure π of a solution is defined as the
excess pressure that must be applied to a solution to
prevent osmosis from taking place.
Osmotic pressure does not depend on the identity of the
solute, but on its concentration.
Osmotic pressure for dilute solutions is proportional to
molarity of the solution at a given temperature(T).
π ∝ C (at given T)
π = C R T
R = Gas constant
C = n2/V
π = n2RT / V
If W2 grams of solute of molar mass M2 is present in the
solution,
n2 = W2/M2

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π = W2RT / M2V
M2 = W2RT / πV
This is widely used to determine the molar masses of
polymers and macromolecules, especially biomolecules, as
they are generally unstable at higher temperatures and
decompose before their boiling point is reached.
If the solutions have the same concentrations (C1 = C2),
then π1= π2.

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BIBLIOGRAPHY:-
1. www.google.com
2. www.slideshare.com
3. www.icbse.com
4. www.wikipedia.com
5. www.yahoo.com