Solution Properties
Colligative Properties
Ahmed Talaat Nouh
Colligative Properties of Solution
What are Colligative properties of a solution...'why colligative'?
 Colligative properties are those properties that depend on the
concentration of the solute particles rather than the kind of solute
particles.
 All colligative properties are related to each other by virtue of their
common dependency on the concentration of the solute molecules.
 From Greek word “collected together”
They include:
 osmotic pressure
 vapor pressure lowering
 boiling point elevation
 freezing point depression
1. Osmotic pressure
 The most important colligative property from a pharmaceutical
point of view is referred to as osmotic pressure.
 The osmotic pressure of a solution is the external pressure that
must be applied to the solution in order to prevent it being diluted
by the entry of solvent via a process known as osmosis.
 Osmosis:
• If two solutions of different concentrations are separated by a
semi-permeable membrane (only permeable to the solvent) the
solvent will move from the solution of lower solute concentration
to that of higher solute concentration.
Osmotic pressure of nonelectrolytes:
 The relationship between osmotic pressure and the concentration of
a non-electrolyte is given for dilute solutions, which may be
assumed to exhibit ideal behaviour, by the van't Hoff equation:
PV=nRT
 where V is the volume of solution, n, is the number of moles of
solute, T is the absolute temperature and R is the gas constant.
 The osmotic pressure of solutions of different nonelectrolytes is
proportional to the number of molecules in each solution.
 The osmotic pressures of two nonelectrolyte solutions of same
molal concentration are identical.
 For example, a solution containing 34.2g of sucrose (mol wt. 342)
in 1000 g of water has the same osmotic pressure as a solution
containing 18.0 g of anhydrous dextrose (mol wt. 180) in 1000 g
water. These solutions are said to be iso-osmotic with each other
because they have identical osmotic pressures.
 But this is not true for electrolytes as they ionize therefore the
number of particles increase  o.p. ↑
Osmotic pressure of electrolytes:
 If the solute is an electrolyte, the previous equation must be modified
to allow for the effect of ionic dissociation, because this will increase
the number of particles in the solution. This modification is achieved
by insertion of the van't Hoff correction factor (i) to give:
PV=inRT
 Although the above equation may be simpler to remember, the
following form of the equation is more useful. This form of the
equation has been derived by realizing that n/V gives the
concentration of the solute in units of molarity, M.
P=iMRT
 Where the value of i approaches the number of ions produced by the
ionization of the strong electrolytes. For weak electrolytes i represents
the total number of particles, ions and molecules together, in the
solution.
2. Vapor pressure lowering:
 To understand why that might occur, let's analyze the vaporization
process of the pure solvent then do the same for a solution.
 Liquid molecules at the surface of a liquid can escape to the gas
phase when they have a sufficient amount of energy to break free of
the liquid's intermolecular forces. That vaporization process is
reversible. Gaseous molecules coming into contact with the surface
of a liquid can be trapped by intermolecular forces in the liquid.
 Eventually the rate of escape (vaporization) will equal the rate of
capture (condensation) to establish a constant, equilibrium vapor
pressure above the pure liquid.
Vapor pressure is the pressure exerted by the vapor at
equilibrium.
 When a non volatile solute is dissolved in a liquid solvent the vapor
pressure of the solvent is lowered. WHY?
 Some of the surface molecules of the solvent are replaced by solute
molecules which do not contribute in the vapor pressure and the
surface area available for the escaping solvent molecules is reduced
because some of that area is occupied by solute particles.
B = Pure solvent
A= Solution
• Lowering of Vapour Pressure:
• Vapour pressure P1 of solvent over a dilute solution equal to
vapour pressure of pure solvent times the mol fraction of
solvent X1
• Because the solute non volatile, so the total pressure = the
pressure of the solvent
• ΔP = Po - P = Vaspour pressure lowering
3. Boiling point elevation
 Boiling point elevation is a colligative property related to vapor
pressure lowering.
 The boiling point is defined as the temperature at which the
vapor pressure of a liquid equals the atmospheric pressure (760
mmHg).
 Due to vapor pressure lowering, a solution will require a higher
temperature to reach its boiling point than the pure solvent.
 The boiling point of pure water is 100°C, but that boiling point can
be elevated by the adding of a solute such as a salt. A solution
typically has a measurably higher boiling point than the pure
solvent.
Phase Diagram for a Solution and the Pure Solvent Indicating the Boiling
Point Elevation
Determination of The boiling point elevation:
 The boiling point elevation DTb is a colligative property of the
solution, and for dilute solutions is found to be proportional to the
molal concentration cm of the solution:
DTb = Kb cm
Where Kb is called the boiling-point-elevation
constant(Ebullioscopic constant).
 Solutions may be produced for the purpose of raising the boiling
point and lowering the freezing point, as in the use of ethylene
glycol in automobile cooling systems. The ethylene glycol
(antifreeze) protects against freezing by lowering the freezing point
and permits a higher operating temperature by raising the boiling
point.
4. Freezing point depression
 The freezing point is depressed due to the vapor pressure lowering
phenomenon.
 Freezing point (or melting point): is defined as the
temperature at which the solid and the liquid phases are in
equilibrium under a pressure of 1 atm.
 The freezing point of pure water is 0°C, but that melting point can
be depressed by the adding of a solute such as a salt.
 The use of ordinary salt (sodium chloride, NaCl) on icy roads in the
winter helps to melt the ice from the roads by lowering the melting
point of the ice.
 A solution typically has a measurably lower melting point than the
pure solvent. A 10% salt solution was said to lower the melting point
to -6°C and a 20% salt solution was said to lower it to -16°C.
Phase Diagram for a Solution and the Pure Solvent Indicating the Freezing
Point Depression
Determination of The freezing point depression:
 The freezing point depression ΔTf is a colligative property of the
solution and for dilute solutions is found to be proportional to the
molal concentration cm of the solution:
ΔTf = Kf cm
Where Kf is called the freezing-point-depression constant (cryoscopic
constant).
 A pleasant application of the freezing point depression is in the making of
homemade ice cream.
 The ice cream mix is put into a metal container which is surrounded by
crushed ice. Then salt is put on the ice to lower its melting point. The
melting of the solution tends to lower the equilibrium temperature of the
ice/water solution to the melting point of the solution. This gives a
temperature gradient across the metal container into the saltwater-ice
solution which is lower than 0°C. The heat transfer out of the ice cream mix
allows it to freeze.
Practical Applications of Colligative Properties:
1.
Preparation of isotonic intravenous and isotonic lachrymal
solutions.
2.
Determination of the molecular weight of solutes or in the
case of electrolytes, the extent of ionization.
3.
They also may be used in experimental physiology as in
immersion of tissues in salt solutions which are isotonic with
the tissue fluids to prevent changes or injuries of the tissues.
Molarity :
M,c Moles (gram molecular weight) of solutes in 1liter of solution
Moles = gm/M.W
Molality :
Moles of solutes in 1000 gm of solvents
Mole Fraction: Ratio of the moles of one constituent of a solution to
the total moles of all constituents
Mole fraction = n1 / n1 + n2
= n2 / n 1 + n 2
n1 = W1 / M1
n2 = W2 / M2
Problems
• Calculate the relative vapour pressure lowering at 20c for
a solution containing 171.2 grams of sucrose (w2) in 1000
grams (w1) of water. The molecular weight of sucrose (
M2 ) in 342.3 and the molecular weight of water ( M1) is
18.02
• Answer:
Moles of Sucrose = n2 = W2/M2 = 171.2 / 342.3 = 0.500
Moles of Water = W1/M1 = 1000/18.02 = 55.5
∆p/p1 = X2 = n2 / n1 + n2
∆p/p1 = 0.05 / 55.5 + 0.50 = 0.0089