15
Foreword
Principles of Freezing-Point Osmometry
When a solute is dissolved in a pure solvent, the following changes in
the solution’s properties occur:
• the freezing point is depressed,
• boiling point is raised,
• osmotic pressure is increased, and
• vapor pressure is lowered.
These are the so-called colligative or concentrative properties of the
solution which, within reasonable limits, change in direct proportion
to the solute concentration; in other words, the number of particles in
solution.
Of the colligative properties, measurement of the freezing point allows
the concentration of an aqueous solution to be easily determined with
great precision.
The freezing point of pure H
O is precisely +0.010°C. One mole of a
non-dissociating solute such as glucose (where the solute does not
dissociate into ionic species, but remains intact), when dissolved in
1 kilogram (kg) of water will depress the freezing point by 1.858°C.
This change is known as the freezing point depression constant for
water. The freezing point depression also depends upon the degree
of dissociation of the solute. If the solute is ionic, the freezing point is
depressed by 1.858°C for each ionic species. For example, if one mole
of sodium chloride were to completely dissociate into two ionic species
(Na+ and Cl-) in 1 kg of water, the freezing point would be depressed
by 3.716°C. However, dissociation is never complete. Interference
between solute molecules reduces dissociation by a factor called the
osmotic coefficient.
In a simple solution such as glucose or sodium chloride in water, the
freezing point can be measured and the unit concentration easily
determined from an equation or a reference table. However, the
equation is unique for each solute. In a more complex solution, all
ionized and non-dissociated species contribute to the freezing-point
depression and the concentration of each solute cannot be easily
determined.
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