What Are Colligative Properties? Definition and Examples   Recently updated !

Colligative Properties
Colligative properties depend on the number of solute particles, not their identity.

In chemistry, colligative properties are characteristics of chemical solutions that depend on the number of solute particles compared to solvent particles, not on the chemical identity of the solute particles. However, colligative properties do depend on the nature of the solvent. The four colligative properties are freezing point depression, boiling point elevation, vapor pressure lowering, and osmotic pressure.

Colligative properties apply to all solutions, but the equations used to calculate them apply only to ideal solutions or weak solutions of a nonvolatile solute dissolved in a volatile solvent. It takes more complicated formulas to calculate colligative properties for volatile solutes. The magnitude of a colligative property is inversely proportional to molar mass of the solute.

How Colligative Properties Work

Dissolving a solute in a solvent introduces extra particles between solvent molecules. This reduces the concentration of the solvent per unit of volume, essentially diluting the solvent. The effect depends on how many extra particles there are, not their chemical identity. For example, dissolving sodium chloride (NaCl) yields two particles (one sodium ion and one chloride ion), while dissolving calcium chloride (CaCl2) yields three particles (one calcium ion and two chloride ions). Assuming both salts are fully soluble in a solvent, calcium chloride has a greater effect on the colligative properties of a solution than table salt. So, adding a pinch of calcium chloride to water lowers freezing point, increases boiling point, lowers vapor pressure, and changes osmotic pressure more than adding a pinch of sodium chloride to water. This is why calcium chloride acts as a de-icing agent at lower temperatures than table salt.

The 4 Colligative Properties

Freezing Point Depression

Freezing points of solutions are lower than freezing points of pure solvents. The depression of the freezing point is directly proportional to solute molality.

Dissolving sugar, salt, alcohol, or any chemical in water lowers the freezing point of the water. Examples of freezing point depression include sprinkling salt on ice to melt it and chilling vodka in a freezer without freezing it. The effect works in other solvents besides water, but the amount of the temperature change varies by solvent.

The formula for freezing point is:

ΔT = iKfm


ΔT = Change in temperature in °C
i = van ‘t Hoff factor
Kf = molal freezing point depression constant or cryoscopic constant in °C kg/mol
m = molality of the solute in mol solute/kg solvent

There are tables of molal freezing point depression constants (Kf) for common solvents.

SolventNormal Freezing Point (oC)Kf (oC/m)
acetic acid16.663.90
carbon tetrachloride-22.9529.8
Freezing point depression constants

Boiling Point Elevation

The boiling point of a solution is higher than the boiling point of the pure solvent. As with freezing point depression, the effect is directly proportional to solute molality. For example, adding salt to water increases the temperature at which it boils (although not by a lot).

Boiling point elevation may be calculated from the equation:

ΔT = Kbm


Kb = ebullioscopic constant (0.52°C kg/mol for water)
m = molality of the solute in mol solute/kg solvent

There are tables of ebullioscopic constants or boiling point elevation constants (Kb) for common solvents.

SolventNormal Boiling Point (oC)Kb (oC/m)
carbon disulfide46.232.35
carbon tetrachloride76.754.48
ethyl ether34.551.824
Boiling point elevation constants

Vapor Pressure Lowering

Vapor pressure of a liquid is the pressure exerted by its vapor phase when condensation and vaporization occur at equal rate (are at equilibrium). The vapor pressure of a solution is always lower than the vapor pressure of the pure solvent.

The way this works is that the solute ions or molecules reduce the surface area of the solvent molecules exposed to the environment. So, the rate of solvent vaporization decreases. The rate of condensation is not affected by the solute, so the new equilibrium has fewer solvent molecules in the vapor phase. Entropy also plays a role. The solute particles stabilize the solvent molecules, stabilizing them so they are less likely to vaporize.

Raoult’s law describes the relationship between vapor pressure and the concentrations of the components of a solution:



PA is the partial pressure exerted by component A of the solution
PA* is the vapor pressure of pure A
XA is the mole fraction of A

For a nonvolatile substance, the vapor pressure is only due to the solvent. The equation becomes:

Psolution = XsolventPsolvent*

Osmotic Pressure

Osmotic pressure is the pressure required to stop a solvent from flowing across a semipermeable membrane. The osmotic pressure of a solution is proportional to the molar concentration of the solute. So, the more solute dissolved in the solvent, the higher the osmotic pressure of the solution.

The van’t Hoff equation describes the relationship between osmotic pressure and solute concentration:

Π = icRT


Π is osmotic pressure
i is the van’t Hoff index
c is the molar concentration of solute
R is the ideal gas constant
T is temperature in Kelvin

Ostwalt and the History of Colligative Properties

Chemist and philosopher Friedrich Wilhelm Ostwald introduced the concept of colligative properties in 1891. The word “colligative” comes from the Latin word colligatus (“bound together”), referring to the way solvent properties are bound to solute concentration in a solution. Ostwald actually proposed three categories of solute properties:

  1. Colligative properties are properties that only depend on solute concentration and temperature. They are independent of the nature of the solute particles.
  2. Additive properties are the sum of the properties of constituent particles and depend on solute chemical composition. Mass is an example of an additive property.
  3. Constitutional properties depend on the molecular structure of a solute.


  • Laidler, K.J.; Meiser, J.L. (1982). Physical Chemistry. Benjamin/Cummings. ISBN 978-0618123414.
  • McQuarrie, Donald; et al. (2011). General Chemistry. University Science Books. ISBN 978-1-89138-960-3.
  • Tro, Nivaldo J. (2018). Chemistry: Structure and Properties (2nd ed.). Pearson Education. ISBN 978-0-134-52822-9.

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