What Are Colligative Properties? Definition and Examples

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. For equal masses of solute, substances with lower molar mass generally produce larger colligative effects because they produce more dissolved particles.

Key Takeaways: Colligative Properties

• Colligative properties depend on the number of dissolved particles, not their identity.
• The four colligative properties are freezing point depression, boiling point elevation, vapor pressure lowering, and osmotic pressure.
• Electrolytes have stronger effects because they dissociate into multiple ions.
• Colligative properties are strongest in dilute, ideal solutions.
• These properties explain antifreeze, road salt, food preservation, desalination, and many biological processes.

Why Colligative Properties Matter

Colligative properties are important because they explain and predict the behavior of many real-world solutions. These effects play major roles in chemistry, biology, medicine, environmental science, and engineering.

One familiar example is road salt. Salt lowers the freezing point of water, preventing ice from forming as easily on roads and sidewalks during winter. Calcium chloride works especially well because it produces more dissolved ions than sodium chloride.

Automotive antifreeze also relies on colligative properties. Ethylene glycol lowers the freezing point and raises the boiling point of coolant mixtures, helping engines operate safely over a wide temperature range.

Colligative properties are essential in biology and medicine. Cells depend on osmotic pressure to regulate water balance. Intravenous (IV) fluids must have the correct concentration so red blood cells neither shrink nor burst. Kidney dialysis, food preservation, desalination, and reverse osmosis water purification also depend on osmotic effects.

Chemists use colligative properties to determine the molar masses of unknown compounds. Measuring freezing point depression or boiling point elevation provides information about the number of dissolved particles in a solution.

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 the number of dissolved particles, 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 = iK_fm

where:

ΔT = Change in temperature in °C
i = van ‘t Hoff factor
K_f = 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 (K_f) for common solvents.

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 = iK_bm

where:

i = van’t Hoff factor
K_b = 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 (K_b) for common solvents.

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:

P_A = X_AP_A*

where:

P_A is the partial pressure exerted by component A of the solution
P_A* is the vapor pressure of pure A
X_A is the mole fraction of A

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

P_solution = X_solventP_solvent*

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

where

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.

van ’t Hoff Factor Explained

The van ’t Hoff factor ($i$i) represents the number of dissolved particles produced by a solute in solution. Because colligative properties depend on the number of particles, substances that produce more particles cause larger effects.

Nonelectrolytes dissolve without breaking apart into ions. For example, glucose dissolves as intact molecules:
$$\text{C}_6\text{H}_{12}\text{O}_6 \rightarrow \text{C}_6\text{H}_{12}\text{O}_6$$

Glucose therefore has a van ’t Hoff factor of approximately:$$i = 1$$

Electrolytes dissociate into ions when they dissolve. Sodium chloride separates into two ions:
$$\text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^-$$

So sodium chloride has an ideal van ’t Hoff factor of:$$i = 2$$

Calcium chloride dissociates into three ions:$$\text{CaCl}_2 \rightarrow \text{Ca}^{2+} + 2\text{Cl}^-$$

Thus:$$i = 3$$

Because calcium chloride produces more dissolved particles, it causes larger freezing point depression and boiling point elevation effects than sodium chloride at the same concentration.

Real solutions often deviate from ideal behavior because ions interact with one another in solution. Ion pairing and incomplete dissociation can reduce the experimental van ’t Hoff factor below the theoretical value.

Colligative Properties vs Noncolligative Properties

Colligative properties depend only on the number of dissolved particles in a solution, not on the chemical identity of those particles. In contrast, noncolligative properties depend on the nature and identity of the substances present.

For example, solutions containing equal numbers of dissolved glucose molecules and dissolved urea molecules produce similar colligative effects because they contain the same number of particles. However, the two solutions may differ in density, viscosity, conductivity, or color because those properties depend on chemical composition.

Requirements for Ideal Colligative Behavior

The simple equations used for colligative properties work best for ideal solutions. Real solutions sometimes deviate from these predictions because of intermolecular interactions and incomplete dissociation.

Ideal colligative behavior generally assumes:

• The solution is dilute.
• The solute is nonvolatile.
• Solute particles do not strongly interact with one another.
• The solvent behaves ideally.
• Electrolytes completely dissociate into ions.

Real solutions may show smaller or larger effects than predicted. Strong ionic interactions, hydrogen bonding, or high concentrations often produce nonideal behavior. For this reason, concentrated electrolyte solutions frequently require more advanced calculations.

Biological Importance of Colligative Properties

Colligative properties are essential for living organisms because cells constantly regulate water balance and dissolved particle concentrations.

Osmosis controls the movement of water across cell membranes. If a cell is placed in a solution with too many dissolved particles (a hypertonic solution), water leaves the cell and the cell shrinks. In a hypotonic solution, water enters the cell and may cause it to swell or burst.

Human blood and body fluids maintain carefully controlled osmotic concentrations. Medical IV solutions must closely match blood osmolarity to avoid damaging cells. The kidneys also rely on osmotic pressure to filter blood and regulate water balance.

Many organisms use colligative properties for survival. Fish living in polar environments produce dissolved compounds that lower the freezing point of body fluids. Plants rely on osmotic pressure to maintain turgor pressure and structural support.

Colligative effects also help preserve food. Salt and sugar reduce the availability of water for microbial growth, slowing spoilage.

Determining Molar Mass Using Colligative Properties

Chemists use colligative properties to determine the molar masses of unknown substances. Because these properties depend on the number of dissolved particles, measuring the magnitude of a colligative effect provides information about how many solute particles are present.

Freezing point depression and boiling point elevation are especially useful for this purpose. In a typical experiment, a known mass of solute dissolves in a known mass of solvent. Measuring the temperature change allows calculation of the solute molality and, ultimately, the molar mass of the unknown compound.

This method is particularly useful for polymers and large organic molecules whose molar masses may be difficult to determine by other techniques.

Common Misconceptions About Colligative Properties

A common misconception is that colligative properties depend on the type of dissolved substance. In reality, they mainly depend on the number of dissolved particles.

Another misconception is that all dissolved substances produce equal effects at equal masses. However, substances with lower molar masses or greater ion dissociation produce more particles and therefore larger colligative effects.

Students also sometimes assume colligative equations work perfectly for all solutions. In practice, the equations are most accurate for dilute ideal solutions. Concentrated solutions and strong electrolytes often show nonideal behavior.

Finally, boiling point elevation and freezing point depression are sometimes assumed to be large effects. In many everyday solutions, the temperature changes are actually relatively small unless the solute concentration is high.

References

• 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.