Last updated: 2026-06-14

Understanding freezing point depression

Dissolving a solute in a liquid lowers the temperature at which it freezes. Like boiling point elevation, this is a colligative property — it depends on the number of dissolved particles, not their identity. It explains why the ocean freezes below 0 °C, why salt clears icy roads, and why antifreeze protects a car radiator.

The freezing point depression formula

$$\Delta T_f = i \cdot K_f \cdot m$$

Symbol	Quantity	Unit
ΔTf	Freezing point depression	°C
i	van't Hoff factor	—
Kf	Cryoscopic constant	°C·kg/mol
m	Molality of solute	mol/kg

The new freezing point is the pure solvent's freezing point minus the depression:

$$T_f = T_f^0 - \Delta T_f$$

Worked example

What is the freezing point of a 1.0 mol/kg solution of table salt (NaCl) in water?

NaCl dissociates into two ions, so i ≈ 2, and for water Kf = 1.86 °C·kg/mol:

$$\Delta T_f = 2 \times 1.86 \times 1.0 = 3.72\ \text{°C}$$

The solution freezes at 0 − 3.72 = −3.72 °C.

Why salt melts ice

Salt scattered on icy roads dissolves into the surface film of liquid water and depresses its freezing point, so the water no longer refreezes and the ice melts. Because the effect scales with particle count, de-icers that release more ions work at lower temperatures:

De-icer	i (ideal)	Practical low temp
NaCl (rock salt)	2	about −9 °C
CaCl₂	3	about −29 °C
MgCl₂	3	about −15 °C

Cryoscopic constants

Each solvent has its own Kf, the freezing-point drop per mol/kg:

Solvent	Kf (°C·kg/mol)	Normal fp (°C)
Water	1.86	0
Benzene	5.12	5.5
Camphor	40	179
Acetic acid	3.90	16.6

Camphor's enormous constant made it the classic solvent for measuring molar masses by freezing-point depression.

Antifreeze

Car coolant mixes ethylene glycol with water to depress the freezing point far below 0 °C, keeping the radiator from freezing and cracking in winter. The same mixture raises the boiling point in summer — the companion boiling point elevation effect — so a single additive protects the engine at both extremes.

Frequently Asked Questions (FAQ)

What is the freezing point depression formula?

Freezing point depression is ΔTf = i × Kf × m, where i is the van’t Hoff factor, Kf is the solvent’s cryoscopic (freezing-point) constant, and m is the molality of solute in mol/kg. The new freezing point is the pure solvent’s freezing point minus ΔTf. For example, 1 mol/kg of sugar in water lowers the freezing point by 1 × 1.86 × 1 = 1.86 °C, to −1.86 °C.

Why does salt melt ice on roads?

Salt dissolves into the thin layer of liquid water on ice and depresses its freezing point, so the water stays liquid below 0 °C and the ice melts. Because the effect depends on the number of particles, salts that dissociate into more ions are more effective — calcium chloride (i ≈ 3) works at lower temperatures than sodium chloride (i ≈ 2). Below about −10 °C ordinary rock salt becomes much less effective.

What is the cryoscopic constant Kf?

The cryoscopic constant Kf is a solvent property giving the freezing-point drop per unit molality. For water Kf = 1.86 °C·kg/mol, so a 1 mol/kg ideal solution freezes 1.86 °C lower. Solvents with large constants, such as camphor (40 °C·kg/mol), show huge freezing-point shifts and were historically used to measure molar masses by freezing-point depression.

How does the van’t Hoff factor affect freezing point?

The van’t Hoff factor i is the number of dissolved particles per formula unit, and freezing point depression scales directly with it. A non-electrolyte like sugar gives i = 1, while NaCl gives i ≈ 2 and CaCl₂ gives i ≈ 3. So at the same molality, CaCl₂ lowers the freezing point roughly three times as much as sugar — which is why ionic de-icers are chosen for cold climates.