What it is and when to use it
Freezing point depression is the drop in the temperature at which a solvent freezes when something is dissolved in it. It is a colligative property: the size of the drop depends on how many solute particles are dissolved, not on what they are. It explains why salt on roads melts ice, why antifreeze protects engines, and how chemists estimate molar mass from a measured freezing point.
Use this calculator for general chemistry problems and lab work where you know the molality of a solution and the solvent's cryoscopic constant. It returns the size of the drop and the new freezing point. It assumes an ideal dilute solution, so it works best at low concentrations.
The formula
The calculator uses ΔTf = i × Kf × m and then new freezing point = pure solvent freezing point − ΔTf.
- ΔTf: freezing point depression in °C (equal in kelvin).
- i: the van 't Hoff factor, 1 for a non-electrolyte and about 2 for NaCl. It defaults to 1 here.
- Kf: the solvent's molal freezing point depression constant. Common values are 1.86 °C·kg/mol for water, 5.12 for benzene and 3.90 for acetic acid.
- m: molality, moles of solute per kilogram of solvent.
Worked example: 0.50 m sugar in water
Sucrose does not dissociate, so i = 1. Water has Kf = 1.86 and freezes at 0 °C.
ΔTf = 1 × 1.86 × 0.50 = 0.93 °C. New freezing point = 0 − 0.93 = −0.93 °C, which is exactly what the calculator displays. As a temperature change that is 1.67 °F.
If the solute were NaCl at the same molality, the ideal i = 2 would double the effect to 1.86 °C, giving a freezing point near −1.86 °C. Real NaCl is a little below that ideal value at higher concentrations.
Common mistakes and how to interpret the result
- Using molarity instead of molality. Freezing point depression is defined per kilogram of solvent, so convert if you only know moles per litre.
- Forgetting the van 't Hoff factor. For salts, leaving i at 1 underestimates the drop by half or more. Measured factors are usually slightly below the ideal integer.
- Using the wrong Kf. The constant belongs to the solvent, not the solute, so water's 1.86 cannot be used for a benzene solution.
- Applying it to concentrated solutions. The linear relationship breaks down as concentration rises, and very high concentrations can reach a eutectic point where solute and solvent freeze together.