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Freezing-Point Depression

Also known as antifreeze effect · colligative freezing

ΔTf=Kf b\Delta T_f = K_f \, b

Worked example: Kf = 1.86, b = 2 mol/kg → dT = 3.72 K — press Try an example to run it live, then adjust anything.

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See also
Freezing-Point Depression with the van 't Hoff Factor — The version that works for salts: road salt and calcium chloride each release more than one particle per formula unit, and this form under-predicts them badly.
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Freezing-Point Depression explained

ΔTfKfb

Dissolved particles get in the way of solvent molecules trying to organize into a crystal, so the solution must be cooled below the pure solvent's freezing point before ice can form. The depression ΔTf is proportional to molality through the cryoscopic constant Kf — for water, 1.86 K·kg/mol, more than three times its boiling-side constant. A 2.0 mol/kg ethylene glycol solution freezes at −3.7 °C, and a fully protected automotive antifreeze mix pushes well below −30 °C by the same mechanism.

The effect is everywhere: road salt melts ice because brine freezes lower than pure water (each mole of CaCl₂ delivers three moles of particles), and hand-cranked ice cream relies on a salt–ice bath dropping below −10 °C. In the lab, camphor's giant Kf of about 40 K·kg/mol made the "Rast method" possible — weighing a solute into molten camphor and reading its molar mass off a melting-point depression measured with an ordinary thermometer.

Freezing-Point Depression formula

ΔTf=Kf b\Delta T_f = K_f \, b
Where
  • ΔTf\Delta T_f= Freezing-point depression (C°)
  • KfK_f= Cryoscopic constant (K·kg/mol)
  • bb= Molality of solution (mol/kg)

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Part of this formula set

Colligative properties

4 formulasBoiling-point elevation, freezing-point depression, osmotic pressure and Raoult's law — the four effects that count particles, not identity.