Freezing-Point Depression

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 ΔT_f is proportional to molality through the cryoscopic constant K_f: 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 K_f 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.