🧊 Freezing Point Depression & Boiling Point Elevation

About this simulation

Dissolving a solute into a solvent disrupts the solvent's ability to form an ordered crystal lattice and lowers its vapour pressure. Both effects are colligative — they depend only on how many solute particles are present, not on what the solute is. The result: the freezing point drops and the boiling point rises, both in proportion to the solute's molality.

How it works

• ΔT_f = i·K_f·m — freezing point depression equals the van 't Hoff factor times the cryoscopic constant times molality.
• ΔT_b = i·K_b·m — boiling point elevation uses the ebullioscopic constant instead.
• Each solvent has its own K_f and K_b, set by how strongly its molecules interact.
• The van 't Hoff factor i accounts for dissociation — NaCl (i≈2) or CaCl₂ (i≈3) shift temperatures more per mole than a non-dissociating sugar (i≈1).

Controls

• Solute particles — how many solute units are dissolved in the beaker (drives molality).
• Van 't Hoff factor — how many ions each dissolved unit splits into.
• Solvent — swaps K_f / K_b and the pure-solvent reference points.
• Ambient temperature — where the beaker currently sits; crystals form below the solution's freezing point, bubbles appear above its boiling point.
• Stir solution — toggles particle motion. Drag to rotate the view, scroll to zoom.

Did you know?

This is exactly why salt is scattered on icy roads and why antifreeze (ethylene glycol) is added to car radiators — the dissolved particles depress the freezing point below the ambient temperature and, in the radiator's case, also raise the boiling point of the coolant.