Calculate freezing point depression, molality, or van't Hoff factor using ΔTf = iKf m, with solvent presets, custom constants, and particle-molality output.
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Quick presets
Temperature
Assumptions
Uses the ideal-dilute colligative relation ΔTf = iKf m. That is appropriate for same-solvent freezing
problems and not for concentrated mixtures or systems where the solvent itself changes phase behavior.
Enter a solvent and two known values Choose the solvent constants, then provide the two known colligative-property values to solve the missing term.
Freezing point depression from colligative properties
A freezing point depression calculator uses the colligative relation ΔTf = iKf m to estimate how dissolved particles lower a solvent's freezing point. It is useful for general chemistry problem sets, solution-property checks, and lab review when you need to move between observed freezing point, molality, and particle count.
Why solute lowers the freezing point
Dissolved particles interfere with the orderly formation of the solid solvent phase, so the solution must be cooled further before freezing begins. Like boiling point elevation, this is a colligative effect that depends primarily on particle count in dilute solution.
That is why the page emphasizes particle molality i × m alongside the observed freezing-point shift. More dissolved particles generally mean a larger depression below the pure-solvent freezing point.
Formula used here
This calculator uses the standard dilute-solution relation ΔTf = iKf m. Kf is the solvent's cryoscopic constant, m is molality, and i is the van't Hoff factor. In practice, the calculator applies this freezing point depression from colligative properties relationship to the user inputs, keeps the units and assumptions consistent, and then surfaces the supporting context needed to interpret the output responsibly.
Subtracts the freezing-point shift from the pure solvent freezing point.
Worked example
For a 1.00 m nonelectrolyte in water, Kf = 1.86 °C·kg/mol and i ≈ 1, so the freezing point depression is 1.86 °C. The solution therefore freezes near -1.86 °C instead of 0 °C.
If an electrolyte contributes more dissolved particles, the shift is larger at the same formal molality. That is why ion-producing solutes are especially effective at depressing the freezing point.
Frequently asked questions
Why must the solution freezing point be below the pure solvent?
Because freezing point depression lowers the freezing point relative to the pure solvent. If the observed temperature is above the pure-solvent value, the standard cryoscopic relation does not describe that state.
Can I use this for antifreeze mixtures or concentrated brines?
Only as a first-pass estimate. Real antifreeze and concentrated salt systems can deviate from the ideal dilute equation and may require experimental property data or a more complete phase-behaviour model.
Why does the calculator ask for van't Hoff factor?
The factor i estimates the effective number of dissolved particles. It lets the page adapt the standard formula for nonelectrolytes and electrolytes without hard-coding a dissociation model.
