Boiling & Freezing Point
Boiling-point elevation and freezing-point depression of a solution from its molality, solvent, and van't Hoff factor, using tabulated Kb and Kf constants.
How to use this tool
See how much a dissolved solute raises a solvent's boiling point and lowers its freezing point. Use it for antifreeze, salting roads, or predicting where a solution will boil or freeze.
What to enter
- Solvent: pick from the list; each carries its own Kb/Kf constants and normal boiling/freezing points.
- Molality: moles of solute per kilogram of solvent (not per litre). A typical lab value is 1.0 mol/kg.
- van't Hoff factor i: how many particles each formula unit splits into: 1 for sugar, 2 for NaCl, 3 for CaCl₂.
Reading the result
You get the solution's new boiling point (raised by ΔTb) and new freezing point (lowered by ΔTf). The shifts scale with both molality and i, so a salt that dissociates into more ions moves the points further.
Worked example
1.0 mol/kg of a non-electrolyte (i = 1) in water boils at about 100.51 °C and freezes at about −1.86 °C.
Solution
Result
Both shifts scale with molality × i, so a solute that splits into more ions moves the points further. The new boiling point is the solvent's normal bp raised by ΔTb; the new freezing point is its normal fp lowered by ΔTf. Kb and Kf are fixed properties of the solvent, not the solute.
Methodology
Boiling-point elevation ΔTb = i·Kb·m and freezing-point depression ΔTf = i·Kf·m, where m is molality, i the van't Hoff factor, and Kb/Kf the solvent's molal constants. New boiling point = normal bp + ΔTb; new freezing point = normal fp − ΔTf.
Sources
- Kb, Kf and normal bp/fp from the CRC Handbook (97th ed.).
Known limits
- i is the ideal (fully dissociated) value; real electrolytes show a slightly lower effective i from ion pairing.
- Valid for dilute solutions; high molality deviates from linearity.