What are colligative properties?

Colligative properties are physical properties of solutions that depend on the concentration of solute particles, not on the type of solute. The four main ones are boiling-point elevation, freezing-point depression, osmotic pressure, and vapor-pressure lowering.

What is the van't Hoff factor i?

The van't Hoff factor accounts for ionic dissociation. For non-electrolytes (sugar, urea) i = 1. For strong electrolytes: NaCl or KCl → i = 2; CaCl₂ or MgCl₂ → i = 3; AlCl₃ → i = 4. Weak electrolytes have a fractional i between 1 and the number of ions.

Why does osmotic pressure use molarity instead of molality?

The van't Hoff equation for osmotic pressure (π = iMRT) uses molar concentration M because it relates to the number of particles per unit volume of solution, which is what drives pressure across a semipermeable membrane. Boiling and freezing calculations use molality because those depend on solvent-mass ratios.

How is vapor-pressure lowering calculated?

By Raoult's Law: ΔP = P° × Xsolute, where P° is the pure solvent's vapor pressure at the given temperature and Xsolute is the mole fraction of the solute. For dilute solutions, Xsolute ≈ m·M_solvent / 1000, where M_solvent is the solvent molar mass.

Can I use this for mixed solvents or multiple solutes?

No. These formulas assume an ideal dilute solution with a single solute in a single solvent. For concentrated or non-ideal solutions, activity coefficients are needed; those require more specialized software.

Colligative Properties Calculator

Enter solvent, molality, and van't Hoff factor to compute boiling-point elevation, freezing-point depression, osmotic pressure, and vapor-pressure lowering in one step.

Solution Parameters

Solvent

Molality, m (mol/kg)

van't Hoff factor, i

Non-electrolyte = 1 · dissociation may give i > 1 · association may give i < 1

Temperature (°C) (for osmotic pressure)

Enter solution parameters and click Calculate to see all four colligative properties.

Solvent: Kb: °C·kg/mol Kf: °C·kg/mol BP₀: °C FP₀: °C

Boiling-Point Elevation

ΔT_b = K_b × m × i

ΔT_b

New boiling point

Freezing-Point Depression

ΔT_f = K_f × m × i

ΔT_f

New freezing point

Osmotic Pressure

π = i × M × R × T

π (atm)

π (kPa)

Approximated with M ≈ m·ρ; assumes dilute aqueous solution density ≈ 1 kg/L

Vapor-Pressure Lowering

ΔP = P° × X_solute

X_solute (mole fraction)

ΔP (mmHg)

New vapor pressure

Inputs used: m = mol/kg, i = , T = °C (R = 0.08206 L·atm/mol·K)

Summary

Colligative properties depend only on the number of solute particles in solution, not their identity. This calculator covers all four: boiling-point elevation (ΔTb = Kb·m·i), freezing-point depression (ΔTf = Kf·m·i), osmotic pressure (π = iMRT), and vapor-pressure lowering (ΔP = P°·Xsolute). Select a common solvent to auto-fill constants, enter molality and the van't Hoff factor, and all four results update instantly.

How it works

Select a solvent from the dropdown or enter custom constants manually.
Enter the molality (moles of solute per kilogram of solvent).
Enter the van't Hoff factor i (1 for non-electrolytes; 2 for NaCl, HCl; 3 for CaCl₂, etc.).
Enter the temperature in Celsius (used for osmotic pressure; default 25 °C).
The calculator applies all four colligative property formulas simultaneously.
Results show the change in each property plus the final value for boiling and freezing points.

Use cases

Chemistry students solving colligative property problems in a single calculation.
Checking lab results against theoretical predictions for all four properties at once.
Comparing how different solvents (water, benzene, cyclohexane) respond to the same solute.
Estimating osmotic pressure for biological or pharmaceutical solutions.
Determining the effect of ionic dissociation (high i) on solution properties.
Teaching general chemistry: seeing all four formulas side by side reinforces the concept.
Quickly finding the new boiling and freezing points after dissolving a salt in water.

Frequently Asked Questions

Last updated: 2026-09-30 · Reviewed by Nham Vu