Raoult's law is the simplest quantitative link between a solution's composition and how readily it evaporates: the vapor pressure of a component above a solution is just its mole fraction times its vapor pressure as a pure substance. It explains why dissolving something in a liquid always suppresses evaporation, and it's the foundation for distillation, humidity calculations, and every other colligative property.

How the formula works

In a pure liquid, molecules at the surface constantly escape into the vapor phase and re-condense, reaching an equilibrium vapor pressure P°. Add a solute, and some of the surface is now occupied by solute particles instead of solvent — fewer solvent molecules are available to escape at any instant, so the vapor pressure drops in direct proportion to how much of the solution is still solvent. That proportion is exactly the solvent's mole fraction, which is why P = X_solvent·P° holds so cleanly for ideal solutions.

The same logic extends to mixtures where both components are volatile: each component's presence at the surface is proportional to its own mole fraction, so each contributes its own mole-fraction-weighted vapor pressure, and the total is just the sum of both contributions.

Inputs and what they mean

Mole fraction must be a value between 0 and 1, and for a two-component solution the solvent and solute mole fractions always add up to exactly 1 — that's what lets the calculator derive one from the other automatically on the VP Lowering tab. Pure vapor pressure (P°) is temperature-dependent, so make sure the value you enter matches the temperature you care about; water's vapor pressure is about 23.8 mmHg at 25°C but roughly 760 mmHg at its 100°C boiling point. For the Two Components tab, both mole fractions must sum to 1 since they describe the same mixture.

Limits and edge cases

Raoult's law is exact only for an ideal solution, where the forces between different molecules are essentially identical to the forces between like molecules — chemically similar liquids like benzene and toluene come very close. Real solutions with strong specific interactions (hydrogen bonding, ion-dipole forces) show positive or negative deviations from the ideal prediction, especially at higher concentrations. The basic single-component formula also assumes the solute itself is non-volatile; if the solute evaporates too, you need the two-component form instead, since it also contributes its own vapor pressure to the total.