Vapor Pressure and Raoult's Law

Core Idea

Every liquid exerts a vapor pressure — the pressure of gas-phase molecules in equilibrium with the liquid surface — that increases with temperature and decreases with stronger intermolecular forces. Raoult's law states that the partial vapor pressure of each component in an ideal solution equals the product of its mole fraction and its pure-component vapor pressure: Pᵢ = χᵢPᵢ°. For non-volatile solutes, the total vapor pressure is simply lowered (vapor-pressure lowering is a colligative property). Real solutions show positive deviations (weaker solute-solvent interactions than pure components, higher vapor pressure) or negative deviations (stronger interactions, lower vapor pressure).

How It's Best Learned

Calculate total vapor pressure over two-component solutions by summing partial pressures from Raoult's law. Compare ideal vs actual vapor pressure diagrams to identify positive and negative deviations and connect them to the relative strength of intermolecular forces between components.

Common Misconceptions

• Raoult's law applies strictly to ideal solutions. Most real solutions deviate; Raoult's law is accurate mainly when the components are chemically similar (e.g., benzene and toluene).
• Vapor pressure lowering by a non-volatile solute depends on the number of solute particles, not their identity — an ionic solute that dissociates into two ions has roughly twice the effect of a molecular solute at the same molality.

Explainer

From your study of intermolecular forces, you know that molecules in a liquid are held together by attractive interactions — hydrogen bonds, dipole-dipole forces, or London dispersion forces. Vapor pressure is the pressure exerted by the gas-phase molecules that have escaped from a liquid surface into the space above it. At any temperature, some molecules at the surface have enough kinetic energy to overcome these attractions and enter the gas phase. When the rate of escape equals the rate of return, the system reaches a dynamic equilibrium, and the pressure of the vapor at that point is the liquid's vapor pressure. Liquids with weak intermolecular forces (like diethyl ether) have high vapor pressures because molecules escape easily; liquids with strong forces (like water) have lower vapor pressures.

Raoult's law describes what happens to vapor pressure when you mix two liquids (or dissolve a solute in a solvent). For an ideal solution, the partial vapor pressure of each component equals its mole fraction in the liquid multiplied by its pure-component vapor pressure: Pᵢ = χᵢPᵢ°. The intuition is simple — if only 70% of the surface molecules are solvent A, then A contributes only 70% of the vapor pressure it would have on its own. The total vapor pressure above the solution is the sum of the partial pressures of all volatile components. When you dissolve a non-volatile solute like sugar in water, the solute contributes zero vapor pressure, so the total vapor pressure drops. This vapor-pressure lowering depends only on how many solute particles are present (it is a colligative property), not on what the solute is.

Real solutions rarely obey Raoult's law perfectly, and the deviations reveal important chemistry. Positive deviations occur when solute-solvent interactions are weaker than the pure-component interactions — the molecules "want to escape" more readily than Raoult's law predicts, so the observed vapor pressure is higher than ideal. An example is ethanol mixed with hexane: ethanol loses its hydrogen-bonding partners, and the weaker ethanol-hexane interactions make both components more volatile. Negative deviations occur when solute-solvent interactions are unusually strong — the molecules are held more tightly in solution, and the observed vapor pressure is lower than predicted. Acetone mixed with chloroform is a classic case: a hydrogen bond forms between chloroform's C-H and acetone's C=O that neither pure liquid can form on its own.

Raoult's law is most accurate when the solution components are chemically similar — benzene and toluene, for instance, interact with each other almost identically to how they interact with themselves. As the components become more dissimilar, deviations grow larger. Recognizing whether a given mixture should show positive or negative deviation is a direct application of your intermolecular forces knowledge: compare the strength of the cross-interactions (solute-solvent) to the self-interactions (solute-solute and solvent-solvent). Stronger cross-interactions mean negative deviation; weaker cross-interactions mean positive deviation.