Raoult's law
/ rah-OOLZ law /
Leave a glass of water out and some of it quietly escapes as vapour — that escaping tendency is the water's vapour pressure. Now dissolve sugar in it. The sugar molecules sit at the surface and partly block water molecules from leaving, so fewer escape. Raoult's law captures this: dissolving something in a liquid lowers how readily the liquid evaporates.
Stated cleanly, Raoult's law says that in an ideal solution the vapour pressure contributed by each component equals that pure liquid's own vapour pressure multiplied by its mole fraction in the mixture. So if water is only 90% of the molecules (mole fraction 0.9), it produces 90% of the vapour pressure it would on its own. The more solute you add, the lower the solvent's vapour pressure falls.
This simple line is the foundation under several everyday effects: why salt water boils a little hotter and freezes a little colder, and how distillation separates liquids by their differing volatilities. Real solutions deviate from the law when their molecules attract or repel each other unevenly, but for dilute solutions Raoult's law is an excellent guide.
Pure water at 100°C has a vapour pressure of 1 atm. Make the water 95% of the molecules by dissolving a non-volatile solute, and its vapour pressure drops to about 0.95 atm — so it must be heated past 100°C to boil.
Solvent vapour pressure = (pure value) × (solvent mole fraction).
Raoult's law is the model for the solvent (the major component) and for ideal solutions; Henry's law is its counterpart for a dilute solute. Both describe vapour pressure over a solution, but they apply at opposite ends — solvent-rich versus solute-dilute.