vapor pressure
Leave a glass of water on the table and it slowly disappears, even though it never boils. The most energetic molecules at the surface keep escaping into the air as vapor. Seal that glass instead, and a balance sets in: molecules leave the liquid as fast as others rejoin it. The push exerted by that escaped vapor pressing back on everything is the vapor pressure.
More precisely, vapor pressure is the pressure of a vapor sitting in equilibrium with its own liquid (or solid) at a given temperature. It measures how eager a substance is to evaporate. Volatile liquids like alcohol, ether, or gasoline have high vapor pressures; oils and honey have very low ones. Crucially, vapor pressure depends only on the substance and the temperature — heat the liquid and far more molecules have the energy to escape, so the vapor pressure climbs steeply.
Vapor pressure quietly governs everyday life: it sets how fast puddles dry, how strongly perfume reaches your nose, and how a fuel evaporates in an engine. It also defines boiling — a liquid boils precisely when its vapor pressure rises to match the surrounding air pressure, which is why water boils cooler up a mountain where the air pushes down less hard.
At 20 °C water's vapor pressure is about 2.3 kPa; at 100 °C it has soared to roughly 101 kPa — exactly the everyday air pressure — and the water boils.
A liquid boils when its vapor pressure equals the surrounding pressure.
Vapor pressure rises with temperature, not in a straight line but exponentially — the Clausius–Clapeyron equation captures this steep curve. Dissolving a non-volatile solute (like salt) lowers a liquid's vapor pressure, the basis of Raoult's law.