Evaporation

Evaporation is the change of a liquid into vapour that occurs only at the liquid's exposed surface, at any temperature — it does not require the liquid to reach its boiling point, and no bubbles form inside the liquid.

Why only some particles escape. The particles of a liquid are in ceaseless, random motion, and at any instant they are not all moving at the same speed — some are slow, some are moderate, and a few are moving unusually fast. Interparticle attractive forces normally hold liquid particles together, but a fast-moving particle sitting right at the surface can carry enough kinetic energy to break free of those attractions and escape into the air as a gas particle. Because this "escape club" is limited to the few surface particles with enough energy at that moment, evaporation nibbles away at the liquid slowly and continuously, rather than turning it to vapour all at once the way boiling does.

Why evaporation cools the liquid left behind. Since it is specifically the highest-energy particles that leave, the particles remaining in the liquid have, on average, less energy than before those particles escaped. Average kinetic energy of the particles is what we sense as temperature, so a liquid that keeps losing its most energetic particles keeps losing average energy too — its temperature drops. This is exactly why sweat evaporating from skin feels cooling, and why water stored in a porous earthen pot stays cooler than the surrounding air.

Why some conditions speed it up. Anything that helps more surface particles gain escape energy, or clears escaped vapour away so the liquid keeps losing particles freely, raises the rate of evaporation:

• Higher temperature gives particles more thermal energy on average, so more of them cross the energy threshold needed to escape.
• Larger surface area exposes more particles to the air at once — the same volume of water spread into a thin puddle evaporates faster than it would inside a narrow glass.
• Stronger wind sweeps escaped vapour away from the surface before it can build up, stopping the air just above the liquid from becoming saturated.
• Lower humidity means the surrounding air already holds less vapour, so it can accept more escaping particles before it, too, becomes saturated.

The reasoning behind evaporation, and why it differs from boiling, follows directly from two facts about particles that hold for any liquid: particle motion is random with a spread of energies, and only particles at the exposed surface have a free path into the air above.

Step 1. Picture the liquid's particles as a crowd in constant motion, each held near its neighbours by attractive forces, but none moving at exactly the same speed as another. At any given instant this spread of speeds means a small number of particles are moving unusually fast compared to the average — these are the only candidates for escape.

Step 2. Only particles sitting right at the surface have open air above them; particles deep inside the liquid are surrounded by other particles on all sides and have nowhere to go even if they happen to be fast. So the "escape opportunity" is doubly restricted — a particle needs both enough energy and a surface position at the same moment.

Step 3. When a surface particle's kinetic energy exceeds the pull of the attractive forces binding it to its neighbours, it breaks free and joins the air as vapour. Because this threshold-crossing happens for only a few particles at a time, the liquid loses mass gradually rather than all at once — this is why evaporation is slow and why it can occur at room temperature, far below the liquid's boiling point.

Step 4. Every particle that escapes was, by definition, one of the more energetic ones present. Removing high-energy particles pulls the average energy of the particles left behind downward, and since temperature reflects that average, the remaining liquid cools. Raising the temperature, spreading the liquid over a larger surface, blowing air across it, or reducing the humidity of the surrounding air each increases how many particles cross the escape threshold per second, which is exactly why those four conditions all speed up evaporation.

Question: A little water is spilt on a tabletop at room temperature and disappears within an hour, with no heating involved. On a windy, sunny day, a wet handkerchief hung outdoors dries faster than an identical one hung indoors in still air. Explain both observations using the particulate nature of matter, and say what they have in common.

Step 1. In both cases, the water's particles are in constant motion with a range of kinetic energies, even though nothing is being boiled — this motion exists at room temperature just as it does at any temperature.

Step 2. For the spilt water, the thin layer spread across the tabletop has a large surface area relative to its small volume, so a large fraction of its particles sit at the exposed surface at any moment. Because so many particles have a chance to be among the "fastest, at the surface" group, particles escape steadily and the whole thin layer evaporates within about an hour, without ever reaching 100°C.

Step 3. For the handkerchiefs, the outdoor one benefits from two separate boosts: sunlight raises its temperature, giving its particles more thermal energy so more of them cross the escape threshold, and wind sweeps the escaped vapour away before the air just above the cloth becomes saturated with moisture. The indoor handkerchief, sitting in cooler, still air, evaporates more slowly because escaped vapour lingers near its surface and the air saturates faster.

Step 4. Both scenarios rely on the same underlying process — surface particles escaping when they have enough energy — but each is sped up by a different factor from the list: surface area for the spilt water, and temperature plus air movement for the outdoor handkerchief. Answer: the spilt water evaporates through its large exposed surface area, and the outdoor handkerchief dries fastest because higher temperature and wind both raise its rate of particle escape.

• Thinking evaporation needs heating to start — it happens at any temperature; heating only makes it faster.
• Mixing up evaporation with boiling — evaporation is slow, surface-only, and occurs at any temperature; boiling is fast, involves the whole liquid, and needs the boiling point.
• Believing the escaping particles leave the remaining liquid warmer — the opposite is true, since the highest-energy particles are the ones that leave, so evaporation cools what stays behind.
• Assuming humid air speeds up evaporation — moisture-laden air is already close to saturated, so it accepts fewer new particles and evaporation actually slows down.

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