system
Given the following post, respond as if you're commenting on Reddit:
human
If the boiling point of water is approximately 100° C, why does water evaporate on an average day?

I'm a third year meachanical engineering student with a basic thermodynamic background, so don't hesitate using scientific terms.

I've considered varying pressures, but the few ideas that I've come up with seem far fetched and not too plausible.

The reason it confuses me, is that a substance cannot change phase without passing it's boiling point and staying at that specific temperature for the duration of the condensation or evaporation. If this is true, why does water evaporate even at low temperatures? Is what I have just mentioned only an approximation?
gpt
A liquid exists at equilibrium with its vapour. At any given point, individual molecules near the surface may have enough energy to break free and enter the gas phase, while gas phase molecules may hit the liquid with little enough energy to be absorbed into the liquid.

When the atmospheric vapour pressure is lower than the equilibrium vapour pressure, obviously water will continue to enter the gas phase until either equilibrium is established, or you run out of liquid water.

Keeping in mind that temperature is the *average* temperature of a substance, individual particles may well have enough energy to enter the gas phase. This also explains why evaporation is an endothermic process (aside from the thermodynamic explanation of the latent heat of evaporation). As the most energetic particles are removed from the liquid, the average energy (and therefore temperature) decreases. The cooler water then absorbs heat from its surroundings, and the process continues.

The boiling point of a liquid is simply the temperature at which its vapour pressure (which increases with temperature) is equal to the ambient pressure. Hence why water boils at much lower temperatures at higher altitudes.