# How can water evaporate at room temperature? [duplicate]

Boiling point of water is 100 degree Celsius. The temperature at which water in liquid form is converted into gaseous form. Then how it possible for water to evaporate at room temperature?

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## marked as duplicate by Brandon Enright, jinawee, Jim, Kyle Kanos, Qmechanic♦Apr 7 '14 at 12:38

The water molecules at the surface of the water don't need as much energy to evaporate. – PPG Apr 4 '14 at 10:53
Boiling is different than simple evaporation.. Water vapours can exist at zero degrees Celsius. – Evil Angel Apr 4 '14 at 10:56
pressure can be low enough for water to evap at room temperature. – geoff Apr 4 '14 at 23:03
Possible duplicate of: physics.stackexchange.com/q/10470 – mikhailcazi Apr 5 '14 at 7:11

Think of temperature as average kinetic energy of the water molecules. While the average molecule doesn't have enough energy to break the inter-molecular bonds, a non-average molecule does.

Water is a liquid because the dipole attraction between polar water molecules makes them stick together. At standard atmospheric pressure (acting somewhat like a vice), you need a comparatively large temperature of 100°C (translating to a high average energy distributed among the micsroscopic degrees of freedom, most relevantly the kinetic ones) for water molecules to break free in bulk, creating bubbles of water vapour within the liquid.

However, at the surface of the liquid, lone molecules may end up getting enough kinetic energy to break free due to the random nature of molecular motion at basically any temperature. On the flip side, water molecules in the atmosphere may enter the liquid at the surface as well, which is measured by equilibrium vapour pressure.

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This is basically the answer I would give. Christoph, would you consider fleshing this out a bit? At the moment it's so brief only people who already know the answer will be able to see what you're getting at. – John Rennie Apr 4 '14 at 11:18
Got a spark but not clear enough. like @JohnRennie said. – tollin Apr 4 '14 at 11:28
@JohnRennie: I expanded the answer and made it community wiki - feel free to improve upon it as yous ee fit... – Christoph Apr 4 '14 at 12:06
Is the "flip side" the reason why evaporation can't be considered a Maxwell's demon of sorts? – Michael Apr 4 '14 at 18:59

Temperature is a measure for how much kinetic energy the molecules in a substance have. If the temperature is high, they are moving pretty fast, if the temperature is low, they are moving a lot slower. If molecules are moving slow, they bundle up and you get a solid. Once you heat it up a bit, the substance starts to become liquid. When you heat it up even more, the molecules will start to move so fast they will spread out into the entire space (gas).

However, this is all averages. In a liquid all molecules are moving, some faster than others. If a molecule happens to break through the 'surface' of the water, it'll have escaped the inter-molecular forces holding the water together and it'll be evaporated. This can also happen with solids, there it is called sublimation.

If you're heating up water, you're adding energy so this process will start to go faster. Then at boiling point, you'll reach the point where molecules will want to start moving so fast they start to form gas bubbles inside the liquid.

disclaimer: this is just what I remember from high school.

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Good memory there. – BMS Apr 5 '14 at 1:31

Imagine spinning a roulette wheel, but instead of dropping in one ball, you drop in 100. They all rattle around at different speeds, like the molecules in water. You can cool them down by spinning the wheel slower, so they bounce about less; heat them up by spinning faster so they bounce more; you can freeze them by stopping the wheel and waiting till they're all stationary; and you can boil them by spinning the wheel so fast that they all fly out of the top.

Now pick up all the balls and throw them back in with the wheel spinning at a moderate speed. If you watch for a while you'll see that although the average speed of the balls is below the "boiling point" where they all fly out the top, every now and again one ball will ricochet off another with enough force to send it flying out of the wheel. If you watch for long enough eventually all the balls will be gone. Your balls just evaporated.

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good thinking. This analogy really helps. – tollin Apr 5 '14 at 8:34
This analogy should be summed up and added to the most visible one. It basically explains that's because sometimes molecules ricochet off other molecules. Thanks for the example. – Carles Alcolea Jul 19 '15 at 15:28

The boiling temperature of a liquid is not the temperature at which it can enter the gaseous state. Rather, it is the temperature at which the saturation vapor pressure $e_s$ is equal to ambient atmospheric pressure. This is why, for example, water boils at lower temperatures at higher altitudes.

Furthermore, water is always evaporating. It is also always condensing. You can picture a cup of liquid sitting in a room. The evaporation rate will be driven by the saturation vapor pressure $e_s$ calculated with the temperature of the liquid water. The condensation rate will be determined by the vapor pressure $e$ of the water vapor in the air. Typically when someone says a liquid is evaporating (or condensing) they are referring to net evaporation. Net evaporation is when evaporation exceeds condensation (liquid is decreasing) and net condensation is when condensation exceeds evaporation (liquid is increasing).

You will typically hear humidity expressed as "relative humidity", which is the ratio of vapor pressure $e$ to saturation vapor pressure $e_s$ and thus when a the relative humidity $\dfrac{e}{e_s} < 1$, net evaporation is occurring and when $\dfrac{e}{e_s} > 1$ net condensation is occurring. In both cases however, evaporation is constantly occurring.

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At the boiling point the gas is produced inside the liquid, but at the surface you constantly have molecules going in and out. If the environment is kept quite dry, then few molecules will come back in with respect to the ones that leave. Off course the higher the temperature, the easier will be for a molecule to get enough energy to break free, but this can happen at any temperature - at the surface.

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It can also be understud by the idea of partial pressure. Water will evaporate in an atmosphere until its partial pressure has reached the vapour pressure given for the ambient temperarture (relative humidity of 100%).

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These answers account for the kinetics of the process. Thermodynamics provides an alternate picture, well suited for any question involving phase transitions. For the system of liquid plus empty volume, the free energy can be lowered by trading some enthalpy (to take molecules from the liquid into the gas phase) for the increase in entropy (all the states available in the gas phase). For a given temperature, this balance determines the equilibrium vapor pressure. There will always be some vapor in equilibrium (at all non-zero temperatures).

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All water moleucles contain energy, in accordance with the temperature. Hot water has enough energy to escape the liquid as vapour.

Even though a body of water is below boiling point, the molecules with more energy (relative to the body of water),rise to the top and can escape, as vapour

Tis easy!

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In Science class, my teacher put a flask of water on his worktable, covered it with a glass dome with a rubber rim on it so as to seal the dome to the desk. He then hooked a hose to the dome, the other end to a vacuum device and proceeded to remove the air from the dome creating a very low pressure zone within the dome... as the pressure dropped, the water began to boil - at room temperature. This is why mountaineers have difficulty getting hot water at higher altitudes.

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Welcome to phys SE! Your information is interesting and relevant, but doesn't quite address the question. – innisfree Apr 5 '14 at 14:01