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The boiling point of water is $\pu{100^\circ C}$. However, we find that even at room temperature, water can evaporate into gas. Therefore, water vapor can exist at temperatures of, say, $\pu{50^\circ C}$.

At what temperature, can you say for certain that all water vapor will turn into liquid? In other words, what is the condensation point of water?

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    $\begingroup$ AFAICT, never: Evaporation occurs at all temperatures. $\endgroup$ – ManishEarth Jun 12 '13 at 5:47
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    $\begingroup$ If I'm not mistaken that is technically not right (although this is more semantics than anything else). Once water has become ice (under normal condition $<0^o$C), the process of ice turning into gas is not evaporation but sublimation. So below freezing there is no evaporation. $\endgroup$ – Michiel Jun 12 '13 at 6:08
  • $\begingroup$ I believe the OP is assuming 1atm, otherwise the temperature values he posted are meaningless. $\endgroup$ – Alex Jun 13 '13 at 0:54
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You're right - water vapor can exist at temperatures such as $\pu{50^\circ C}$. This is a phenomenon known as evaporative cooling, where molecules of water with higher kinetic energies tend to "release" themselves from the system, and as a result, less and less water molecules are held within that system.

Temperature is a sort of proxy for kinetic energy, and vice versa. Reference, also, the Maxwell-Boltzmann distribution, which reveals that at $\pu{100^\circ C}$, for example, not all molecules in the system possess the kinetic energy that they would at $\pu{100^\circ C}$, but rather the average exists at that temperature.

To answer your question, never in a real scenario.

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From a chemical engineering perspective (where we do a lot with steam), the answer is dependent on if you're in a closed system or not. If I put some water in a closed container, it would evaporate only enough such that the gaseous water would reach the vapor pressure at that specific temperature. All temperatures above $\pu{0 K}$ have a non-zero vapor pressure, so you could say that water evaporates at all temperatures, at least for a while, and then it will be in equilibrium with the liquid (or solid) state. However, if you are not in a closed system (say, outside with a cup of water), then the water will continue to evaporate to reach the vapor pressure, but the gas will continue to escape, so the water will never stop evaporating.

The opposite, and the answer to your question, is also true. If you have circumstance when the surrounding pressure due to water vapor is higher than the vapor pressure of water at that temperature (like a rain cloud cooling down), then the water will start to precipitate. Of course, like all equilibrium, even if the current net phase shift is from gas towards liquid water, some of the liquid will still be evaporating - it's just that more gas will be condensing.

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In other words, what is the Condensation Point of water?

The dew point is the point at which liquid water would condense from the atmosphere, and the frost point is the temperature at which ice would form from the atmosphere. Both are functions of temperature, pressure, and the water content of the atmosphere.

At what temperature, can you say for certain that all water vapor will turn into liquid?

At absolute zero there would be no gaseous water.

For temperature above 0 Kelvin and 273.15 Kelvin (0 Celsius) or below then water condensing from the atmosphere would form ice.

For temperature between 0 Celsius and 100 Celsius liquid water would condense.

Above 100 Celsius water wouldn't condense unless the gas phase of the system was under pressure.

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I wasn't sure about this myself! But after thinking about it logically, I realized that when a gas condenses into a liquid at around 50 Degrees, I knew that when a liquid condensed, it was one step below, so when it condensed it would become a solid. When it freezes (condenses) into a solid, the temperature is around freezing (0 Degrees Celcius) and (32 Degrees Farenheit)

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  • $\begingroup$ Could you back this answer up with some sources? $\endgroup$ – M.A.R. ಠ_ಠ Dec 16 '15 at 22:14

protected by Community May 16 '18 at 19:19

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