Evaporation is said to cause cooling effect because it absorbs energy from surroundings to change its phase from that of a liquid to gas. I am in doubt as to why would the surroundings be ready to give up energy more than the water molecules. Wouldn't 'lower temperature of the water molecules in comparison to the surroundings' or other such conditions which might involve specific heat capacity or conductivity etc. be mandatory to ensure that the heat is transferred from the surroundings to the water molecules and not the other way round ?
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$\begingroup$ See: physics.stackexchange.com/questions/416161/… $\endgroup$– Andrew SteaneCommented Jun 19, 2019 at 11:53
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1 Answer
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Evaporation requires the liquid molecules to overcome the intermolecular attractive forces and escape to the surroundings. Only the most energetic molecules, near the surface of the liquid, have enough kinetic energy to overcome the attractive forces. As these highly energetic molecules escape from the liquid (i.e., evaporate), the average kinetic energy of the molecules in the liquid decreases and temperature is nothing but a measure of the average kinetic energy of the system. Hence, temperature decreases with evaporation.
Now, for more number of liquid molecules to have energy enough to overcome the intermolecular forces, the average energy of the liquid must be high. The liquid can gain more energy by absorbing energy from its surroundings. But according to the second law of thermodynamics, this is possible only if the temperature of the liquid is lower than that of its surroundings. Such a flow of energy occurs until the liquid and its surroundings are at the same temperatures (or until all the liquid evaporates).
So obviously, evaporative cooling is possible only if the liquid is cooler than its immediate surroundings. A very good example of that would be sweating, in which the sweat absorbs energy from our body and evaporates.
Edit:
When you perspire the sweat will be at the same temperature as that of the body. However as the more energetic molecules escape the liquid, the average energy (and hence temperature of the liquid) decreases and this is turn causes the sweat to absorb more heat from the body.
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$\begingroup$ The average energy need not be high for a large number of liuid molecules to possess the desired energy to overcome attractive forces.The surface molecules just need to attain enough energy which is possible without much of change in the average energy of the whole liquid. That is the reason why evaporation akes place below boiling point. Right? $\endgroup$– ZamCommented Apr 25, 2019 at 6:47
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$\begingroup$ Sweating is not a good example for explanation of this specific question unless a reasoning is given as to why the sweat droplets are at a lower temperarure than the body so as to be able to absorb energy from the same and evaporate. $\endgroup$– ZamCommented Apr 25, 2019 at 6:50
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1$\begingroup$ Yes, you are correct. What I wanted to convey was, if the average energy is higher then more number of molecules would have the energy required to overcome the intermolecular forces. $\endgroup$ Commented Apr 25, 2019 at 6:53
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$\begingroup$ I thought that was explained in the initial part of the answer. Anyway, I have added an explanation. $\endgroup$ Commented Apr 25, 2019 at 7:18
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1$\begingroup$ I said only some molecules possess enough kinetic energy to escape the intermolecular forces. Liquid molecules follow the Maxwell-Boltzmann velocity distribution and there are always some molecules which possess kinetic energy more than the average. You can read more about it here: en.m.wikipedia.org/wiki/Maxwell–Boltzmann_distribution $\endgroup$ Commented May 5, 2019 at 13:11