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# Water

issy | 02:49 Sat 26th Mar 2005 | Science
If the boiling point of water is 100 degrees Celsius, why does it evaporate at room temperature? For instance, if I leave a drop of water on my coffee table before I go to bed, why is it not there in the morning?

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No best answer has yet been selected by issy. Once a best answer has been selected, it will be shown here.

All liquids constantly evaporate.  The hotter the liquid is, the faster it evaporates of course.  The "boiling point" is simply the temperature at which a certain substance can no longer exist as a liquid, and must become a gas.  For water this is 100�C when the water is evaporating into air of "standard" pressure and humidity.  Under these conditions you can't get 101�C water, for example.  The more heat you apply, the faster the water will evaporate, but a thermometer put in the water would show a constant 100�C.

This can be explained when you consider the individual water molecules.  The molecules are always moving and bouncing off each other - very quickly, but some moving faster than others.  The hotter the water, the faster the average speed of the molecules.  There is also an overall attractive force holding the water molecules close together (holding them in liquid form).  When one of the faster water molecules is moving towards the surface of the water, it may have enough speed to overcome the attraction of its neighbouring water molecules and escape (evaporate) into the air.  Similarly, molecules of water in the air can enter a pool of liquid water.  Under normal conditions there will be more water molecules leaving a pool of liquid water than entering it from the air, so the liquid slowly evaporates.

You'll also notice that, because it is always the faster/more energetic molecules that evaporate first, there is a resulting decrease in the average speed and energy of the molecules remaining in the liquid.  This corresponds to an overall decrease in the temperature of the liquid.

Now I read it, I realise that my answer is actually quite vague, inaccurate and has gaping holes.  If anyone would care to offer the real explanation that would be marvellous.

Your question confuses the boiling point of water with evaporation.  The two terms have only a relative relationship. Evaporation is a surface effect.  The temperature of a body of water represents an AVERAGE value of the kinetic energy of the water molecules in the water.  Some of those molecules have a "higher than average" kinetic energy and break free of the surface to become water vapor.  As this happens over time, the remaining body of water becomes slightly cooler (as water
molecules with higher kinetic energy have "left").  That is why sweat evaporating from your body provides a mechanism to cool you off.

This occurs at any temperature of water between the melting point and the
boiling point -- with more evaporation occurring with a body of water that has a higher average temperature (that is, closer to the boiling point).  It can even happen below the freezing point (via sublimation -- when water goes directly from a solid (ice) to water vapor) and this explains why the ice cubes in your freezer sometimes seem to "evaporate".

All of this is assuming that you are at standard pressure of 1 atmosphere.
As the atmospheric pressure changes, water will boil and freeze at different temperatures. (With help from Newton Archives).
The way I understand it, when any liquid substance is in equilibrium with  a gas phase (in this case, air), a certain amount of the substance is also present in gas phase.  In fact there is always an exchange between water vapor and liquid water taking place ( in both directions).  At equilibrium, the amount of water in the gas phase is called the partial pressure of water.  At a given temperature and pressure in a closed system that is at equilibrium, the rate of water vapor changing to liquid water (condensation) equals the rate at which liquid water transforms to water vapor (evaporation), resulting in a constant partial pressure of water vapor and a constant volume of liquid water.  In this case (equilibrium), the air is saturated with water vapor (100% relative hunidity).  However, if the system is not in equilibrium and the air is not saturated with water vapor, the rate of evaporation is greater than that of condensation, resulting in the  reduction of the volume of liquid water (the disapearance of your water drop on the coffee table).  This partial pressure of water is dependent on temperature (increases with increasing temperature), and it just so happens that at 100 deg C, the partial pressure of water is 1 atmosphere.   So if the over all pressure of the system is 1 atmosphere, liquid water will not be stable and all  of the water will be transformed to water vapor, hence boiling.  So at 1 atmosphere pressure, at 100 deg C, the liquid water will remain at this temperature until all the water has transformed to water vapor.  This also explains why water boils at a lower temperature at higher altitudes where atmospheric pressure is less than 1 atmosphere.

Continued-

So basically water boils when its vapor pressure is equal to the total pressure of the system.  As for the cooling effect of evaporation mentioned above, this is due to evaporation being an endothermic reaction (heat is added for the reaction to take place).  Heat is taken from the system for the reaction (phase change ) to occur, resulting in a reduction in the temperature of the system.

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