If I leave a glass of water out, why do only the surface molecules vaporize? If I leave a glass of water out on the counter, some of the water turns into vapor. I've read that this is because the water molecules crash into each other like billiard balls and eventually some of the molecules at the surface acquire enough kinetic energy that they no longer stay a liquid. They become vapor. 
Why is it only the molecules on the surface that become vapor? Why not the molecules in the middle of the glass of water? After all, they too are crashing into each other. 
If I put a heating element under the container and increase the average kinetic energy in the water molecules to the point that my thermometer reads ~100°C, the molecules in the middle of the glass do turn into vapor. Why doesn't this happen even without applying the heat, like it does to the surface molecules?
 A: From a thermodynamic point of view, at fixed pressure, the vaporization takes place when the temperature exceeds the temperature of change of state $ Tc (P ) $
Within the liquid, the pressure that is to be taken into account is the hydrostatic pressure. This pressure is a little greater than 1 bar and the associated vaporization temperature is 100 ° C.
On the surface (thickness of some mean free path), the environment of the molecules is different. the pressure to be taken into account is the partial pressure of water vapor, which is related to the moisture content of the air. If the humidity is less than 100%, this pressure is well below 1 bar and evaporation takes place at a much lower temperature.
A: There's a fundamental difference between a liquid changing to a gas at the surface vs. in the bulk: the formation of new surface area, which costs energy. 
Net evaporation from the surface is spontaneous whenever the relative humidity is less than 100% because energy fluctuations enable surface molecules to detach into the gas phase, as you describe. Here, the total surface area doesn't change. In contrast, gas formation within the bulk (i.e., the formation of an evaporative bubble) requires the formation of an additional liquid-gas interface, which carries an energy cost because bonds tend to be unsatisfied at interfaces.
In fact, the energy penalty from having to form the surface of a nucleating vapor bubble is so important that we generally can only achieve boiling by  (1) providing an existing surface or (2) waiting a large amount of time for a large gas cluster to randomly assemble or (3) superheating the liquid to increase the driving force for boiling (or a combination of these).
A: The water molecules in the liquid attract each other. Their thermal velocity distribution allows some molecules to be fast enough to overcome this attraction. If it happens to a molecule at the surface to be kicked by such a fast molecule, it may be kicked with an impulse stronger than the attractive forces, and therefore leave the liquid. The same kick inside the liquid would be passed on to other molecules very efficiently.
If a gas bubble forms inside the liquid, the reduced attraction between the molecules in the gas is part of the energy penalty to be paid for the bubble formation in the form of heat from an external heat source.
