Vapor Pressure and Equilibrium I am still fairly new to physics, and my questions might come foolish to the more professional; but these are honest questions that I feel restrained in understanding. I believe I partially got the concept of vapor pressure. If it is the pressure excerted by the liquid particles that have gained the needed kinetic energy to enter a gaseous state (in a vat, implicitly).  I inferred that it could gain this excess kinetic energy by random forces imposed on it by the neighboring particles; but I am yet to understand the concept of equilibrium and how it is indeed related to vapour pressure? I ask this here because I believe I have searched much without this site, and generally all other sites are worded similar to each other letting impass to the solution I need. It has been pointed out to me that the equilibrium state is so reached only when the vapour pressure is equal to the atmospheric pressure. But if the vapour pressure was so defined as above, wouldn't it be one with the atmospheric pressure? Do we consider the particles of the atmosphere individual from those of the vaporized liquid? If so, wouldn't they be coexisting in the same site? Thank you in advance.
 A: It has been pointed out to me that the equilibrium state is so reached only when the vapour pressure is equal to the atmospheric pressure.
That is when a liquid boils, vapour bubbles are formed within the body of the liquid.
If the external pressure is lower than atmospheric eg at the top of a mountain, the boiling point temperature is reduced.
The graph below is of saturated vapour pressure in millibars (1000 millibars = 1 atmosphere) against temperature.

If the atmospheric pressure fell to $200\,\rm mb (= 1/5$ atmosphere) then the boiling point of water would be just above $80\,^\circ \rm C$.
Suppose that you put a beaker of water at $20\,^\circ \rm C$ in a box with dry air at atmospheric pressure in it then the water would evaporate (liquid to vapour at the liquid surface) and the vapour pressure would increase.
However at the same time some of the water vapour molecules would hit the water surface and become part of the liquid phase.
Initially the rate at which water molecules escape from the liquid phase is greater than the rate of return back to the liquid phase.
However eventually there are sufficient molecules in the vapour phase such that the rate at which molecules escape from the liquid is equal to the rate at which they return to the liquid - a dynamic equilibrium is set up and the pressure exerted by the vapour above the liquid is called the saturated vapour pressure.
If the temperature of the water is increased the rate at which molecules escape the liquid phase increases until the vapour pressure is such that the dynamic equilibrium (molecules leaving = molecules returning) is again set up resulting in the saturated vapour pressure being higher than before.
