In describing vapor pressure, I have often heard the following intuition: Assume a closed container that is partially filled with water. Above the water lies a vacuum. The water will vaporize until the pressure exerted on the surface of the water no longer allows the water to leave the solution. Is this even correct?

From Wikipedia, vapor pressure is defined as "the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases (solid or liquid) at a given temperature in a closed system." From this equilibrium perspective, the vapor pressure is reached when the rate of vaporization is equal to the rate of condensation.

From the first definition, is the building pressure really forcing the water molecules down to an appreciable extent? Or is vapor pressure simply a convenient way of describing such an equilibrium.

  • $\begingroup$ The first paragraph is not a definition, just an intuitive picture of what we mean by vapor pressure. Furthermore, it is not described by the thermodynamics, because the system is initially out of equilibrium (liquid + vacuum). But it helps because we intuitively understand what will be the dynamics and in the end the final (equilibrium) state (which is describe by the thermodynamics). $\endgroup$ – Adam Oct 5 '13 at 22:11
  • $\begingroup$ Both descriptions are very similar, the misconception lies the idea of "pressing the water molecules" down. The equilibrium is between the condensing and evaporating molecules. The water does not need be pressed down. $\endgroup$ – Alexander Oct 5 '13 at 22:30

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