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Homework Help: Vapor Pressure vs Atmospheric Pressure

  1. Sep 27, 2014 #1
    1. The problem statement, all variables and given/known data
    If the boiling point is the point at which vapor pressure > atmospheric pressure, so all of the water molecules can break free and fly into the atmosphere (i.e. overcoming the atmospheric pressure), then why is it that when atmospheric pressure > vapor pressure, the atmospheric pressure is pushing down on the water such that no vapor can escape too?

    I guess the question im asking is: Is vapor pressure independent of external pressure?

    2. Relevant equations

    3. The attempt at a solution
    So Vapor pressure is the result of the phenomenon when the surface molecules of a liquid gain enough kinetic energy to leave the liquid and become a gas. The force per unit area of these molecules is the vapor pressure.

    So it makes sense as T goes up, Vapor Pressure should increase since on average more molecules have the energy to break free from their intermolecular bonds.

    From an intermolecular bond perspective this makes sense. But form a pressure perspective im still a bit confused. If the boiling point is the point at which the vapor pressure > the atmospheric pressure so all the molecules can keep leaving the surface, then how is it that molecules can leave the surface via evaporation when pAtmosphere > pVapor Pressure?

    wouldnt the atmospheric pressure push all of the water molecules down into liquid? Just as all of the water molecules push against the atmospheric gas molecules to free themselves when pvapor pressure > patmosphere?


    Also if im comparing the vapor pressure of a solid vs that of a liquid (sublimation vs evaporation), then if I decrease the temperature of both solid and liquid at the same rate BELOW the melting point of the solid (but at different pressures), apparently the vapor pressure of the liquid will decrease faster than that of the solid, to the extent where Vapor Pressure of the Solid > Vapor PRessure of the Liquid.

    What's going on here? I assume its a different case because we're in two differnet phases?
    Last edited: Sep 27, 2014
  2. jcsd
  3. Sep 27, 2014 #2
    Being strict with the term, the boiling point (T) would verify Atmospheric Pressure = Vapor Pressure. So, at sea level (P = 1 atm), the vapor pressure of water has the value 1 atm (760 mm Hg, or whichever system should we use). And this is known for every component, where P-T graphs are plotted, showing that, for instance, pentane vapor pressure curve reaches those 760 mm Hg at 36.1 ºC (more volatility).
    About your question on how is it that molecules can leave the surface via evaporation when Atmospheric Pressure > Vapor Pressure, I supposse it has to do more with statistical mechanics for gases, and it's probably not that simple as the molecules of atmosphere press down the molecules of the liquid (gas), and they wouldn't come up. Since solids are a more stable phase, models would show that the probability that a molecule of the solid comes out of it in a point below its fusion point (sublimation) is low, but non zero. In the end, we would be talking about Quantum Mechanics and its strange world...
    To summarize: everything's quiet only at absolute zero.
  4. Sep 27, 2014 #3

    This is not quite correct. At the atmospheric boiling point, the vapor pressure is high enough to allow bubbles to form within the liquid by physically pushing the atmosphere back. However, if the vapor pressure is less than the atmospheric pressure, this does not prevent liquid from evaporating at the upper surface of the liquid, and then having molecules diffuse away into the overlying air. At the interface between the liquid and the air, the partial pressure of the liquid will be equal to the equilibrium vapor pressure of the liquid. The partial pressure of the vapor at the interface will be higher than in the bulk of the air, so the vapor can diffuse away into the air. This is what happens during ordinary evaporation.

    No. The vapor molecules diffuse into the atmosphere.

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