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Vapor pressure and atmospheric pressure

  1. Jun 28, 2013 #1
    the book says vapor pressure is independent of atmospheric pressure, if that is the case then why must the vapor pressure of a liquid be equal to the atmopheric pressure for it boils. I pictured atmospheric pressure as a column of air sitting above a liquid and therefore pushing down the gas molecules from the liquid ,thus affecting its vapor pressure
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  3. Jun 29, 2013 #2


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    When you consider a column of air sitting above a piston below which there is the water, then the vapour can only form if its pressure is equal or higher than the pressure of the air. If it is smaller, the vapour molecules will be pressed back into the liquid. This does not mean that the vacuum pressure changes appreciably with external pressure. It is rather a statement about mechanical equilibrium.
  4. Jun 29, 2013 #3


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    Vapor pressure, not 'vacuum pressure', whatever that is.
  5. Jun 29, 2013 #4
    The last sentence is where I'm getting confused. There was a problem in my book that said that you have a pot of boiling water at sea level and boil it. Then you take that same pot of water to mount everest and boil it, and the boiling point it lower at higher altitude. Due to higher altitude the atmospheric pressure is lower. I saw this as the atmospheric pressure exerting less force on vapor pressure so you would need less energy/temperature to get the water to boiling point, where atmospheric pressure = vapor pressure. So my reasoning is, less atmospheric pressure = less vapor pressure = lower boiling point. But there's a problem in my book that says that vapor pressure only depends on temperature and particles dissolved in solution. I understand that temperature and vant hoff factor affect vapor pressure but why doesn't atmospheric pressure affect it too?
  6. Jun 29, 2013 #5
    vapor pressure of a liquid is the pressure at which vapors of the liquid are saturated. if you increase that pressure by just one tad bit by introducing additional vapors into the fixed volume, the molecules will be so close together that you will have a liquid instead of a gas.

    so , :) a liquid will boil when the pressure exerted by its vapors (vapor pressure) equals the atmospheric pressure. this simply means that when you have an equal amount of molecules of that liquid escape and exist in the gas phase as there are molecules of air, we observe boiling.

    so, at sea level, you have more molecules of air in a fixed volume (higher pressure) then you do high up on a tall mountain. this means that high up on the mountain the liquid we are trying to boil will need to have a fewer number of its molecules escape in the gas phase than it did at sea level in order to equal the number of molecules of air in a fixed volume. hence you will observe boiling at a lower temperature.

    so as you can see the atmospheric pressure does not affect the vapor pressure (see definition at top of my post) but it affects the temperature of the boiling point
  7. Jun 29, 2013 #6
    Got it, thanks so much!
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