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Vapor Pressure and non ideal gases

  1. Mar 21, 2014 #1

    I understand that vapor pressure is independent of initial pressure, and depends only on temperature.

    However, is this true of a non ideal gas at high pressures?

    (I am specifically interested in the vapor pressure of a meg/water mixture at approx. 100 bar),

  2. jcsd
  3. Mar 21, 2014 #2
    What do you mean by initial pressure? Vapor pressure is a function of temperature and the concept is applicable to non-ideal gasses as well.
  4. Mar 21, 2014 #3
    Hi Jamese. Welcome to Physics Forums!!!!!

    What is meg, methylene glycol?

    If you have a two component system, then the saturated vapor pressure of the gas phase is a function both of temperature and concentration of the liquid phase, even in the ideal gas region and even if the liquid is an ideal solution. Is this what you are dealing with here?

    Last edited: Mar 21, 2014
  5. Mar 21, 2014 #4
    Thanks for the replies.

    I will take a step further back...because I think I might understand even less than I originally thought!

    I have a cylinder full of water and air at a particular temperature, if the air is saturated then the air has a particular water vapor pressure.

    If I then compress the cylinder to a higher pressure whilst the temperature remains constant, is the vapor pressure of the water vapor still the same?

  6. Mar 21, 2014 #5
  7. Mar 21, 2014 #6
    OK thanks.

    Is the vapor pressure directly related to a grammes/actual volume in the cylinder?
    And is the value for grammes/actual volume the same for both cylinder situations?
  8. Mar 21, 2014 #7
    If you mean the partial pressure of the water vapor, then no, as long as there is any liquid water present in the cylinder.
  9. Mar 23, 2014 #8
    I found a graph for air water mixtures.

    At the lower end of the pressure curves the grams of vapor per actual volume remains constant, regardless of the pressure.

    However, as pressure increase the lines start to curve, showing that this relationship is no longer constant and there is less grams of water vapor per actual volume,

    What causes this?


    Attached Files:

    Last edited: Mar 23, 2014
  10. Mar 23, 2014 #9
    See if you can calculate the points on the graph using the ideal gas law and the vapor pressure of water at the different temperatures. This should be easy to do. I haven't gone over the graph in detail, but I do know that at pressures above ~10 atm for this system, you start to get deviations from the ideal gas law.

  11. Mar 23, 2014 #10
    So for pressures less than 10 atm whilst this system behaves as an ideal gas, then the amount of water vapor that air can hold is constant (on an actual volume basis) and varies only as function of its temperature.

    In the less than 10 atm situation is the partial water vapor pressure the same at all system pressures?

  12. Mar 23, 2014 #11
    The graph you showed plots water mass per air mass (called the mixing ratio) . That's different than water mass per air volume (called absolute humidity). Also note that you shouldn't say "grams of vapor per actual volume". grams is a unit, not a physical quantity. The physical quantity is called mass.
  13. Mar 23, 2014 #12
    OK thanks noted about the units.

    I converted the mass per mass graph, to give the mass of vapor per actual volume.

    This then showed that the mass of vapor per actual volume remains constant at varying system pressures less than 10 atm.
    This is because the system is behaving as an ideal gas at pressures up to approx. 10 atm and at pressures above this real gas behavior comes into play (as Chester pointed out to me).

    Forgive me for the probably repeating myself, but the baby steps approach really help with my understanding.

    Next question; In the less than 10 atm situation is the partial water vapor pressure the same at all system pressures?

  14. Mar 23, 2014 #13
    Yes. Roughly. But, beyond the ideal gas region, the concept of partial pressure loses some of its meaning. That's because, at higher total pressures, the mole fraction of water in the gas phase is no longer equal to the vapor pressure of pure water at the system temperature divided by the total pressure (except, of course, for single component water).
  15. Mar 24, 2014 #14
    OK thanks.

    In the low pressure ideal gas situation; does the partial pressure of the water vapor depend on the particular gas in the system?

  16. Mar 24, 2014 #15
    Well, air is essentially non-condensible. But, if you have a mixture of water and a soluble species
    like ethanol, then there can be a substantial amount of ethanol in the liquid phase, and the partial pressure of water vapor will be less than its pure liquid vapor pressure.

  17. Mar 24, 2014 #16
    OK thanks.

    Am I correct in believing what you describe with respect a soluble species is described by Raoult's law?
  18. Mar 24, 2014 #17
    Yes, for an ideal solution.

  19. Mar 25, 2014 #18
    In two mails up you mentioned the solubility of something in the liquid phase affecting the partial pressure.

    Will the solubility of a gas, for example carbon dioxide (in gaseous form) in water also effect the partial pressure of the water vapor?

  20. Mar 25, 2014 #19
    A little, but not much. Even air dissolving in water affects its vapor pressure slightly, but the solubility of air in water is very low (and thus air is usually considered a "non-condensible").

  21. Mar 25, 2014 #20
    At higher pressures (100 bar plus) do we have to be more concerned with dissolved gasses, or is the effect still negligible?

    If I have the vapor pressure of a vapor in a gas at atmospheric system pressure, or low system pressures where vapor pressure is a constant.
    How do I then find the vapor pressures at higher system pressures, where the real gas effects come into effect?

    Thanks for the continued assistance, I am finding this very informative!
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