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I Difference between Partial pressure and vapor pressure

  1. May 23, 2016 #1
    Partial pressure must be less than or equal to the vapor pressure if there is no liquid present. However, when both vapor and liquid are present and the system is in the phase equilibrium, the partial pressure of the vapor must be equal to the vapor pressure and system is said to be saturated.

    I am totally confused with above paragraph. Please help. I understand vapor pressure's definition and also understands it in the context of Cavitation. when local pressure in the flow becomes below the vapor pressure cavitation occurs.

    Please help
     
  2. jcsd
  3. May 23, 2016 #2

    DrClaude

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    It's hard to tell what it is you don't understand. The vapor pressure is the partial pressure at equilibrium of a substance in the gas phase when the conditions (T,P) are such that the substance should normally be in the liquid phase (as per the phase diagram). Therefore, the second sentence
    is true by definition. If the system is not saturated, then the system is out of equilibrium. Note that out of equilibrium, the system can either be undersaturated or oversaturated.

    This has to be true at equilibrium, because if the partial pressure is greater than the vapor pressure, the system is metastable and will relax to liquid + vapor phases, until the equilibrium (partial pressure = vapor pressure) is reached.
     
  4. May 23, 2016 #3

    BvU

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    Partial pressure is pressure if there is only one component present in the vapour. When there are more components (think of e.g. air above water) the paragraph makes sense, I hope ?
     
  5. May 23, 2016 #4
    In an ideal gas mixture (say water and air), each species behaves as if the other species is not even present. I'm sure you are aware that if you have liquid water and pure water vapor present in equilibrium, the water vapor is at the equilibrium vapor pressure. In the case of a gas mixture of water vapor and air in contact with liquid water, the partial pressure of the water vapor in the gas phase will be equal to the equilibrium vapor pressure. However, if you have a gas mixture of water vapor and air, and the partial pressure of the water vapor in the gas phase is less than the equilibrium vapor pressure, you can't have any liquid water present.
     
  6. May 23, 2016 #5
    I got your point sir. Thank you.

    Why we consider only partial vapor pressure instead of actual atmospheric pressure comparing with vapor pressure?
    Because in case of boiling we have read that when vapor pressure becomes equal to the atmospheric pressure; boiling occurs.
     
  7. May 23, 2016 #6
    That's because, if you look at the bubbles under the water, they contain pure water vapor without any air. So there is enough pressure within the bubbles for them to grow (i.e., allow the volume below the liquid surface to expand by pushing back the atmosphere above). Also, when boiling is occurring, the system is not in thermodynamic equilibrium. If it were, and the pressure were 1 atm, there would have to be no air present. Or, if the total pressure were higher than 1 atmosphere, the system could be in equilibrium with air present.
     
  8. May 23, 2016 #7
    Ok I am going through some previous discussions to understand your reply properly. Thank you for your support.
     
  9. May 23, 2016 #8
    This is one of the reply given by you

    "The water vapor in the rising bubbles is virtually at equilibrium with the liquid water that surrounds it. This is essentially the same situation that exists in a closed system containing only water. If the liquid water temperature is 100 C, the water vapor pressure inside the bubbles is 1 atm.

    The air above the liquid in the open container is not at thermodynamic equilibrium with the mixture of liquid water and vapor bubbles. The only thing that the air does is control the overall pressure acting on the mixture (typically, 1 atm.). In order for the water to boil (i.e., for bubbles to form below the surface), the temperature has to be high enough for the vapor pressure inside the bubbles to push back the atmosphere so that the bubbles can grow. This happens when the equilibrium vapor pressure matches the atmospheric pressure above the mixture.

    Chet"


    Now my point is atmospheric pressure is playing a significant role is boiling. We say in mountains water boils below 100 degree Celsius because of low atm pressure. So why we only mention partial vapor pressure (it is only 3% of atmospheric pressure).
     
  10. May 23, 2016 #9
    As I said, it is not a thermodynamic equilibrium situation. Read what I said carefully: inside the bubbles, you have 100% water vapor, not 3%.
     
  11. May 23, 2016 #10
    Ok, the word partial vapor pressure is used for; bubble's vapor (i.e. 100% vapor) not the vapor above the liquid water surface.

    Vapor above the liquid surface is in the thermodynamic equilibrium with liquid water i.e. called vapor pressure even in the open system.

    So when atmospheric pressure changes it is the vapor pressure that changes. So we require less partial vapor pressure inside bubble for boiling.
     
  12. May 23, 2016 #11

    BvU

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    This isn't going very good at all:
    No. The vapour phase above the water has a certain pressure that is the sum of the partial pressures of the various components (O2, N2, CO2, etc. and some water). Depending on the humidity, the partial water vapour pressure is normally only a small fraction of the total pressure.
    In the bubbles under the surface of boiling water the partial pressure is the pressure is the vapour pressure -- there is only one single component present.
    No. As Chet says: this is generally not an equilibrium situation. As long as the vapour pressure that is associated with the temperature of the water is higher than the partial pressure of the water in the vapour phase, evaporation will continue to happen. That's why the laundy dries on a washing line, streets dry up after rain, a cup of coffee goes empty after a few days, etc.
    No. The equilibrium vapour pressure depends on the temperature only. Boiling occurs when the temperature is so high that the equilibrium vapor pressure of the liquid is higher than the total pressure above the liquid.
     
  13. May 23, 2016 #12

    russ_watters

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    You have invented a new term: "partial vapor pressure". You should not do that if you want to understand the two terms you were asking about. That term has no meaning and inventing it and trying to use it will just confuse you.
     
  14. May 23, 2016 #13
    Sir, partial pressure of vapor above liquid surface; I was calling it partial vapor pressure. That's all.
     
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