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Phase diagram of water in real life

  1. Nov 17, 2017 #1
    Hello,

    All substances that are either in the liquid or solid state have a small vapor pressure which implies that they are always slowly (solids especially) evaporating and turning into a gas (vapor).

    The phase diagrams of water and carbon dioxide indicate only states (P , V) of pressure and temperature that are in thermodynamical equilibrium but in real life, both water and dry ice are not in equilibrium: water is present both as a liquid (in oceans and lakes) and as a vapor at the same conditions of pressure P and temperature T.
    That said, how is the phase diagram of water useful since water is not in stable equilibrium? Is the phase diagram useful only for water in a laboratory setting?

    Thanks!
    fog37
     
  2. jcsd
  3. Nov 17, 2017 #2
    If the system is closed, such that the water is in a sealed container, and is held at constant temperature and pressure, then the system is in equilibrium. Specifically, this could be considered an Isothermal-isobaric ensemble since N, P, and T are fixed constants.
     
  4. Nov 17, 2017 #3

    rbelli1

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    This answers your question

    When you "move" from one region to another in the phase diagram you have a change in energy but not P or T. It is at the exact same point in the diagram thus the quotes around move.

    BoB
     
  5. Nov 18, 2017 #4
    Thank you.

    rbelli1, I am grasping how moving from one region of the diagram to another does not change the P and T. Each point in the diagram has a different P and V.
    And as NFuller mentioned, the equilibrium is achievable within a closed container but water on earth is not inside a closed container and that is why I am wondering when the water phase diagram is useful except of the closed container situation...
     
  6. Nov 18, 2017 #5

    Borek

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    Staff: Mentor

    It definitely tells you when to expect water to freeze (think ice on puddles) and when to expect water to condense and whether it will condense into ice or water (think clouds), so it seems to be pretty useful outside of the lab settings.
     
  7. Nov 18, 2017 #6
    Thanks Borek.
    • Just to make sure, in any phase diagram, regardless of the single substance state (solid, liquid, gas), the temperature T is the temperature of the substance itself (measured by inserting a thermometer in it) and the pressure P is the total external pressure acting on the substance, correct? A solid, a liquid and a gas also have an internal pressure pushing towards the outside. How does that factor in? Is the internal pressure always equal to the total external pressure since we talking about stable equilibrium states?
    • As far water goes, is the provided partial pressure of water in the air (mixture of gases) always equal to the saturated vapor pressure of water? I don't think so. If that was the case, RH=100% and evaporation would never be possible. At The higher the temperature the higher the water partial pressure will be and the higher the reachable saturated vapor pressure will be. The pressure in the table below should just be the partial vapor pressure (not the saturated vapor pressure).
    T, °C P, torr P, atm
    0 32 4.5851 0.0060
    5 41 6.5450 0.0086
    10 50 9.2115 0.0121
    15 59 12.79 0.0168
    20 68 17.54 0.0231
    25 77 23.76 0.0313
    30 86 31.84 0.0419
    35 95 42.20 0.0555
    40 10 55.3 0.0728
    45 11 71.92 0.0946
    50 12 92.5 0.1218​

    When dynamic equilibrium is reached, the water vapor pressure is the maximum pressure the water vapor can have at that specific temperature, correct? The relative humidity is 100% when the water partial pressure is saturated, i.e. maximum. When further cooled, the airborne water vapor will condense.
    • In the discussion of dew point, condensation, saturation the presence of the other gas constituents in the air seem to not have an effect, correct? Well, but when we say that water freezes at p=1 atm and T=0 C, we are including the partial pressures of all other gases in determining the freezing behavior of water, correct?
    • From the pure water phase diagram, we see that at a temperature of 20 °C and absolute pressure of 1 atm (NTP), the water vapor pressure is just 0.00231 atm. When we study the phase diagram of water, we would use 1 atm (not 0.00231 atm), and obtain that water is supposed to be a liquid (not a vapor). However, in real life, vapor seems to be the direction/phase water is going for (evaporation would evaporate all water if it wasn't for the water cycle)...
     
  8. Nov 18, 2017 #7
    In cases where the gas phase consists of a mixture of air and water vapor, the equilibrium is established when the partial pressure of water vapor is equal to the equilibrium vapor pressure of water at the liquid temperature. This applies even if the total pressure (due mostly to the air) is 1 atm.
     
  9. Nov 19, 2017 #8
    Thank you Chestermiller.

    So the other air molecules (nitrogen, oxygen, etc.) present in the air can, at most, slightly delay when dynamic equilibrium (condensation=vaporization for water) is reached due to collisions with the water vapor molecules. The partial water vapor pressure is only affected by T (same goes for all other gases in the air).

    But, just to make sure, at a certain specific temperature T, the water vapor pressure would continue to increase until it reaches its saturated vapor pressure value (the highest pressure it can be for that temperature) once evaporation rate = condensation rate.
    But if stating the partial pressure at a certain T means stating the saturated vapor pressure, that implies that equilibrium is always reached at a certain temperature....Is that the case?
     
  10. Nov 19, 2017 #9

    Borek

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    Not sure what your question is - given enough time every system goes to equilibrium.
     
  11. Nov 19, 2017 #10
    Sorry for the lack of clarity. Let me reformulate my question better:

    The water vapor pressure (i.e. partial vapor pressure) at the temperature T=20 C is 0.02 atm (see table above). Is that tabulated vapor pressure referring to the saturated vapor pressure of water vapor?

    From the phase diagram of water, the boiling temperature is 100 C when P=1atm. That pressure of 1 atm is the total air pressure (due to all gases in the air) above the liquid interface and not just the partial vapour pressure which is only 0.86 atm at that temperature. Is that correct?
     
  12. Nov 19, 2017 #11
    Yes.

    No. The vapor pressure of water at 100 C is 1 atm. When water boils at 100 C (in contact with the atmosphere), the bubbles that form within the bulk of the liquid are pure water vapor with virtually no air inside them. These bubbles are able to form because their pressure is high enough to allow the surrounding liquid water to push back the air atmosphere above. If the room is not hermetically sealed (and if you have enough liquid water), if you keep adding heat to the liquid, the air will be purged from the room, and you will eventually have only water vapor.
     
  13. Nov 19, 2017 #12
    Thank you!

    But if the pressure of 1 atm (boiling at T=100 C and P=1 atm) is solely and entirely due to water vapor, it would mean that the actual total external pressure must be higher than 1 atm since there are other gases and vapors that also contribute to the overall pressure with their partial pressures. It was said before that the variable pressure P in the phase diagram is the total external pressure and not just the water vapor pressure...
     
  14. Nov 19, 2017 #13
    I never said that, and, it's not correct. The pressure P on the phase diagram is the equilibrium vapor pressure of water. In cases where other gases are present (e.g., air), we can still use the phase diagram if we regard equilibrium for water to be established if the partial pressure of the water vapor in the gas mixture above the liquid is equal to the equilibrium vapor pressure. But, at 100 C, if the total pressure is 1 atm and equilibrium is present, the gas space must contain only water vapor. Therefore, if it is possible for gases to escape the room, all the air will eventually be purged from the room (provided there is still liquid water present).
     
  15. Nov 19, 2017 #14

    Borek

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    Imagine a cylinder filled with water, with a movable piston resting on the water surface. External pressure is that of the atmosphere - 1 atm. You start to heat water. Once you get to 100 °C water starts to boil and the vapor pushes the cylinder up. But, as the cylinder moves pressure of water vapor never exceeds 1 atm.

    In reality there is no well defined boundary between water vapor and the air, but the principle remains the same.
     
  16. Nov 19, 2017 #15
    Ok thanks.

    That makes sense. So The pressure P on the vertical axis is the saturated vapor pressure of water and at T=110 oC the saturated water vapor pressure must be 1 atm.

    However, in real life, when we see a pot of water boiling, there is air above the water (which contains a variable amount of water vapor, on average around 1% at sea level, and 0.4% over the entire atmosphere) so the contribution from water vapor is usually very small and does not contribute with 1 atm. The measured outside pressure of 1 atm derives from all the partial pressures of the gases/vapors in the air together. that means the water is boiling at T=110 oC even when the water vapor pressure is much much less than 1 atm.
     
  17. Nov 19, 2017 #16

    Borek

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    No, it must be higher. 1 atm is at 100 °C.

    Note:
    • by definition boiling means water starts to evaporate not only at the surface, but also under the surface
    • despite being defined as force/surface pressure is a property of a point in space
     
  18. Nov 19, 2017 #17
    That is because the water vapor in the air is not at equilibrium with the liquid water. Only at the surface between the liquid water and gas phase is the water partial pressure equal to the equilibrium vapor pressure (and, at this location the partial pressure of the oxygen and nitrogen in the air is zero). The water vapor generated by the evaporation diffuses from the surface to the bulk air, driven by the partial pressure difference between the surface and the bulk air. Gradually, by this process, the air in the room gets "humidified."
     
  19. Nov 20, 2017 #18
    Thanks again.

    1) "That is because the water vapor in the air is not at equilibrium with the liquid water". I agree.

    2) "Only at the surface between the liquid water and gas phase is the water partial pressure equal to the equilibrium vapor pressure (and, at this location the partial pressure of the oxygen and nitrogen in the air is zero). " Why are the partial pressure of oxygen and nitrogen zero at the interface and only the water vapor exerts a pressure? That would explain how the only vapor pressure present at the interface is the water vapor pressure and its value is exactly 1 atm. But what happened to those other gases? Did they get pushed out of that region of space near the air-liquid interface?

    3) The water vapor generated by the evaporation diffuses from the surface to the bulk air, driven by the partial pressure difference between the surface and the bulk air. Gradually, by this process, the air in the room gets "humidified." I agree.
     
  20. Nov 20, 2017 #19
    Yes. Even below the boiling point, the partial pressure of water vapor at the interface is equal to the equilibrium vapor pressure, so the partial pressures of the other gases total 1 atm minus the equilibrium vapor pressure of water.
     
  21. Nov 20, 2017 #20

    russ_watters

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    One point I'm not sure if you have clear about how boiling works:

    Boiling happens where heat is applied. In a pot of water on an electric range, that's at the bottom of the pot, where the pot touches the heating element. So the water that is actually boiling is the water touching the hot bottom of the pot. This water, clearly, only has water around it; no other gases. The pressure at the bottom of the pot is atmospheric pressure plus the weight of the water above, which makes the boiling temperature just slightly above 100C. The bubbles form, then rise to the surface (and cool to 100C).

    That's what differentiates boiling from evaporating: it happens at the bottom of the pot, not at the surface. So it doesn't matter what the composition of the air above is; it just matters how fast you add heat. Turn off the burner and the water will still be at 100C, but it will evaporate from the surface at a rate dependent on the composition of the air.
     
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