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Can water be compressed so much that it freezes?

  1. Apr 25, 2009 #1
    Intuitively, it makes sense to me that it would. But then again, weird stuff can happen under extreme circumstances...

    What would actually happen? Would the temperature of the substance actually increase or decrease under ENORMOUS pressure (i.e., that the molecules could barely move, if possible)?
     
  2. jcsd
  3. Apr 25, 2009 #2
  4. Apr 25, 2009 #3
    Yeah, the temperature part of it confused me. I recognize temperature is directly proportional to pressure, but how come the pressure can't be so intense that the molecules are squeezed so tightly that they almost have no kinetic energy?

    I understand the equations, but conceptually, I don't.
     
  5. Apr 25, 2009 #4
    For a substance to be a liquid, its atoms or molecules have to be able to move around each other with ease. The more you compress atoms, the less room they have to move.
     
  6. Apr 25, 2009 #5

    Mapes

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    Temperature is only directly proportional to pressure for an ideal gas at constant volume. Depending on the material, a pressure increase can cause an increase or decrease in temperature. Water is one of the materials that gets hotter when pressure is applied, but the increase is much less than directly proportional.
     
  7. Apr 25, 2009 #6

    DaveC426913

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    Freezing the water requires that the molecules have a chance to crystalize. Crystalizing requires that the molecules have enough elbow room to do so.

    It's not that water wil NOT freeze under extrreme pressure, it's just that water is very complex.

    There are actually almost a dozen forms of ice water. Ice I (naturally-formed water-ice) is only one.
     
  8. Apr 25, 2009 #7

    Mapes

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    A good rule of thumb is that if any solid phase has a smaller molar volume than the liquid phase(s) (and I don't know of a material where this isn't the case), then the material will freeze under pressure, regardless of temperature. It's the energetically favorable thing to do.
     
  9. Apr 26, 2009 #8
    To add to what DaveC says, according to this chart, at 300K, or 27C, water is a type 6 solid at 10^9 Pascal pressure.

    phase.gif

    As you might recall, water expands on freezing. But pressure acts to reduce volume. This is the opposite of what you would want to get it to crystalize in the normal manner. Under higher pressure, at room temperature, water is a different form of solid--whatever type 6, 8 9 and 10 are.
     
    Last edited: Apr 26, 2009
  10. Apr 26, 2009 #9
    Squeezing something does not necessarily make it change temperature. The solidification of a liquid upon pressure has nothing to do with changes in temperature. It is purely because of the applied pressure.
     
  11. Apr 26, 2009 #10

    DaveC426913

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    Squeezing something does actually change its temp, yes. Look at adiabatic heating, or look at pounding a nail into wood.

    But that's not what solidifies a liquid, no.
     
  12. Apr 26, 2009 #11
    So, we couldn't have oceans deeper than 100 km on Earth because the water at the bottom would freeze under pressure.
     
  13. Apr 26, 2009 #12
    For sure, but the tone of the discussion up until this point had been that some how squeezing it made it colder so it then freezes.

    Just to be clear: under pressure, the equilibrium state can change, thus even at the same temperature it is possible to cross phase transitions by applying pressure.
     
  14. Apr 26, 2009 #13

    DaveC426913

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    I did not get that impression at all.
     
  15. Apr 26, 2009 #14

    Mapes

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    No, I don't agree with this; the Sandia article, for example, mentions that the water heats up to >100°C. Any material that expands with temperature will get hotter upon being pressurized.

    I agree with this part.
     
  16. Jan 12, 2010 #15
    If we look at PV=nRT, as pressure increases so does temperature. So why for water would the temperature decrease upon the addition of pressure.
     
  17. Jan 12, 2010 #16

    Mapes

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    PV=nRT is a relationship for ideal gases. We are talking about liquid and solid water.

    Nobody is saying that it does. We are discussing how solid water can be the stable phase at >0°C when great pressure is applied.
     
  18. Jan 12, 2010 #17
    dacruick,

    This thread is discussing liquid water not water vapor.

    For water vapor:

    If you compress water vapor to a pressure at or beyond 22.09 MPa the water vapor is considered to be in the supercritical pressure range and the Ideal gas law cannot be used. When you hear of a powerplant operating in the supercritical range that means the boiler is operating at a pressure greater than 22.09 MPa.

    Also, for real gases the Ideal gas law should not be used (or be used with caution). You should use a modified version of the Ideal gas law that incorporates the compressibility factor, Z. You could also use a more elaborate equation of state such as, Benedict-Webb-Rubin or Beatie-Bridgeman.

    Thanks
    Matt
     
  19. Jan 12, 2010 #18
    PV=nRT is the ideal gas equation and tells us nothing about phase change. Also, you can have an isothermal process (expansion/compression of a near ideal gas, for instance).

    This is an easier phase diagram to understand
    http://www.jamstec.go.jp/xbr/2deepstar/02deepenv/figures/page24_blog_entry6_1.png

    So the answer is, it depends on what you call freeze and what temperature you're at.
     
  20. Jan 13, 2010 #19

    Mapes

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    Last edited by a moderator: Apr 24, 2017
  21. Jan 13, 2010 #20
    As far as I remember, I was thought once that one caracteristic of fluids is that they are incompressible.

    So is that only an idealisation? and they actually can be, given enough pressure? A bit confused here... But maybe its not the same thing and it just means the volume of a fluid cannot ever change?

    cheers
    Frederic
     
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