# Can water be compressed so much that it freezes?

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)?

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.

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.

Mapes
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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?

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.

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.

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.

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.

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.

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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.

DaveC426913
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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.
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.

So, we couldn't have oceans deeper than 100 km on Earth because the water at the bottom would freeze under pressure.

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.

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.

DaveC426913
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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.
I did not get that impression at all.

Mapes
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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.

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.

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.

I agree with this part.

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.

Mapes
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If we look at PV=nRT, as pressure increases so does temperature.

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

So why for water would the temperature decrease upon the addition of pressure.

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.

dacruick,

If we look at PV=nRT, as pressure increases so does temperature.

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

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.

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.

Mapes
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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

Mapes
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It is an idealization made because liquids* are far, far easier to deform by shearing than by compressing. All real materials are compressible to some extent.

*(Note: fluids includes gases, which are relatively easy to compress.)

im pretty sure if you look at the triple point chart that chmdude has up it tells all you need to see. increasing the pressure, you cannot make ice. But if you look in the bottom left, apparantly you can increase the pressure of steam to make ice.

but here is carbon dioxide and you can increase the pressure of this to phase change it from liquid to solid

ahh i see, so the chart above just didn't go to a high enough value along the y-axis.?

Mapes
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ahh i see, so the chart above just didn't go to a high enough value along the y-axis.?

It's a weird case, where pressure could cause ice (normal ice, in a hexagonal crystal structure) to melt into water, then resolidify at enormous pressure in a different crystal structure, all at a temperature of 0°C. Pretty amazing.

It's a weird case, where pressure could cause ice (normal ice, in a hexagonal crystal structure) to melt into water, then resolidify at enormous pressure in a different crystal structure, all at a temperature of 0°C. Pretty amazing.

Exactly. What we call "frozen water" isn't the only solid form of water, it's just the crystal structure that happens to form at atmospheric pressure. If you compress anything below its supercritical temperature it will eventually become solid (or a glass).

However, are you sure your phase diagram is accurate? IIRC once you pass the supercritical points of temperature and pressure, that is all that exists. I could be wrong and/or my P-chem professor left something out.

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Is it just me or am I missing an important consideration?
Pressure applied with a reasonable amount of thermal release(radiative, convective, conductive) will permit and demand a reduction in the pressure sample temperature; to my understanding that is.

In other words, again to my understanding, is that if you placed water in a "piston" compression arrangement that has no ability for thermal expenditure, freezing is impossible.
An extreme example, of course, and not possible per se, but hopefully illustrative.

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Mapes
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Is it just me or am I missing an important consideration?
Pressure applied with a reasonable amount of thermal release(radiative, convective, conductive) will permit and demand a reduction in the pressure sample temperature; to my understanding that is.

I'm not sure what you mean by "pressure sample temperature." (The temperature of the pressurized sample?). What is "thermal release"?

Pressurizing a material will generally increase its temperature; to simplify the discussion, we've been referring to processes occurring at constant temperature. It would be interesting, granted, to consider the temperature effects of adiabatic pressurization and solidification.

I'm not sure what you mean by "pressure sample temperature." (The temperature of the pressurized sample?). What is "thermal release"?

Pressurizing a material will generally increase its temperature; to simplify the discussion, we've been referring to processes occurring at constant temperature. It would be interesting, granted, to consider the temperature effects of adiabatic pressurization and solidification.

Yes, the temperature of the pressurized sample.
Thermal release: Perhaps a bad use of terms on my part, but if the temperature(thermal energy) of the pressurized sample is unable to "migrate" beyond the physical domain of that sample, the temperature of that sample can not be reduced.

The "thermal release" then, in my context, is referring to the the external environmental conditions which permit or disallow dissipated heat in all it's forms. Such as a heat sink, if you will.
Without heat dissipation, freezing is impossible.

That's my contention. Not saying I'm right, but those are my humble thoughts.

Exactly. What we call "frozen water" isn't the only solid form of water, it's just the crystal structure that happens to form at atmospheric pressure. If you compress anything below its supercritical temperature it will eventually become solid (or a glass).

However, are you sure your phase diagram is accurate? IIRC once you pass the supercritical points of temperature and pressure, that is all that exists. I could be wrong and/or my P-chem professor left something out.

The supercritical point is just a point where the line of first order transitions between liquid and vapor ends. The critical point itself is point where there is a second order transition. There the heat capacity diverges with a certain universal critical exponent.

But this has nothing to do with the other phase transitions in the phase diagram.

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.

QUESTION: Does this hold true for all elements?