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Q & A: freezing point of heavy water

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Most recent answer: 05/06/2016
Q:
How the prezence of deuterium in the water influences the water's freezing temperature?
- Vic
usa
A:
It's easy to guess that deuterated ("heavy") water will have a higher freezing point than ordinary water. The reason is that the more massive bound deuterons have less zero-point energy than do the less massive bound protons, for reasons easily understood from basic quantum mechanics. That means their energy can drop more as they freeze into a more rigid pattern. Sure enough, Wikipedia says that heavy water has a freezing point of 3.8C.   My guess is that the freezing point change is very nearly linear in the percent deuterated.

Mike W.

(published on 04/02/2011)

Follow-Up #1: enriching deuterium by freezing

Q:
Does this mean DOH will also freeze quicker than HOH? Can this be used to separate deuterated water out of ordinary water?
- Jean (age 14)
A:
That's a really interesting thought. I'll give a rough answer, based on the approximation that each D or H in the ice or water has the same energy (enthalpy, to be precise) regardless of whether the nearby sites have D's or H's. From the 3.8C freezing T of D2O together with the latent heat of freezing of H2O (and the assumption that D2O has roughly the same thermodynamic parameters) you can calculate how much the free-energy  of moving a single D to ice from liquid is below that of the same move for H. I get about 7*10-23 J. Now the thermal energy scale (absolute T* Boltzmann's constant) at around 0C (or 3.8C, all the same at this level of approximation) is around 4*10-21 J. So the free-energy preference for freezing the D's is, in thermal energy units, under 2%. The relative enrichment of the D concentration in the ice in a partially frozen solution is under 2%.

Now you might wonder if, by holding the temperature at say 1C, you could get some ice to form that only contains D2O or DHO, since H2O doesn't freeze at all at that T. The loss of entropy in separating the D from the H would prevent any ice from forming in ordinary low-D water at that point. It takes a lot of energy to unmix things.

Mike W.

(published on 04/03/2011)

Follow-Up #2: How to raise the Freezing point of water?

Q:
Is there anything that would raise the freezing point of water so things would freeze above 32 degrees?
- Matthew (age 27)
Oakridge
A:
Matthew,

  For something to dissolve in water,  the water must prefer being in the solution to staying apart as pure water. In technical terms, the solute lowers the chemical potential of the liquid water.  Unless the solute lowered the chemical potential of the solid ice even more, adding some solute will always favor the liquid, lowering the freezing point. Since the solid ice is a regular pattern of molecules, small amounts of solute break up the pattern, and are excluded from the ice. A few examples are sea water (with salt), sugar water, and water with alcohol. In all three cases, adding the solute lowers the freezing point. Here is a link to a previous question that was asked on our site that gives a more thorough explanation.

The only exception to this argument would be some molecule that fit into the ice structure even better than ordinary water. Actually, water made with the deuterium isotope of hydrogen rather than ordinary hydrogen fits the bill. This "heavy water" actually freezes at 3.8C (39F) rather than 0C.

 Thanks for a very interesting question! -Zach (+mbw)



(published on 12/28/2009)

Follow-Up #3: raising the freezing point

Q:
Is there a way to either 1) raise the freezing temperature of water or 2) raise the melting temperature of ice? Basis: if salt decreases the freezing point, what can be added to increase the freezing point?
- Ted (age 40)
Denver, CO
A:

The true thermodynamic freezing and melting points are identical, so raising either one means raising the other.

Any solute that doesn't get incorporated into the ice lowers the freezing point, for basic thermodynamic reasons, as we discuss elsewhere. () So you have to ask what sort of molecule would become part of the ice without changing its properties enough to make us want to call it something besides ice. Heavy (deuterated) water is the obvious example, so I've put your question in this thread.

Mike W.


(published on 09/05/2013)

Follow-Up #4: freezing at 40F ?

Q:
Hi, I have been following this chain as I build a backyard skating rink every year and am frustrated by warmer than ideal winters. I thought that since you can lower the freezing point of water with additives like salt, you might be able to raise the freezing point of water with other additives. From the chain, it seems like this is not the case, except for heavy water, which does not appear to be an economical solution. Is there another liquid with ice like qualities that freezes at 40 degrees F or so and that is non-toxic, relatively inexpensive, etc.?
- Abe (age 37)
Massachusetts
A:

You're right that no additive will raise the freezing point of water, except for the unacceptably expensive deuterated water. (Deuterated water also is toxic to plants and animals.)

I'm sorry to say that I can't think of any good water substitute for your purposes. Perhaps some reader will think of one, but that seems unlikely to me. 

Here's a possible idea, but I'm not sure it's practical. When the temperature is around 32°F, your rink is too warm. Your house is too cold until you heat it. Maybe you could get a large heat pump to pump heat from your rink to your house. This general type of heat pump is a very efficient way of heating houses in moderately cold weather. I don't know how well one could be adapted to your pond and how much it might extend your skating season. Maybe you could interest some local geothermal HVAC installer in this project.

Something similar, run in reverse, should work well for swimming pools during the period when the pool isn't hot enough yet but the house is too hot. The pool application would almost certainly be doable, since it's easier to exchange heat with circulating liquid water than with ice.

Mike W.


(published on 11/27/2013)

Follow-Up #5: separating deuterium from hydrogen

Q:
Hello,I was reading your Heavy Water discussion (d2O = Deuterium (Oxide?] Water).Thank you in advance for taking the time to share your knowledge and thoughts on these topics.I was wondering - I need to remove Deuterium out of my drinking water ... would it be safe to say that I could do one of the following to do so easily and efficiently?Suggestion 1. - Freeze an ice cream container full of drinking water from my well. Throw it in the freezer and then pull out of freezer and put in a controlled container that lowers the temp to 4 degrees Celsius and let thaw and drip into another empty pail to collect all the light water (H2o - >3.2 degrees celcius). once the ice has melted while in a 4C the rest of the ice can be safely considered to be D2O yes? your thoughts?2. OR, if we add a ton of sea salt to the water to lower the freezing point to far more than D2O then we should be able to do the same process as the first but invertedreversed freezingmelting points of course.What do you think? Please keep in mind that money and water supplies are not a concern as I have enough of those to accommodate the needs to do so - within reason of course.Many thanks in advance!Best Regards,Derrick
- Derrick (age 44)
Edmonton Alberta Canada
A:

As we pointed out in the answers up-thread, the freezing/melting process only enriches the concentration slightly. It's not very easy to repeat often enough to get good separation of the deuterium. There's some discussion of the techniques used here:  http://fas.org/nuke/intro/nuke/heavy.htm.

Mike W.


(published on 05/06/2016)

Follow-up on this answer.