Antwain
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Question about thermodynamics
Ok, so say you have ice at -5 or 0*C. You use this to cool something and just in changing phase it takes up 6kJ/mol.
Thats quite a bit, considering that heating that same liquid water to ~70*C (paper towel calc.) will be necessary to make it absorb that much more
energy.
So what I am wondering is this...
If you put salt with that ice it will eventually form a solution. It may take ages but it should eventually dissolve. Now, dissolving salt in water
is not an exothermic process (especially) and so what I am wondering is whether a cold salt solution somehow has a very low enthalpy or if instead the
process of the ice melting even below zero is where all the energy comes from (ie. from the surrounds). Further, will a salt water mixture that has
been put in the freezer be as effective as removing heat as ice cubes from that freezer?
I realise that Ice/salt water mixtures are usually used, and that I may have just answered my own question but I still can't rationalise it. Phys chem
was never my strong point.
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JohnWW
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As well as mixing common salt NaCl) with ice at 0ºC or just below being endothermic, resulting in a drop in temperature of the resulting solution,
the dissolution of salt in it also depresses the freezing-point of the resulting mixture, by an amount proportional to the moles of ions dissolved (up
to the saturation point). Hence the result is a liquid solution of salt substantially colder than 0ºC. I think that temperatures down to -10ºC can
be obtained in the laboratory for cooling purposes by his means. However, a salt water solution frozen in a refigerator would not be any advantage,
because freezing expels all or most of the salt from solution, and besides the presence of salt in solution makes only a minor difference to the
specific heat. This is why pack-ice that forms on the surface of seas in high latitudes in winter, e.g. the Arctic Ocean, consist of fresh water.
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Antwain
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Ok, I didn't know it was endothermic, but clearly it is not *very* endothermic (think ammonium, nitrate or both). I didn't mean freezing the "freezing
mixture", just getting it very cold. you can go below -10*C, surely. Before they redefined the Fahrenheit scale (because it was stupid and arbitrary -
and wrong) 0F was supposed to be the coldest salt water mixture possible (of course because Mr F was a fool, as shown by his coming up with such a
stupid system, defined by body temperature for gods sake, he got it wrong by a few degrees). I don't have a 'feel' for the F system but since the
redefinition it can be worked out. -40C = -40F, 1C = 4/9F => 0F = -40 + (4/9)* 40 = -22.22
Sounds like the NaCL.2H2O barrier to me 
Anyhow, where was I before my rant? Yes, so if you have *not* frozen salt water at 0 or -5 that will not have the heat absorbing capacity of ice
then, if the specific heat is not grossly changed.
Presumably then, the ice is there to soak up the heat and the salt water is there because ice is a very good insulator and you want to conduct heat
out of stuff. Sound good?
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DerAlte
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You need a phase diagram to see this clearly. Sorry, I haven't got one for the NaCl/H20 system. Read up about the phase rule.
The easiest way to explain is: The water is initially ice at 0C or lower. The salt is added. At the interface the salt dissolves becuase salt and
water freeze at a lower temp due to the freezing point being lower than pure H2O.
In order for any solid to melt, heat is absorbed - the latent heat of fusion, 70 cal/g for ice to water. The mixture gets colder to allow the ice to
melt as it absorbs heat. It will do this until something intervenes to stop the process. At ~-13C a crystalline hydrate forms, NaCl.2H20, once the
salt solution is of sufficient density.
The formation of this gives out whatever the latent heat of formation is for that hydrate. It tends to stop the cooling. Further cooling will occur
only when the salt solution is sufficiently dilute to prevent this.
Assuming ice is in excess, using up all the salt will also stop the process. You now have an process going on in which the formation rate of the
hydrate competes with the cooling due to water melting on dissolving with the salt. Crystallization takes time and the solution can become
supersaturated. The temp can go lower, down to about -20C IIRC.
Another limit to the process is the separation of the eutectic of salt and water at -23C. This is probably the ultimate limit. Rather complex, I'm
afraid! And I didn't even mention entropy...
I was playing with some dry ice yesterday. It sublimes at -78C, and using acetone in a mixture, a liquid bath can be maintained at around that temp
until the solid CO2 evaporates. Liquid nitrogen used to be used for our temperature chambers for testing electronic systems.
Regards, Der Alte.
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kilowatt
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Here's a phase diagram for NaCl/H2O.
[Edited on 28-10-2007 by kilowatt]
The mind cannot decide the truth; it can only find the truth.
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Antwain
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I have to agree that CO2/acetone or N2 are far preferable to water ice, but outside uni its ice unless i can be bothered making a trip to BOC and
forking out $$$ for dry ice.... and while there is a 2x2x1m box subliming away at uni 
Also its coming into summer here so being able to cool mixtures will be more important.
So, in fact, adding ice at 0*C to salt will create a solution which is COLDER than 0*C?
would ammonium sulfate of chloride work better? the (impure) sulfate only costs a couple of times as much as NaCl
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DerAlte
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Ttry CaCl2 - it gets a bit lower that NaCl, IIRC. Der Alte
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vulture
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In theory, CaCl2 gets you to -53C. I couldn't get it lower than -38C. Which is still impressive though, quite the ice jacket forming on the beaker.
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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