Sciencemadness Discussion Board - Endothermic freezing mixtures

Endothermic freezing mixtures

Formatik - 2-5-2008 at 00:09

This webpage states the following ratios of compounds yield a maximally achievable low temperatures:

NH4Cl............................30 g/100 g cold water..........................-5°C
NaNO3............................75 g/100 g cold water.........................-5
CaCl2·6H2O..................41 g/100 g ice......................................-9
KCl.................................30 g/100 g ice................................-10.5
Urea.......................... ..10 g/100 g ice....................................-10.8
NH4NO3....................14 g/100 g ice.....................................-13.6
NaNO3.......................15 g/100 g ice.....................................-13
NH4Cl........................25 g/100 g ice....................................-15.4
NH4NO3....................60 g/100 g cold water..........................-16.8
NaCl...........................33 g/100 g ice....................................-21.3
CaCl2·6 H2O................81 g/100 g ice....................................-21.5
NaBr............................66 g/100 g ice....................................-28
Ethanol.........................105 g/100 g ice .................................-30
CaCl2·6 H2O................123 g/100 g ice...................................-41
CaCl2·6 H2O................143 g/100 g ice..................................-55
KOH............................31 g/100 g ice....................................-63
H2SO4 (w = 66 %).........91 g/100 g ice...................................-90
MgCl2·6H2O................85 g/100 g ice....................................-94

Has anyone tried out colder ones like CaCl2.6 H2O, H2SO4, or MgCl2.6H2O?

I've tried CaCl2 (not the hexahydrate, anhydrous) about 143 g to 100 g cracked ice ratio and it didn't seem much colder than ice.

Without special equipment, what would be a good way to determine how cold those mixtures get? Maybe with the known boiling points of simple gases, before (l) -> (g), e.g. CO2, B.p. -78°. HC#CH, B.p.: -84°. HCl, -85°. CH3.CH3, -88°. N2O, -88°. CH3.CH2.CH3, -42°. NH3, -35°, etc. try and condense the gases.

If not then easier could be melting points of some common liquids before (s) -> (l).

[Edited on 2-5-2008 by Schockwave]

ShadowWarrior4444 - 2-5-2008 at 00:19

Hmmm, mixing CaCl2 with ice should result in an exothermic response due to the Heat of Hydration--this is why it is used in de-icing salts.

What you have here would seem to be a list of salts that decrease the freezing point of water to the specified temp. (Aside from the ones that arn't salts, that is.)

[Edited on 5-2-2008 by ShadowWarrior4444]

Formatik - 2-5-2008 at 01:49

Quote:

Originally posted by ShadowWarrior4444
Hmmm, mixing CaCl2 with ice should result in an exothermic response due to the Heat of Hydration--this is why it is used in de-icing salts.

Right; it warms up, I also noticed just adding the CaCl2 to regular water causes the solution to get very warm (heat of solution: -162 cal/g).

H2SO4 also heats up with H2O, but here they are saying it gets even much colder (with ice).

S.C. Wack - 2-5-2008 at 03:09

In my crude and small scale experiments using thermocouples, styrofoam cups, CaCl2 hexahydrate from big bag of Prestone ice melt and boiling water, and snow, the lowest temp that I got was -38C IIRC. This was after cooling the snow and salt to -10C before mixing. Less cooling of the components produced a corresponding result, it seems that there is just a certain number of degrees of cooling and that's that. So it all depends at what temperature that you start out at. I tested all of the different ratios stated in the various literature and there was no difference in my hands.

12AX7 - 2-5-2008 at 06:53

Yes which is why it specifies CaCl2.6H2O. Ditto MgCl2, which I think is more hygroscopic.

Tim

ShadowWarrior4444 - 2-5-2008 at 10:02

What is the mechanism behind the cooling of water in this method? (What does the thermal energy go toward.)

Does the presence of a hygroscopic salt catalyze the melting of ice, thereby increasing the rate at which thermal energy is absorbed?

(Note: a slight bit of research indicates that this is true.)

Though some of the substances in the list call for cold water instead of ice--what is the mechanism behind this? Is energy required to dissosciate the ions?

[Edited on 5-2-2008 by ShadowWarrior4444]

12AX7 - 2-5-2008 at 13:55

The heat is used up by solvation energy. Solid ice acts as a "cold reservoir" and dissolves in the solution, cooling it. Apparently, some solubles break up ice, absorbing what was crystalline binding energy. Likewise, some materials (urea, ammonium nitrate, etc.) become very cold on dissolving, as the crystal structure is broken up.

I'm amazed KOH is on the list; NaOH becomes very hot when it dissolves! I can't imagine any big difference between the two salts that would make such a diametrically opposed result.

Tim

ScienceGeek - 2-5-2008 at 17:12

Funny the reaction of Ammonium Nitrate and Barium Hydroxide Octahydrate isn't mentioned.

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Equation:

Ba(OH)2•8H2O (s) + 2NH4NO3 (s) → Ba(NO3)2 (aq) + 2NH3 (g) + 10H2O (l)

This better demonstrates the enormous increase in entropy

12AX7 - 2-5-2008 at 18:14

Hmm, barium nitrate isn't too soluble (I don't remember how soluble the hydroxide is). And ammonia would be dissolved in the water, for the most part. Certainly that's why it makes a paste. I guess most of the entropy comes from freeing the water of hydration and solvating the ammonia.

I've seen that reaction in person, it's quite frosty.

Tim

MagicJigPipe - 2-5-2008 at 18:20

Seems like the CaCl2 method could be a simple replacement for acetone/dry ice for condensing ammonia, methylamine or chlorine.

I wonder how practical this would be in the lab.

Perhaps the barium hydroxide and ammonium nitrate method would get colder if cooled before hand?

[Edited on 5-2-2008 by MagicJigPipe]

chemrox - 2-5-2008 at 22:10

My CaCl2 solution is in the freezer at -14*C which is about where I want it for the use I have in mind. I used really cheap dessicant from an RV supplier.

Formatik - 5-5-2008 at 12:49

Quote:

Originally posted by S.C. Wack
In my crude and small scale experiments using thermocouples, styrofoam cups, CaCl2 hexahydrate from big bag of Prestone ice melt and boiling water, and snow, the lowest temp that I got was -38C IIRC. This was after cooling the snow and salt to -10C before mixing. Less cooling of the components produced a corresponding result, it seems that there is just a certain number of degrees of cooling and that's that. So it all depends at what temperature that you start out at. I tested all of the different ratios stated in the various literature and there was no difference in my hands.

In this website they could also get only -37.7 deg. with CaCl2.6H2O (no ratios mentioned), they mention that the literature says -55 deg. is obtainable, but that even by several attempts they could not get it this cold.

Quote:

Originally posted by ShadowWarrior4444
What is the mechanism behind the cooling of water in this method? (What does the thermal energy go toward.)

The difference between salts and the hydrates, is that there is little to no heat from reaction with H2O. An explanation I've seen for the cooling is that the crystalline structure still requires energy to break down and that warmth from the surroundings is absorbed and so the temperature sinks. MgSO4.7H2O could also get pretty cold.

Note: the salts "melt" ice, not so much from heat, but because it depresses the feezing point, in other words, it allows the water to remain liquid at sometimes much colder temperatures.

Quote:

Originally posted by 12AX7 I'm amazed KOH is on the list; NaOH becomes very hot when it dissolves! I can't imagine any big difference between the two salts that would make such a diametrically opposed result.

I've tried the KOH and ice ratio and it gets pretty cold, colder than ice judging by touch of the beaker.

I tried condensing NH3 gas with it, but it might need some pressure in order for the gas to not escape out of the test tube it was going into and so that the gas can be brought into contact with the cold much better.

S.C. Wack - 5-5-2008 at 14:38

They didn't say anything about precooling. The lowest temperatures at least with CaCl2 require it. I have no doubt that the lower temperatures can be attained, but this will require precooling to -20-30C.

I doubt their claim that the cooling with ether and dry ice is ratio-dependent as ether is clearly just a medium for forming a slush bath.

Formatik - 5-5-2008 at 16:08

Quote:

Originally posted by S.C. Wack
They didn't say anything about precooling. The lowest temperatures at least with CaCl2 require it. I have no doubt that the lower temperatures can be attained, but this will require precooling to -20-30C.

No they didn't. Good to know.

Quote:

I doubt their claim that the cooling with ether and dry ice is ratio-dependent as ether is clearly just a medium for forming a slush bath.

They were just speculating there.

jimmyboy - 5-5-2008 at 16:21

The list you have doesn't seem like endotherms at all?? - just a list of ionic salt solutions and their freezing points..

maybe I am misreading your question

There is a big difference between solutions with a low freezing point and an endothermic reaction

[Edited on 6-5-2008 by jimmyboy]

Formatik - 5-5-2008 at 18:47

This thread is about endothermic mixtures (achieving lower temperatures), freezing point depression was a secondary but still somewhat related item of mention.

Formatik - 5-5-2008 at 21:21

I've finally found an old reference that mentions the hexahydrate (Römpp's Chemie Lexikon, 7th ed., p. 1677) for -55 deg. CaCl2.6 H2O the ratio given is 59 g of it to 100 g ice. A few others mentioned here:

BaCl2............22 g/100g ice..........-7.8
NaCl..............26 g/100 g water.....-10
CaCl2.6H2O...71 g /100 g water...-12
NH4SCN........57 g/100g water...-16
(NH4)2SO4.....38g/100g ice......-19
NaBr..............39 g/100 g ice.... -28
MgCl2............22 g/100 g ice....-33
H2SO4 (66.1%)...48g/100g snow...-37

And the CO2: solid CO2 with: alcohol: -72; chloroform: -77; acetone: -86; ether: -100.

They also mention using a Dewar flask (thermos can should work) in order to prepare such mixtures to protect against external warmth influences. Like one of the webpages emphasizes, there is some variation in the reports of these mixtures. Ether and dry ice got -82.5 by the experimentors in the previous website. Though they used a basic shallow glass container for cooling, maybe with a Dewar flask it could have gotten even colder.

Formatik - 6-5-2008 at 20:52

I've tried the 91 g 66% H2SO4 and 100 g ice ratio. Before adding the acid to the ice, the acid was cooled to -3° in the freezer overnight and the bottom half of the glass container wrapped in aluminium. After adding the acid to the ice chunks, it got cold within moments and reached the maximum low temperature of the thermometer (-20°C). I wasn't able to liquify propane with it. After a while the temperature rose to -16°C.

I also added 40 g MgSO4.7H2O to 100 g ice chunks, then stirring to see how cold it would get and the lowest was -8°C. Another 40 g and stirring brought it down to -9°C. Finally, 70 g more added and it went back up to -8°C. Nothing special that a simple NaCl/ice bath can not reach.

DJF90 - 7-5-2008 at 07:13

One of the important bits of information about an ice bath is that the ice should be crushed rather finely and then mixed with the correct amount of salt. Adding the salt to chucks of ice will work, but is nowhere near as effective.

[Edited on 7-5-2008 by DJF90]

StevenRS - 7-5-2008 at 19:18

I tried the H2SO4/ice mix, and used shaved ice (very fine). It got very, very cold, much colder than my thermometer would go. It condensed SO2 no problem, and it barely froze isopropanol alcohol. Thats around -90C, I think. Also, the sulfuric and ice were around -13 to start with.

I wonder if you could use different solvents, and different salts to condense liquid air? Use water/salt to freeze a gas like H2S, then use the frozen H2S (or something else) to freeze another gas, mix it with a salt, and so forth to very, very low temps?

I know the number of steps involved would be large, and some device to hold all the mixtures, each with reducing volume, would be hard (if not impossible to build) but imagine the rewards!

Liquid O2 made with reusable solvents, salts, and a freezer!

Imagine the fun you could have with that...

Solid ozone... even better....

Any thoughts?

Formatik - 7-5-2008 at 20:12

Here's a result of the ethanol: 105 g EtOH (200 proof, denatured with isoPrOH) mixed with 100 g pounded ice. This mixture went below the maximum readable low of the thermometer (-20°C), stayed there for a while and later rose. 100 g acetone with 100 g pounded ice also went beyond -20°C, but it did not stay there as long as the EtOH. Both compounds were precooled in the freezer prior to mixing. Methanol, isopropanol, acetonitrile, or potentially better even more volatile substances such as tetrahydrofuran, diethylamine, or ether probably yield similar results.

I've also found an old reference called Elements of Chemistry (1842) that mentions precooling also, here they talk about cooling with other freezing mixtures as a form of precooling to attain the lowest temperatures of the calcium chloride hexahydrate. They also recommend snow or pounded ice.

Another chiller is conc. HNO3 and ice: 1 part HNO3 (60%) with 1 part snow can get as low as -56ºC (source: Photographische Chemie und Chemikalienkunde (1896), p. 26-27).

From the same source: 1 part of conc. HCl (d= 1.18) and 1 part snow: -37ºC. And also the sulfuric acid: 1 part H2SO4 (diluted by some water), 3 parts snow: -50ºC. 5 parts crystallized CaCl2 and 4 parts snow: -50ºC.

Moderatley diluted H2SO4 when poured onto snow can reach a low of -40 to -50ºC (Grundriss der Experimentalphysik und Elemente der Chemie (1896), p. 253).

From another source (Lehrbuch der klinischen Osmologie als funktionelle Pathologie und Therapie (1902), p. 301): 1 part by mass water and 1 and ½ parts by mass powdered KSCN reaches -34.5ºC. For comparison: 1 part snow and 1 part NaCl (by mass): -21.3ºC. 1 part snow and 3 parts crystallized CaCl2 (by mass): -33 º. 1 part by mass of snow and 1 part by mass of diluted H2SO4: up to -50º.

And lastly, from another reference (Lehrbuch der chemischen Technologie der Energien (1906), p. 227-8): 2 pbw (parts by weight) dilute H2SO4 with 3 pbw. of snow or ice: -30 deg.; 5 pbw. NaCl and 5 pbw. NH4NO3 with 12 pbw ice or snow: -32ºC; 10 parts H2SO4 (65.3% conc. at 0 deg.) with 11 parts ice or snow: -37.0 deg.C.; 1 part HCl acid (spec. gravity: 1.18 at 0 deg.C.) with 1 part snow or ice: -37.5 deg.C.; 4 parts potash with 3 parts ice or snow: -48.5ºC; 1 part HNO3 (60% N2O5 at 0 deg.) with 2 parts ice or snow: -56.0 deg.

There were also some others, like 6 pbw. Na2SO4 with 4 pbw NH4Cl and 2 pbw. KNO3 solubilized in 4 pbw of dilute HNO3 goes from +10 to -23ºC or 6 parts by weight (pbw) Na2SO4 and 5 pbw NH4NO3 solubilized in 4 pbw dilute HNO3 goes from +10 deg. to -40 deg.C., etc.

More specs on the hexahydrate:

From: Elements of Chemistry (1877), p. 368.

This following is from an older CRC handbook about the proportions of the hexahydrate or acid:

Gathering from that, it seems like that what should be used for the hexahydrate at least is snow instead of ice (* or finely pulverized ice); ratio for the -55º: 1 to 0.70, thus 143 g added to 100g snow. So far everyone’s suggestions have been excellent (e.g. precooling, pulverized ice, SO2 testing).

Quote:

Originally posted by StevenRS
I tried the H2SO4/ice mix, and used shaved ice (very fine). It got very, very cold, much colder than my thermometer would go. It condensed SO2 no problem, and it barely froze isopropanol alcohol. Thats around -90C, I think. Also, the sulfuric and ice were around -13 to start with.

Did it liquify or solidify? If it only liquified then it reached below the B.p. of -10°C, however if it solidified that would mean it went below -72.4°C (M.p. of SO2). SO2 would probably be the best way to test for lower temperatures because it easily liquefies, and then if it gets colder will solidify.

Quote:

I wonder if you could use different solvents, and different salts to condense liquid air? Use water/salt to freeze a gas like H2S, then use the frozen H2S (or something else) to freeze another gas, mix it with a salt, and so forth to very, very low temps?

I know the number of steps involved would be large, and some device to hold all the mixtures, each with reducing volume, would be hard (if not impossible to build) but imagine the rewards!

Liquid O2 made with reusable solvents, salts, and a freezer!

I think it's in the realm of possiblity to find a mixture at the extreme end of endotherms as there are in exotherms - and that's what this thread is all about, but these seem to me to have been less studied. Thermochemistry seems more interested in the hotter spectrum.

Quote:

Imagine the fun you could have with that...

Solid ozone... even better....

Any thoughts?

Imagine being able to make your own dry ice. :cool:

[Edited on 8-5-2008 by Schockwave]

StevenRS - 8-5-2008 at 03:41

Neither became truly solid, the SO2 was not cooled as well at all as the Isopropanol, which started to become very viscous, and started to freeze.

Formatik - 8-5-2008 at 22:28

I've done a couple more. First, 20 g of finely powdered NaOH was added to 90 g of finely crushed ice in a beaker wrapped in Al foil, then quickly stirred. This mixture would go about to -10ºC, but not below. Next, the same was done except 20 g powdered KOH was added to 90 g of finely crushed ice; this went as low as -16ºC. Neither hydroxides were precooled.

Not last but not least, 100g of around 54% HNO3 (precooled in the freezer) was added in the same manner to 100 g of finely crushed ice. By adding the acid, the temperature went immediatley below the maximum readable of about -20ºC. Later it rose to about -20ºC, and more.

And lastly, an addendum to the ethanol: adding 100 g pounded ice to 105 g ethanol (denatured) which was not precooled in a beaker got no lower than about -18ºC.

[Edited on 8-5-2008 by Schockwave]

Formatik - 16-5-2008 at 17:26

Condensing could take a while so it might not be very accurate for determining how cold a mixture gets beyond a thermometer. Even about 3.5 mL 99% isopropanol in a glass tube put into around 200 mL liquid N2 (-196 deg.) took around over 1 minute to become solid. Right after melting (m.p. -89.5°C) the alcohol was a thick viscous liquid, similar to glycerin in consistency.

Liquid N2 freezing mixtures (and the source):

hexane / liquid nitrogen: -94 deg.
methanol / liquid nitrogen: -98 deg.
pentane / liquid nitrogen: -131 deg.

I've tried the methanol/N2 one as it's used to chill chemical reactions. When the liquid N2 is in excess the methanol is left behind as a frozen solid after the nitrogen has boiled off. When methanol is in excess then the N2 boils off and dances on it, similar to water. Methanol vapors also get pushed out of the container in the process.

nodrog19 - 16-6-2008 at 19:04

i found this table in Organic Laboratory Techniques, 3rd Edition

ice- water -0C
100g ice 33g NaCl -21C
70g ice 100g CaCl2*6H2O -54C
dry ice- ethanol -72C
dry ice- acetone -77C

Formatik - 8-11-2009 at 13:03

There's been some more talk on freezing mixtures over at the German forum, so I've decided to add some more information to this thread. Disregard the improbable values for MgCl2 and H2SO4.

More clarity on the acid mixtures comes from Gmelin: 3 parts snow and 1 part sulfuric acid containing 1/5 its mass of water reaches -32.5°. If the materials are precooled to -7°, then the temperature sinks to -51°. Since the 18th century similar mixtures of snow and dilute H2SO4 have been used to freeze mercury. Nitric acid mixes have also been used to freeze mercury. HNO3 with 1.5 H2O when mixed with two times its mass of snow, starting at 0 deg. causes a temperature depression of -56°.

A corroboration reference for mixing of snow with absolute alcohol causing a depression of 0° to -30° is Handbuch der biochemischen Arbeitsmethoden by E. Abderhalden. If precooling in a freezer, I feel the need to reiterate never to leave a flammable unstoppered, so no explosion can result.

Concerning the mixture of ether and dry ice. After looking through old literature, this one was a favorite freezing mixture. Being able to reach temperatures occasionally colder than the dry ice itself, and a drop or two of the mixture is said to be able to produce blister burns on the skin. But not normally reach near -100°.

In Gmelin it's stated that Faraday with solid CO2 and ether got -77°, for some strange reason B. Schwalbe got a wide range, from -77 to -97°. Anyways, I think the value of -100° might actually come from Faraday achieving -103° using an airpump for the same mixture, and Schwalbe reaching -110° also using an air pump. E. Abderhalden states solid CO2 and alcohol in a vacuum causes the temperature to sink below -100°. The increased evaporation being responsible for the heightened cooling.

I oversaw the reference earlier in Elements of Chemistry, stating some guy named Walker did experiments using freezing mixtures against other freezing mixtures to cool even further. With 8 parts snow and 10 parts dilute H2SO4 (I'd like to know exactly how dilute) when frozen (using CaCl2-mixes) goes from -54.4° to -68.3°. A sulfuric acid-water freezing point curve is below.

[Edited on 9-11-2009 by Formatik]

The ratio

Formatik - 11-11-2009 at 16:15

A system of chemistry, Vol.1 by J. Murray, who seems to have seen the original reference of Walker, states the ratio is 10 parts sulphuric acid diluted with half its weight of water, with 8 parts snow. In total, this is about 28% conc (d= 1.20). What this mixture has reached, by the old literature, it is said to have been the coldest temperature measured (well, without dry ice, liquid gases, compression devices).

densest - 17-11-2009 at 15:56

A type K thermocouple (the most common type AFAIK) is good down to liquid air temperatures, and the chromium-nickel alloys are poor heat conductors. Many inexpensive (less than $US 50) multimeters come with a bead thermocouple with teflon insulated wire. Such a system is very useful to measure cooling mixtures. A sleeve made of thin teflon tubing closed at one end could protect the thermocouple from salts & acids. To make such a sleeve, heat some teflon tubing until it softens and starts to become transparent or translucent. Remove it from the heat and squeeze it gently but firmly closed with pliers, a clamp, or a vise. Don't squeeze it to a thickness below about 1 1/2 wall thicknesses. The tubing will fuse together giving you a condom for your thermocouple.

When snow falls where I live I'll experiment more with freezing mixtures... the Canadian members should have access to some good OTC chemicals for this purpose

I have some teflon tubing with ID about 1/8" (3 mm or so) which I would be willing to cut to 10-30 cm lengths and mail to people for postage + a trivial amount.

Sedit - 17-11-2009 at 17:25

Formatik, using snow which I feel is essential due to the orderness of bonding being broken all at once, never once reached temperatures no matter what concentration I tryed. HCl and snow would out do H2SO4 everytime reaching -40s IIRC(I have a thread here or at science forums somewhere). I tryed shaving ice but got nothing. I attempted to condense ammonia with the HCl/snow mix but got limited results. I was hoping someone was able to have success with CaCl2 hydrates as this looked very nice indeed.

Might I suggest if you get snow before me to attempt a concentrated solution of CaCl2 on snow instead of trying to get the two powders to mix quick enough. I have had a strong suspicion all year that this may achieve the results documented.

Formatik - 17-11-2009 at 22:14

Quote: Originally posted by densest

A type K thermocouple (the most common type AFAIK) is good down to liquid air temperatures, and the chromium-nickel alloys are poor heat conductors. Many inexpensive (less than $US 50) multimeters come with a bead thermocouple with teflon insulated wire.

Unfortunately, I found out about thermocouples and multimeters too late.

Quote:

Such a system is very useful to measure cooling mixtures. A sleeve made of thin teflon tubing closed at one end could protect the thermocouple from salts & acids. To make such a sleeve, heat some teflon tubing until it softens and starts to become transparent or translucent. Remove it from the heat and squeeze it gently but firmly closed with pliers, a clamp, or a vise. Don't squeeze it to a thickness below about 1 1/2 wall thicknesses. The tubing will fuse together giving you a condom for your thermocouple.

Be careful heating teflon, since very toxic gases especially perfluoroisobutene are released if it's overheated/decomposed.

Quote:

When snow falls where I live I'll experiment more with freezing mixtures... the Canadian members should have access to some good OTC chemicals for this purpose

If no snow is available, one could also try with pounded, pulverized ice that isn't wet.

Quote: Originally posted by Sedit

I attempted to condense ammonia with the HCl/snow mix but got limited results. I was hoping someone was able to have success with CaCl2 hydrates as this looked very nice indeed.

For something like liquefying ammonia, the literature says to use large amounts of cooling mixture. For saftey reasons, the CaCl2 hydrate might be the better idea.

Quote:

Might I suggest if you get snow before me to attempt a concentrated solution of CaCl2 on snow instead of trying to get the two powders to mix quick enough. I have had a strong suspicion all year that this may achieve the results documented.

There's an interesting reference I've found on this: Über die Kältemischung aus Chlorcalcium und Schnee (= concerning the freezing mixture of calcium chloride and snow) by Dr. Hermann Hammerl. (from: Sitzungsberichte der Kaiserlichen Akademie der Wissenschaften, Vol. 78, Pt. 2. 1879. Pgs. 59 to 79). This is a somewhat lengthy paper on just the cooling mixture of calcium chloride and snow alone. I have seen quite some variation when it comes to the ratio of the mixture, but this paper is well detailed. So I trust the ratio given there which is: 1 gram CaCl2.6H2O with 0.70 gram of snow, for lowest cooling value and temp. of -54.9°C. On pg. 74, and of this temperature he says: "in order to always reach this temperature minimum from mixing snow and calcium chloride, most importantly the crystallized calcium chloride is in a finely powdered form, and if possible be cooled below 0°C. The snow should not be wet, but crumbly, and also cooled below 0°C." He said that he and his colleges - through experimentation - always got below -50°C when mixing in the proper proportions.

Freezing point depression

nickd - 18-11-2009 at 06:15

Whilst there is some some variation in freezing point depression with different substances, you may remember that freezing point depression is a property that dependes on the RATIO of the solute to the solvent particles rather than the nature of the actual substance itself
.The depression effect of freezing point changes are related to the number of particles and does'nt involve solute /solvent interaction. (It is related to the molality of the solutions as represented by Raoult's law)

Formatik - 18-11-2009 at 17:25

These mixtures involve several phenomena. For one of the reasons why the temperature drop occurs in the chloride mixture- the enthalpy of solution - solvent-solute interactions do have an influence.

DJF90 - 20-11-2009 at 09:04

Nickd: You are *partially right*, in that the depression effect of freezing point (and also temperature dependance of solubility) are related the the ratio of the solute/solvent (i.e. the mole fractions). However, this is only true when you consider the system as an *ideal* solution; one where the enthalpy of mixing is zero, and the gibbs free energy relies only on the entropic term.

However, this is only a model. A step up would be to treat it as a "regular" solution, which has an "ideal" entropy of mixing and a non-zero entalpy of mixing.

Even better would be to use the "electrolyte" solution model, which is based on Debeye-Huckel theory.

Panache - 15-12-2009 at 06:55

just a stupid thought i always had regarding ice salt baths, what would happen if the ice and salt you were combining were already at a temperature below -31. Would the system warm to -31, perhaps this is where an intelligent design advocate needs to chime in and let us know.

Sedit - 4-1-2010 at 21:51

I tryed adding finely crushed CaCl2 from damprid that had been moistened and sat out for a while till it reclumped where it was grounded down again till it reclumped.

I ground it once more and added it to a bowl full of fine ice. It use to be snow but its been sitting in the cooler for a week for I guess its best to just say fine ice at this point because thats how it would behave where as snow would get alot lower I believe.

Anyway one with the shitty data. I did not keep all numbers since im just playing but after the addition the lowest temperature I was able to achieve was -31C. I now see that judging from the chart up there that importance over the amount of CaCl is a must so I guess I will take better measurements next time. My main goal was to test the temperature since I want to rerun the Li[NH3]4 experiments and thats seems cold enough to be effective. I just didn't want to mess around with another HCl/Ice bath for obvious reasons. It was good that at one point there was small leak in the NH3 generator and the small amounts of HCl around quickly showed it but I still want nothing to do with a bowl of HCl even if its been diluted.

An Unusually Philosophical Post

S.C. Wack - 24-10-2010 at 14:03

By Mr. Richard Walker:

Observations on the best Methods of producing artificial Cold.
Phil. Trans., 270 (1795)

On the Production of artificial Cold by Means of Muriate of Lime.
Phil. Trans., 120 (1801)

Attachment: phil_trans_270_1795.pdf (1.8MB)
This file has been downloaded 980 times

Attachment: phil_trans_120_1801.pdf (757kB)
This file has been downloaded 901 times

Random - 24-10-2010 at 15:29

How low the temperatures could go if I would put NaCl in freezer at -20 and take it out after 2 days and then mix it with -10 ice. Could it go below -20 then?

not_important - 25-10-2010 at 13:18

Freezing mixtures generally do not go colder than the eutectic, and the NaCl-H2O eutectic point is −21 C.

Formatik - 12-12-2010 at 15:48

Quote: Originally posted by S.C. Wack

By Mr. Richard Walker:

Observations on the best Methods of producing artificial Cold.
Phil. Trans., 270 (1795)

On the Production of artificial Cold by Means of Muriate of Lime.
Phil. Trans., 120 (1801)

Wow, that's pretty interesting. What Walker is saying in the second document is that below -91 has been reached, being held there for 20 seconds. The -68 mixture was also actually 8 parts conc. H2SO4, 4 parts water, 1 part grain ethanol cooled to air temperature. Then he took 10 parts of this mixed acid, mixed it with 8 parts snow and chilled it with the CaCl2 mixture. Walker is probably some kind of original source for the idea H2SO4 and ice gets -90 C, mentioned on the first post in this thread.

Also I think interesting, the effect of precooling on calcium chloride mixtures, described on p. 133. It doesn't matter if snow or ice powder are used either, they reach the same temperatures in the mixtures.

I haven't had much time for experimentation, but one thing I did do a while back was to condense propane with dry ice. Then adding dry ice to the propane, and attempted to freeze 91% isopropanol, but it only got viscous after a couple minutes. The liquid propane was dangerous to handle not only because of flammability and stench, but because it causes severe burns. A drop or so of the liquid propane on my hand caused a severe frostbite burn that took months to heal. I thought since it was claimed SO2 and dry ice get below dry ice itself, if I were to take something more volatile it could yield similar or even better results.

vulture - 13-12-2010 at 15:18

I found that mixing dry ice with iPrOH, acetone, pentane and diethylether all stops at -78C. There seems to be no difference whatsoever.

blogfast25 - 15-12-2010 at 14:09

Quote: Originally posted by DJF90

One of the important bits of information about an ice bath is that the ice should be crushed rather finely and then mixed with the correct amount of salt. Adding the salt to chucks of ice will work, but is nowhere near as effective.

[Edited on 7-5-2008 by DJF90]

Yes, absolutely: fineness of the mixtures is everything.

Another very common exothermic reaction: neutralisation of NaHCO3 with any strong acid. An entropic effect apparently due to the escaping CO2...

Formatik - 17-12-2010 at 23:49

Some more action from me. I've quantified it well compared to last attempts from some time ago. I repeated the mixture of dry ice and liquid propane mixed in a plastic thermos can, and after mixing them I got a reading of -65 C (Note: I condensed the propane with some dry ice, and then added some more dry ice). It didn't get much colder than that for some time, so I used an aquarium air pump set on high and used a tube to bubble air through it, and this caused the temperature to drop rapidly. After maybe something like 30 minutes of bubbling, I was able to eventually get -80 C. But, the same can be done with acetone.

I did the same thing using acetone. 158g solid CO2 pieces and 102g acetone when mixed in a plastic thermos can started out at -64 C, and it only decreased slowly after one minute to -65 C, then another minute to -66 C. So, then I bubbled in air with the same pump and this promoted temperature drops. The decrease was by 3 C within about a minute, then by 1 C per minute for 6 minutes to get to -75 C. Then after bubbling in air an additional 17 minutes, the temperature eventually dropped to -80 C. Then the reading was fluctuating slightly between -81 and -80 C, and it was like this up to 18 more minutes, after which I added 10g X 2 dry ice pieces as I kept bubbling an additional 13 minutes after dry ice addition, after which time it then reached -83 C. After 5 minutes, I added 28g solid CO2 during the bubbling. And then temperature readings started fluctuating and increasing. But I then removed the air pump. And took the multimeter (it was on the cold concrete ground of the side garage). As I moved it to a warmer area in the house, I noticed colder readings that went below -83 C, I later took the air pump to this mixture and just let air hit the surface/go over the liquid instead of bubbling in, and I was able to get steady readings of maximally -88 C, some fluctuations went below but these were due to moving the multimeter during measurement.

I used a miserable air pump source. I think a better pump or even a vacuum, ought to produce promised literature values.

I wasn't sure if I were to be able to get colder than dry ice (-78 C) alone at first, but the trick to cool down dry ice freezing mixtures faster by blowing air through it seems to work pretty well, this could be useful to cool those kind of mixtures much more rapidly.

Liquid nitrogen can possibly be made colder on similar principle by blowing anhydrous helium or hydrogen through or over it.

Some pictures are below.

Liquid propane mixed with dry ice and air bubbling in.

Propane and dry ice multimeter reading.

The set-up of acetone and dry ice.

acetoneandCO2reading.png - 55kB

Acetone and dry ice after bubbling air through it for under 30 minutes.

Steady cold readings.

[Edited on 18-12-2010 by Formatik]

peach - 12-4-2012 at 11:24

I'm not sure if it has been pointed out in the thread thus far, but to achieve these really low temperatures using frigorific mixtures, it is important that they are being cooled by a supplementary source.

Quote:

These artificial processes for generating cold are much more effectual when the materials are previously cooled by immersion in other frigorific mixtures.

- Elements of Chemistry

This means, for instance, that the sulphuric and ice needs to be sat in a second container of say, KOH and ice. I have tried the sulphuric method three or so times before, with it sat in a dewar, and never been able to get the temperature below that of the ice when it came out of the freezer.

I tried again today with KOH and ice, as follows:

A Thermos, it's inner and outer cap were placed in the freezer, along with 100g of crushed ice and 31g of KOH (all kept as separate components). They were removed hours later, when the inside of the flask was reading -20C. The thermometer was poked to the base of the Thermos, the cold, measured 31g of KOH dumped in, then the 100g of crushed ice, fresh out of the freezer. I put the cap back on and gave it a good shake, checked the thermometer was at the base and replaced both the inner and outer caps.****

I then sat and wrote down the temperature on a minute by minute basis for half an hour.%%%

These are the values I obtained (in Celsius), per minute:

-20 (0 minutes, freezer temperature)
-14
-3
4
7
2.8
-0.6
-3.2
-5.4
-6.8
-8.6
-9.8
-10.8
-11.9
-12.6
-13.5
-14
-14.5
-14.9
-15.5
-15.8
-16.1
-16.4
-16.6
-16.8
-17.2
-17.3
-17.5
-17.6
-17.6
-17.6

As the trend was now stabilising, I left it for much longer periods of time. Here are the results that occurred:

45 minutes
-18.7

66 minutes
-17.7

75 minutes
-17.4

158 minutes
-16.9

207 minutes
-15.5

226 minutes
-15

321 minutes
-12.7

423 minutes (some 7 hours after combining the two)
-10.2

The mixture rapidly rose in temperature when shaken together, then (soon after) fell in a de-accelerating curve for half an hour. From there, it has sat in a very gradual linear rise back to room temperature. I've included the numbers such that you can plot them yourself in Libre / Open / Microsoft Office if you so wish.

Enthalpy of fusion (water): 6.01 kJ/mol
Enthalpy of solution (KOH): -57.61 kJ/mol

This is a 31g / 100g KOH to ice mix, as specified in the table, and measured on a digital kitchen scale (reasonably accurate for this task).

100g of H2O: 5.6 molies
31g of KOH: 0.55 moley moles

Enthalpy of fusion for 100g H2O: 33.6 kJ
Enthalpy of solution for 31g KOH: -31.6 kJ

Those numbers are close, the KOH is giving out almost as much heat as the ice absorbs in melting (based on their normal melting and dissolving figures as separate entities). The mixture as a whole, based on those values, gives a total heat absorbing capacity of 2kJ, versus the 5.6 kJ the ice alone would have had at 0C, melting on it's own; 2 to 3 times less.

From the graph, it is clearly evident that the test is not wanting of patience, as the linear rise means the process is now complete - there will be no more cooling. I wanted to point this out as, when I first saw the numbers in the table, I was optimistic I could have some fun with those, in a practical sense. However, the process is not particularly applicable to practical work (e.g. running condensers or baths), because those baths do need stacking. With the mixtures I've tested not even being able to get below -20C, they are actually performing worse than regular table salt / ice baths, and using more expensive and corrosive depressants. On top of that, my freezer can manage -30C. And, on top of that, the baths (unlike dry ice and LN2) do not just need topping up, they need emptying and refilling (not something you'd want to be doing if cooling something that is going to go bananas when the bath / condenser is restacked).

Following the logic of bath stacking, with the freezer at -30C and the baths sat in there, one could start with the two CaCl2 baths and then stack into these the KOH bath. I suspect that would reach (or at least get closer to) the specified -63C.

I believe these mixtures were originally developed more as temperature reference standards, perhaps for applications such as calibrating thermometers. I commend the original researcher, but they not a substitute for dry ice where heat loads must be absorbed.

I also believe that a much more profitable expenditure of time would be to simply freeze an ethanol mixture in the freezer, to gain a -30C slush. There are some long chain alcohols, I can't remember which, that solidify as native species in this temperature region, and absorb a fair amount of heat in their state change (I think I found these on matweb or in the Rubber Bible, the CRC book). There are also metal alloys that may be of some amusement.

***Photos
%%%Graphs. Image shack just imposed a 500 photo limit on their free accounts. As mine had over a thousand in it, quite a few of my photos have been replaced with 'violated rule' tags. I've uploaded these as rough pictures, I don't know how long those will stay on there.

Formatik - 23-4-2012 at 22:56

Quote: Originally posted by peach

I'm not sure if it has been pointed out in the thread thus far, but to achieve these really low temperatures using frigorific mixtures, it is important that they are being cooled by a supplementary source.

It was mentioned in several places, but yes, precooling is important and crucial to obtaining literature lows with frigorific (endothermic) mixtures. I've also found this to be the case.

It was also mentioned the effect of increased evaporation being able to cool the mixture even more by blowing air through it or using a vacuum. Though hasn't been discussed with relevance to endothermic mixtures (just freezing mixtures involving dry ice). But I have been able to cool a regular hydrated MgCl2-ice mixture a few degrees more just by blowing air through it.

Quote:

I tried again today with KOH and ice, as follows:

A Thermos, it's inner and outer cap were placed in the freezer, along with 100g of crushed ice and 31g of KOH (all kept as separate components). They were removed hours later, when the inside of the flask was reading -20C. The thermometer was poked to the base of the Thermos, the cold, measured 31g of KOH dumped in, then the 100g of crushed ice, fresh out of the freezer.

....

45 minutes
-18.7

I got something similar, but slightly warmer using a mixture of: 20g fine KOH and 90g finely crushed ice reaching -16 C, when neither were precooled.

Quote:

Following the logic of bath stacking, with the freezer at -30C and the baths sat in there, one could start with the two CaCl2 baths and then stack into these the KOH bath. I suspect that would reach (or at least get closer to) the specified -63C.

Right, the data is not so straight-forward as it as been put on some tables. I also think the KOH mixture could reach that low by cooling it with other freezing mixtures. Though I have pulverized KOH in one of my cooling experiments and that was no fun. Extremely irritating in dust form.

I suspect if there's any truth to it (the only reference I think I've seen for it is the webpage first mentioned in this thread), that this is the case for MgCl2.6H2O-ice mixtures being able to reach the purported lows. In which case it would have to be cooled with other cooling mixtures, or basically just dry ice to potentially reach temperatures below that of dry ice.

Quote:

I believe these mixtures were originally developed more as temperature reference standards, perhaps for applications such as calibrating thermometers. I commend the original researcher, but they not a substitute for dry ice where heat loads must be absorbed.

This has been more of an exploratory thread for the sheer experimentation and knowledge. But if I was to use a cooling mixture my first choice would be a cheap salt and ice (or snow), dry ice with or without solvent (then air blown through it to cool it down faster and better), or liquid nitrogen.