Can a NaCl (aq) solution have a supercooling effect?

[Hammad Shahid]

TL;DR Summary

I had an aqueous solution of NaCl, and I observed it to have a super cooling effect.

Today we had a chem lab on colligative properties. The lab was focused on seeing how the freezing point of water would change w/ the addition of solutes.

Now, the first part was to measure the freezing point of water. This was done by placing an apparatus of DI room temp water into a salt-ice bath at -10C.

There was a probe inserted into the apparatus and connected to a LabQuest device. The data was plotted on the machine.

I quite clearly saw the supercooling effect, and then the temperature of the water returned back to 0C (and stayed constant).
This I understand is due to the difficulty of arranging into ice crystals in pure water.


Now, the textbook mentions that if there are impurities in water, then the ice crystals can form around the impurities (therefore, making the supercooling effect pretty much negligible).

However, this is where I ran into the problem:-
The second part was measuring the freezing point of the same volume of water with table salt mixed into it.

Now I wasn't expecting a supercooling effect here. To my surprise, the solution decreased to -5C and then rapidly jumped to -2C, where it stayed constant.

How do I explain this?

Thank you.

 

[Borek]

Now I wasn't expecting a supercooling effect here.

Why?

 

[Hammad Shahid]

Why?

From what is taught in class, if there are impurities in water, the water can form crystals around it easier.

But now I'm confused: We're learning about the freezing point depression, and apparently having solute particles interferes with the forming of crystals in solution?

 

[Borek]

I see where your confusion comes from. By impurity in this context we typically mean anything that makes the mixture nonhomogeneous, like small solid particles - any kind of dust for example. Dissolved salt doesn't fit, as the solution is homogeneous.

 

[Hammad Shahid]

I see where your confusion comes from. By impurity in this context we typically mean anything that makes the mixture nonhomogeneous, like small solid particles - any kind of dust for example. Dissolved salt doesn't fit, as the solution is homogeneous.

Ok, that makes a lot of sense. But then, why doesn't the dissolved salt also provide places to form crystals around? Is it too small?
And would the dust also not interfere with crystal forming, at least to some degree?

I am having trouble understanding how solute can both help form crystals and interfere with forming crystals.

 

[chemisttree]

And would the dust also not interfere with crystal forming, at least to some degree?

I am having trouble understanding how solute can both help form crystals and interfere with forming crystals.

Probably not true to say that solutes, at least ALL solutes lower the freezing point of water under all circumstances.

https://sciencing.com/raise-freezing-point-water-5211895.html

 

[TeethWhitener]

But then, why doesn't the dissolved salt also provide places to form crystals around?

No. The individual ions in a salt solution are solvated by water molecules. A heterogeneous particle such as a dust particle gives ions a surface on which to anchor, lowering the energy barrier for ions to escape their water solvation shells and begin the crystallization process.

 

[Nik_2213]

Long time but, IIRC, when I did something similar in phys-chem, we had to ensure there was a nucleation centre. IIRC, a gentle rub of a glass rod for a cm or two on inside of glass beaker *usually* sufficed. But you had to repeat this for each test...

A BIL who'd been involved in protein X-ray crystallography had some amusing tales about super-saturated solutions that were so 'clean', they totally refused to crystallise progressively. They'd sit for days, even weeks, then set solid overnight. Upside, microscopic inspection of slush, picking several nice crystals out of the mess, often provided viable 'seeds' for try#2...

The related process, super-heating at boiling point, is why 'anti-bumping' granules and beads were devised. Without such, a big beaker of heated water may do a splendid geyser burp-- Be Not There !

If you work very gently, a rotary evaporator may manage without. If, like me, you had to push lab's bigger rotary evaporator to the limit, urgently distilling dozens of litres of methanol for our auto-analysers because it had been delivered in the wrong drums, was now contaminated with drums' varnish, multiple anti-bump beads were essential...

 

[Hammad Shahid]

No. The individual ions in a salt solution are solvated by water molecules. A heterogeneous particle such as a dust particle gives ions a surface on which to anchor, lowering the energy barrier for ions to escape their water solvation shells and begin the crystallization process.

Ok, so if there are no particles at all, how to water molecules escape from the solvation shell in the first place? Wouldn't the water molecules be in a lower energy shell still solvated around the ions rather than crystallization?

 

[Nik_2213]

Please read the wiki on supercooling, with reference to crystal homogeneous nucleation.
https://en.wikipedia.org/wiki/Supercooling
Roughly translated, you can only super-cool so far and so long, before the liquid 'spontaneously' sets to an amorphous 'glass'.

IMHO, possible triggers include quantum tunneling, back-ground radioactivity and cosmic rays...