Sciencemadness Discussion Board - Salting out acetic acid?

Salting out acetic acid?

UndermineBriarEverglade - 4-11-2024 at 14:42

I have vinegar concentrate, 45% by weight. That's below the eutectic point for freezing, and I have no distillation setup. Could I salt it out to get past the eutectic point? If I understand this correctly, I would need a salt that's very soluble in water and insoluble in acetic acid.

I have seen extraction with DCM and a salt, but I don't have any DCM. I have acetone, isopropyl alcohol, and denatured spirits. Sedit salted out some amount of AcOH + IPA with Epsom salt but later moved on to DCM so I'm not sure if the method was fruitful.

Precipitates - 10-11-2024 at 22:50

Perhaps, but you will need to experiment.

Depending on the concentration of acetic acid you desire, and how much of it you would like to concentrate, you could just forcibly dry it out with a drying agent e.g., anhydrous copper sulphate - but this becomes less practical for larger volumes.

Mateo_swe - 11-11-2024 at 07:10

Why not buy some glacial acetic acid, its on ebay and other places not hard to find.

UndermineBriarEverglade - 11-11-2024 at 10:32

Mateo, I don't want to attract attention by buying precursor chemicals online. Everything must be OTC.

I added MgSO₄ to a IPA/vinegar concentrate solution and was unable to see any layer separation. Added salt, no change.

I tried using MgSO₄ as a drying agent, since it's cheap and reusable. I added 500g of 45% vinegar concentrate to 305g of MgSO₄·.48H₂O (only partially dried) and got 157g of product. Density is 1.0563 g/mL, suggesting either glacial or 45%. It bleaches pH paper (but so did the 45%). I really ought to get some indicator so that I can titrate. It has reached -21C without freezing, so it's less than 77% by weight. My freezer doesn't actually go down to the eutectic point so this whole enterprise might have been a mistake.

Notes:

Add the MgSO₄ to the vinegar, not the other way around. It formed a strong cake and I broke a stirring rod.
Chill the vinegar beforehand. Forming this much MgSO₄·7H₂O is exothermic. It got up to 70C or so. Powerful smell.
Filtering this much salt was frustrating, hence the 70% yield. I think it'd be better to add just enough salt to comfortably get past the eutectic point, and then get the rest of the way via fractional freezing.

Phase-diagram-for-acetic-acid-and-water-5-621640101.png - 29kB

[Edited on 2024-11-11 by UndermineBriarEverglade]

Keras - 12-11-2024 at 08:57

I don’t think you can separate acetic acid from water.
The best way is to neutralise your acetic acid with sodium bicarbonate/carbonate/hydroxide, boil until you get the salt and then dissolve the salt in a stoichiometric amount of concentrated sulphuric acid. That should give you sodium sulphate and pure acetic acid, assuming you have access to 98% sulphuric acid.

averageaussie - 13-11-2024 at 15:35

Quote: Originally posted by UndermineBriarEverglade

Mateo, I don't want to attract attention by buying precursor chemicals online. Everything must be OTC.

I added MgSO₄ to a IPA/vinegar concentrate solution and was unable to see any layer separation. Added salt, no change.

I tried using MgSO₄ as a drying agent, since it's cheap and reusable. I added 500g of 45% vinegar concentrate to 305g of MgSO₄·.48H₂O (only partially dried) and got 157g of product. Density is 1.0563 g/mL, suggesting either glacial or 45%. It bleaches pH paper (but so did the 45%). I really ought to get some indicator so that I can titrate. It has reached -21C without freezing, so it's less than 77% by weight. My freezer doesn't actually go down to the eutectic point so this whole enterprise might have been a mistake.

Notes:

Add the MgSO₄ to the vinegar, not the other way around. It formed a strong cake and I broke a stirring rod.
Chill the vinegar beforehand. Forming this much MgSO₄·7H₂O is exothermic. It got up to 70C or so. Powerful smell.
Filtering this much salt was frustrating, hence the 70% yield. I think it'd be better to add just enough salt to comfortably get past the eutectic point, and then get the rest of the way via fractional freezing.

[Edited on 2024-11-11 by UndermineBriarEverglade]

What i'm gathering from this is that your salting out method failed (density of either 45% or glacial + melting point being low enough to be less than 77% so its 45%)

Your best bet is definitely either gonna be synthesis (as already stated) or distillation, which you said you couldn't do.
so my next best guess would be either to try a different salt (calcium chloride might work?) or using another solvent to form a different azeotrope with either the water or acetic acid. I have no experience with this however, and cant really help you because azeotrope tables confuse me.
you also might be able to find something that reacts with water but not acetic acid.
anhydrous copper sulfate might also work, can be dried by heating. (http://www.sciencemadness.org/talk/viewthread.php?tid=17976)

UndermineBriarEverglade - 14-11-2024 at 10:29

More of an attempt to dry it than a salting-out really. What I don't understand is that the MgSO₄ released heat on addition, so it was doing something. I might repeat the experiment at smaller scale with a larger excess of MgSO₄ and more careful addition.

I attempted the sodium acetate method before (with vinegar, charring from sugar) and found that more than the stoichiometric quantity of acid was required to wet all the sodium acetate. That method is probably easier with distillation. I only have 93% sulfuric. Is the idea that the sodium sulfate forms the heptahydrate and absorbs any water that was present?

Texium - 14-11-2024 at 12:39

Quote: Originally posted by UndermineBriarEverglade

Everything must be OTC.

A friendly word of advice: if you go strictly by this policy, you likely will not get very far in home chemistry. Especially not if you don't even have a distillation apparatus. From what I've observed, the self-imposed limitation of "everything OTC" usually leads to struggling for months or years to purify a handful of simple chemicals, then realizing you still can't do that much, getting bored, and quitting. Building up a lab from scratch using nothing but kitchenware and OTC chemicals seems like a cool idea on paper, but gets exhausting in practice. Just buy GAA online. It's not a suspicious chemical and isn't on any lists. You'll save money and time and have a superior quality reagent.

Precipitates - 14-11-2024 at 20:34

I would second that.

You can either have good laboratory equipment to prepare these reagents i.e., distillation apparatus, which is kind of basic from an organic chemistry point of view, or just buy these chemicals directly.

It's very difficult for everything to be OTC, and not have the proper apparatus to purify or produce new chemicals - a double hinderance. I've been there before (as likely many of us have), but quickly (although probably not quickly enough), got tired of it.

I still think you could dry the acetic acid with a drying agent - but multiple runs would be required, adding a bit, recovering the purer acetic acid, adding more agent, recovering again etc. You can measure the density for a rough indication of purity, but I wouldn't be surprised if 10+ separate additions of drying agent would be required.

Is it worth it - probably not. It may just turn out to be a rather boring fruitless endeavour.

clearly_not_atara - 14-11-2024 at 21:02

I think you can obtain anhydrous acetic acid by half-neutralizing the acetic acid and evaporating to form solid sodium diacetate (soluble to about 119g / 100 ml), which can be dried and will release acetic acid at high temperatures. This doesn't require anything but dilute acetic acid and sodium bicarbonate or carbonate. You do need to figure out the stoichiometry, but I don't think it has to be exact.

For 100 mL of 45% acetic acid density is 1.05 so I would estimate there are about 47 grams of acetic acid so about 28 grams of sodium bicarbonate should give about 48 grams of precipitate yielding about 20 grams of anhydrous acid. The theoretical yield is only 50%, but you could reuse the sodium acetate to make more diacetate.

RU_KLO - 15-11-2024 at 04:48

Quote: Originally posted by clearly_not_atara

I think you can obtain anhydrous acetic acid by half-neutralizing the acetic acid and evaporating to form solid sodium diacetate (soluble to about 119g / 100 ml), which can be dried and will release acetic acid at high temperatures. This doesn't require anything but dilute acetic acid and sodium bicarbonate or carbonate. You do need to figure out the stoichiometry, but I don't think it has to be exact.

For 100 mL of 45% acetic acid density is 1.05 so I would estimate there are about 47 grams of acetic acid so about 28 grams of sodium bicarbonate should give about 48 grams of precipitate yielding about 20 grams of anhydrous acid. The theoretical yield is only 50%, but you could reuse the sodium acetate to make more diacetate.

Without knowing starting acetic acid concentration, how could you, for example by titration, know the end point to get the diacetate?

(My undestanding is that neutralization of acetic acid with carbonate to ph 7 aprox, will shield sodium acetate)

Maybe, perform a titration of acetic acid with carbonate, then use half of carbonate fot the main stock of acetic acid?

[Edited on 15-11-2024 by RU_KLO]

UndermineBriarEverglade - 15-11-2024 at 09:31

Texium & others - I appreciate the advice about OTC. I know it's difficult, but I am trying to avoid prison for my main interest of explosives. It will be safer if nobody knows I'm interested in chemistry at all. Unfortunately my OTC glassware source doesn't have distillation columns. I'm only pursuing acetic and nitric acids "on the side" and don't want to sink too much effort into them.

Sodium diacetate is interesting, never heard of it. I do think this route would be lossy. It would also (as usual) benefit from a distillation apparatus, because I imagine the acetic acid would quickly evaporate.

averageaussie - 17-11-2024 at 13:44

Quote: Originally posted by UndermineBriarEverglade

More of an attempt to dry it than a salting-out really. What I don't understand is that the MgSO₄ released heat on addition, so it was doing something. I might repeat the experiment at smaller scale with a larger excess of MgSO₄ and more careful addition.

I attempted the sodium acetate method before (with vinegar, charring from sugar) and found that more than the stoichiometric quantity of acid was required to wet all the sodium acetate. That method is probably easier with distillation. I only have 93% sulfuric. Is the idea that the sodium sulfate forms the heptahydrate and absorbs any water that was present?

Magnesium sulfate hydrating is exothermic, so it was indeed absorbing the water. if you have access to a pot, a vacuum pump and some silicone you can make a basic vacuum drier, just add anhydrous magnesium sulfate or calcium chloride in the same container as your thing you want dried, put on the lid and start the vacuum. there are plenty of tutorials online for this.

if you only have 93% sulfuric, could you not dilute it down? I cant guarantee that it will work diluted however, but it might just be worth a shot.

magnesium sulfate and similarly sodium sulfate do chemically bond to water, forming the heptahydrate. worth noting is that magnesium and sodium sulfate are both insoluble in ethanol, if that helps.

kmno4 - 22-11-2024 at 03:52

You cannot "extract" water from AcOH/H2O mixtures with aid of MgSO4, because it forms stable solvates with AcOH.
In another words, AcOH will remove water from MgSO4xH2O.
BTW recently I obtained ~80% AcOH from some store and - if I have time - I am going to try extractive distillation with EtAc using D-S trap and some additional modification. I just prefer to have 99% AcOH than 80% one.

averageaussie - 24-11-2024 at 14:54

hm, thats annoying.
So concentrated acetic acid is a dehydrating agent then, no?
is there a numerical measure of how strong of a dehydrating agent it is, or is it just "chemical x wants water more than chemical y"?
if there is a numerical measure, it shouldn't be particularly difficult to find a chemical with bigger number ™ and use that to dehydrate it.
or, better yet, find a chemical that actually destroys the water but doesn't attack the acetic acid (something like P2O5 or maybe sodium, thats a big maybe though and sodium probably directly reacts with the acetic acid, but you get the idea)
also, what did you buy 80% acetic acid as?
as for concentrating acetic acid for Everglade, so far the best solution still seems to be distillation, and I do have to agree with Texiums advice - working only OTC is extremely limiting and boring.

averageaussie - 24-11-2024 at 15:01

just realised that (for sodium being used to dehydrate) the newly formed NaOH will react with the acetic acid, neutralising it :/ whoops

kmno4 - 4-12-2024 at 07:01

Quote: Originally posted by kmno4

... to try extractive distillation with EtAc using D-S trap and some additional modification. I just prefer to have 99% AcOH than 80% one.

Ha, it works. Unfortunately, Vigreux column was needed (40 cm) becuse separation of ethyl acetate-water azeotrope was highly imperfect without fractionation. Practically anhydrous AcOH was obtained, "freezing" completely below 10 C. About 30 cm3 of water (saturated with NaCl) was separated from ~140 g of 80% AcOH, so everything fits. The azeotrope contains ~10% of water, so dehydrating of larger amoun of AcOH takes some time.
Nice method of removing of smaller amounts of water, especially from highly concentrated (>90%) AcOH.

Fery - 6-12-2024 at 08:04

Buy distillation apparatus and do what kmno4 suggested you.

https://www.chemeurope.com/en/encyclopedia/Azeotrope.html

Quote:

Use of azeotropes to separate zeotropic mixtures
Sometimes azeotropes are useful in separating zeotropic mixtures. An example is acetic acid and water, which do not form an azeotrope. Despite this it is very difficult to separate pure acetic acid (boiling point: 118.1°C) from a solution of acetic acid and water by distillation alone. As progressive distillations produce solutions with less and less water, each further distillation becomes less effective at removing the remaining water. Distilling the solution to dry acetic acid is therefore economically impractical. But ethyl acetate forms an azeotrope with water that boils at 70.4°C. By adding ethyl acetate as an entrainer, it is possible to distill away the azeotrope and leave nearly pure acetic acid as the residue.

kmno4 - 6-12-2024 at 15:12

By the way, instead of EtAc, one can use rather inexpensive ether: MTBE.
It forms also one water azeotrope in the system H2O-AcOH-MTBE, with b.p. ~55 C. It is then easier to separate, but it contains only ~4 % of H2O. Other substances also can be used, but they are much more expensive than mentioned EtAc or MTBE.

[Edited on 6-12-2024 by kmno4]

Mateo_swe - 17-12-2024 at 06:19

If you are concerned with "flying under the radar" maybe you can build a distilling apparatus from metal.
Either DIY one together yourself or ask a skilled metalworker to weld together a distiller.
Its common with those who like to make moonshine to make these things.
A simple retort might work but a little more advanced one with cooling water would be better.
With a distilling apparatus you could try the suggestions above and be able to make glacial acetic acid.

A easier route would be to pay someone to buy some glacial AcOH for you.
But i guess you need someone you can trust completely if you live in a place where intrest in hobby chemistry and smallscale energetic materials can send you to jail.