Sciencemadness Discussion Board

Unkown product, Glacial Acetic acid distillation

walruslover69 - 21-12-2017 at 20:49

I recently acidified sodium acetate with concentrated sulfuric acid and distilled to produce glacial acetic acid. The product smelled very strongly of acetic acid, I put it in the freezer (~ -5 degrees) to get a rough estimate of purity ( if it froze I would know that it is at least 85% acetic acid). It did not freeze so i added anhydrous magnesium sulfate in an attempt to dry it out. Upon adding the MgSO4 it appeared to react quite vigorously with the acetic acid solution. The MgSO4 bubbled with the same intensity as if i had added a base or metal oxide. This is extremely weird and I have no idea why this is happening. I double checked to make sure I had not mislabeled my MgSO4 and added something different, so i added a small amount of regular Walmart Epsom salt, and it reacted in the same way although not as strongly. Do anyone have any thoughts about what could be going on? or possible impurities that would cause this. Magnesium sulfate is usually extremely stable for the most part only ever reacting to form a precipitate. I am very perplexed. I would love to hear people's suggestions.

JJay - 21-12-2017 at 21:25

Fascinating. Perhaps the magnesium sulfate dried out the glacial acetic acid, causing it to freeze in the low temperature, giving off its heat of crystallization and... I am not sure what happened, but that sounds like a fun thing to try.

walruslover69 - 21-12-2017 at 21:33

the acid isn't freezing, the MgSO4 reactions until it all dissolves. I just got a laboratory distillation apparatus so i am working my way through synthesizing and purifying a whole list of common solvents and reagents.

Dr.Bob - 22-12-2017 at 05:13

You might be making Mg(OAc)2 and H2SO4 from the reaction of MgSO4 and HOAc. The H2So4 would then react with the water left violently. Just does not sound like a great way to dry it, unless you can redistill it afterwards to remove all of the salts and mess. The best way to get the acetic acid pure is to redistill from more acid again. With a simple distillation (no fractionating column or other complex techniques), you sometimes need multiple runs to get to the desired purity. Commercial systems (like distilleries) use long columns and process controls to achieve their purities.

unionised - 22-12-2017 at 05:35

Dry MgSO4 reacts quite vigorously with water.
Bubbles of gas trapped in the solid may look like gas being evolved.

The reaction to produce H2SO4 and Mg acetate isn't plausible- look at the acid strengths.

walruslover69 - 22-12-2017 at 08:26

I only took a small sample of maybe 5 ml to dry with magnesium sulfate just to get an idea of how much water was in it. I plan on re distilling the bulk of it in order to purify it. trapped bubbles being released as it dissolves is probably the most likely cause, I will drop some of the dry magnesium sulfate into pure water to see how it compares later.

unionised - 22-12-2017 at 08:30

Incidentally, I'm not sure MgSO4 is a good choice of drying agent; it may be soluble in acetic acid.

walruslover69 - 22-12-2017 at 08:53

That is probably part of the problem, now i feel quite stupid. I will probably try calcium or copper sulfate later this afternoon. Calcium sulfate is barley soluble in water so I can't imagine it would be soluble in acetic acid. or I might just give up and re distill anyways.

NEMO-Chemistry - 22-12-2017 at 08:59

Calcium Sulphate is plaster of paris... Although i assumed it would react with acetic acid??

walruslover69 - 22-12-2017 at 09:05

Calcium sulfate is a fairly potent desiccant. It is the conjugate of a strong acid (H2SO4) and a strong base (Ca(OH)2) so it is neutral and doesn't react.

NEMO-Chemistry - 22-12-2017 at 11:34

Quote: Originally posted by walruslover69  
Calcium sulfate is a fairly potent desiccant. It is the conjugate of a strong acid (H2SO4) and a strong base (Ca(OH)2) so it is neutral and doesn't react.
It will with water though... Although i think HCl dissolves it, i cant remember i havnt used it for a long time as a desiccant.

aga - 22-12-2017 at 13:33

Quote: Originally posted by unionised  
The reaction to produce H2SO4 and Mg acetate isn't plausible- look at the acid strengths.

Could you please explain that further, pretty please ?

Acetic acid (as per wiki) has a couple of 'strengths' , namely a log p of -0.322 and pka of 4.76

What does that actually mean ?

unionised - 23-12-2017 at 05:36

Log p isn't an acidity it's a partition coefficient.

The pKa (4.76) is effectively the pH at which half of it is ionised.
The pKa for sulphuric acid is about -3 (it's a strong enough acid that it's difficult to measure properly)


Dr.Bob - 23-12-2017 at 05:59

The reaction would clearly be in equilibrium, and shifted away from the direction I pointed, but the reaction of water with the acid could shift it back some, as it is nearly irreversible. My point is that when you mix inorganic salts in solution, you rarely get them back pure, they tent to form mixtures of many salts that make purifying your material tough, thus the need for distillation or other methods of purification. Drying agents work best when they are 1) quite insoluble in the solvent used, 2) non reactive with the solvent and other solvates, and 3) there are no other ions or reactives in solution to react with them.

Thus sodium sulfate will work well in ethyl acetate to dry a hydrocarbon product in solution, but drying acetic acid with potassium carbonate is not so good. I don't know that MgSO4 will react with acetic acid, but it is theoretically an issue.

Likewise drying HBr with sulfuric acid can reduce the HBr to bromine, but I think it works to dry HCl with it. Many cases are not cut and dry. For instance, drying ethanol is very tough, as it either dissolves or complexes with many drying agents. Only a few work, like 3A sieves and Alumina.

unionised - 23-12-2017 at 07:34

Quote: Originally posted by Dr.Bob  
The reaction would clearly be in equilibrium, and shifted away from the direction I pointed, but the reaction of water with the acid could shift it back some, as it is nearly irreversible. My point is that when you mix inorganic salts in solution, you rarely get them back pure, they tent to form mixtures of many salts that make purifying your material tough, thus the need for distillation or other methods of purification. Drying agents work best when they are 1) quite insoluble in the solvent used, 2) non reactive with the solvent and other solvates, and 3) there are no other ions or reactives in solution to react with them.

Thus sodium sulfate will work well in ethyl acetate to dry a hydrocarbon product in solution, but drying acetic acid with potassium carbonate is not so good. I don't know that MgSO4 will react with acetic acid, but it is theoretically an issue.

Likewise drying HBr with sulfuric acid can reduce the HBr to bromine, but I think it works to dry HCl with it. Many cases are not cut and dry. For instance, drying ethanol is very tough, as it either dissolves or complexes with many drying agents. Only a few work, like 3A sieves and Alumina.

"The reaction would clearly be in equilibrium, and shifted away from the direction I pointed"
Yes, by about 6 or 7 orders of magnitude. Practically speaking, that means it just doesn't happen.
"but the reaction of water with the acid could shift it back some, as it is nearly irreversible."
Not really, the reaction of conc H2SO4 with water can be reversed by simple distillation.

"My point is that when you mix inorganic salts in solution, you rarely get them back pure, "
True, but irrelevant since there's only 1 inorganic salt.

" thus the need for distillation or other methods of purification."
That's about the only thing you said that's true and relevant.

"Thus sodium sulfate will work well in ethyl acetate to dry a hydrocarbon product in solution,"
True, but distillation might be better + easier.
Also, since water doesn't mix with hydrocarbons, drying then isn't usually a problem unless you want them absolutely anhydrous, and sodium sulphate isn't good enough for that.


" I don't know that MgSO4 will react with acetic acid, but it is theoretically an issue. "
If you don't know then why post it as a possibility, without checking.
And,at best, it's hypothetically possible. You should find out what "a theory" means in science.

"Likewise drying HBr with sulfuric acid can reduce the HBr to bromine,"
No, it oxidises it.

"drying ethanol is very tough, as it either dissolves or complexes with many drying agents. Only a few work, like 3A sieves and Alumina. " ironically, magnesium sulphate works reasonably well- though you still need to distil the product.
It's possible to dehydrate ethanol by distillation alone without using a drying agent at all, but its more suitable as an industrial process than a lab one.

walruslover69 - 18-4-2018 at 07:02

I recently preformed the synthesis of glacial acetic acid again. I started with store bought white vinegar and neutralized it with sodium carbonate. I then boiled the solution down to the sodium sodium acetate powder. I am aware now that this is not a good method because of sugars that are added to the vinegar. I then added sulfuric acid to the sodium acetate and distilled over the acetic acid. For the first half of the distillation i was obtaining what i believe to be relatively pure glacial acetic acid. Towards the end of the distillation, white fumes started to be produced. I turned off the heating several minutes later. To my complete surprise there was a thin deposit of sulfur on the inside of my condenser.

Does anyone know what might be going on. I am familiar with sulfuric acid being reduced to sulfur dioxide by sugar or carbon but never to elemental sulfur. My best guess would be that the sugar is reducing the sulfuric acid to sulfur dioxide and then there is somehow hydrogen sulfide produced that reduces the sulfur dioxide to elemental sulfur.

LearnedAmateur - 18-4-2018 at 11:45

Sugars shouldn’t be able to reduce sulphuric acid at all, I wouldn’t have thought. I’ve tried the dehydration of sucrose to carbon and never got even a funny smell, let alone gas evolution or yellow solids.. I doubt there’d be any ‘special’ (other than sucrose, glucose, fructose, dextrose, etc.) sugars either, if at all; I was under the impression that white vinegar was literally just diluted GAA? Are you sure it’s not distilled malt vinegar, I.E does it state for food or cleaning use?

[Edited on 18-4-2018 by LearnedAmateur]

aga - 18-4-2018 at 11:56

Both the production of sodium acetate then GAA with conc sulphuric have always been easy and straightforward in my limited experience.

It is pretty much impossible to even guess without photos, especially the sulphur thing.

Edit:

Well, i could guess, and most likely be correct, but those dark, bleak days are behind me.

[Edited on 18-4-2018 by aga]

walruslover69 - 18-4-2018 at 12:38

It was walmart brand "distilled white vinegar". When i was baking the sodium acetate down it turned a very slightly golden brown and smelled like it was caramelizing. I have read about some other people on the forum noticing that there is sugar in it. I assumed when i first attempted it that it was just diluted acetic acid as well.

doing it straight from pure sodium acetate is a fairly straightforward process. Its the added sugar in the vinegar that seems to cause problems. I didn't take a photo before I cleaned it, but it looked exactly like this. https://lab.whitequark.org/images/sulfur-chloride-synth/try-...

[Edited on 18-4-2018 by walruslover69]

Ubya - 18-4-2018 at 14:37

sometimes vinegar has sulfites(additive E221 to E228), and in acidic solution they produce sulfur dioxide.
what produced elemental sulfur i have no idea

XeonTheMGPony - 18-4-2018 at 16:10

do concentrate acetic do not re distill from h2so4 unless you enjoy carbon! re distill over anhydrous sodium acetate.

Be for you run the acetic acid generating phase heat the living hell out of every thing to thoroughly dry it!

I've run more then one synth and still perfecting it, but that is what I had found in literature.

cheaper and easier way is get acetone free nail polish remover, make sure it is the ethyl acetate type.

Distill this to get just the ethyl acetate. Next add sodium hydroxide and reflux for a hour or so, distill off the now free ethanol (This will be very dry and pure! Handy way to get the best bang for buck in time and money spent 2 reagents for the price of one!).

Check the PH of your remaining solution, if still strongly basic add a bit more ethyl acetate and reflux, rinse and repeat till no longer basic, distill off fresh batch of ethanol.

Remove the now saturated sodium acetate solution and dry it, once dry fluff up crystals and put in micro wave to really dry it.

Dissolve into clean dry hot methanol use little methanol as possible to dissolve it, hot filter any thing not dissolved, allow to slowly cool.

Filter out your nice pure sodium acetate crystals and rinse with some freezing cold methanol.

Re-microwave dry after air drying thoroughly.

set aside in air tight container. Now distill some sulfuric acid to concentrate to 98% (And if very serious make some sulfur trioxide to get 100%)

Oven bake your glass wear at 150 for an hour, once cool enough to handle assemble and charge apparatus!

Profit?

[Edited on 19-4-2018 by XeonTheMGPony]

[Edited on 19-4-2018 by XeonTheMGPony]

aga - 18-4-2018 at 23:28

Even with brown or red wine vinegar you can make white, mostly pure, sodium acetate crystals simply by recrystallising them once.

The idea of making ethanol at the same time sounds much sexier.

LearnedAmateur - 19-4-2018 at 01:22

Yeah I was going to suggest that, growing NaOAc crystals to crush, dry, and use seems like the best option from impure acetic acid. Oh yeah, and if you’ve got ethanol from the mix too and don’t need it as a solvent then you can always oxidise it for more acetic acid.

woelen - 19-4-2018 at 01:45

Sugars certainly are capable of reducing sulphuric acid. Most of the reaction is dehydration, but there is a side reaction, especially when the acid gets hot. In this side reaction you get SO2. Getting all the way down to S, or even H2S, however, would surprise me.

LearnedAmateur - 19-4-2018 at 03:38

Well I never knew that, what are the products of oxidation? I assume one or more of the hydroxyl groups get oxidised to carboxylic acids or does it go all the way to CO2 + water?

woelen - 19-4-2018 at 04:33

The result of the oxidation is a real mess. With such complex oxidations you will have a (useless) mix of products. In some cases, the chain of carbon atoms may break apart, at other places a carbonyl group may appear. At the end-points of chains carboxylic acids may be formed and there almost certainly will be formation of some CO2 as well.

LearnedAmateur - 19-4-2018 at 06:17

I could guess that it would be quite varied just by looking at glucose for instance, all those hydroxyl groups to react, but I’ve never been one to study saccharides in the first place though, they just don’t really seem that interesting and are a little to complex in their behaviour for my liking. Thanks for the info anyway