Sciencemadness Discussion Board - Acetate Salts From Vinegar For Synthesis of Acetic Acid

Acetate Salts From Vinegar For Synthesis of Acetic Acid

Lytzu - 9-1-2020 at 14:33

I would like to prepare glacial or near glacial acetic acid for use as a reagent. I think I've read all the relevant threads at least half a dozen times and am starting to feel confident that it is something I can accomplish. However, none of these threads seem appropriate for this post as, for now, I will be focusing specifically on acetate salt production. If this would be better merged to a previous thread, the most recent may be applicable.

Anyway, from what I've read the simplest way for the home chemist to produce acetic acid is to "salt out" an acetate from vinegar. This can be achieved by the addition of sodium carbonate, sodium bicarbonate, calcium carbonate, etc, which reacts with the acetic acid in the vinegar to form an acetate salt. Excess water can then be boiled off. A second reaction of the acetate salt and sulfuric acid turns the salt back into acetic acid.

I have a very small amount (<=10g) of sodium carbonate on hand from an old chemistry kit I got as a kid and never really used. Eventually I would like to try this with other bases but for now I'll work with what I have.

My plan is to follow the chemical equation 2CH₃COOH (acetic acid) + Na₂CO₃ (sodium carbonate) = 2NaCH₃COO (sodium acetate) + H₂O + CO₂
at 3/100ths of a mole. From the molar masses I calculated this to be 72g 5% vinegar + 8.58g sodium carbonate will yield about 4.92g sodium acetate. The only reason I am doing the experiment on such a small scale is because of the limiting factor of the amount of sodium carbonate I have. After I try this I will obtain either more sodium carbonate or sodium bicarbonate.

The scale I bought is coming in early next week and in the meantime I would appreciate the help determining if this is a viable plan. I will be updating this thread as I progress beyond this one experiment, which is part of why I thought I'd make my own post. Further experiments of this subject will likely include producing acetates from 30% vinegar if I can find it (the depot lied, said it was there and it wasn't) and 85-90% photography indicator stop bath solution. As I mentioned before I would also like to try with sodium bicarbonate, and perhaps calcium carbonate and others. For now I am focusing on this acetate salt step of production but will eventually progress to reverting the acetate back to acetic acid via sulfuric acid.

Thanks in advance. If anything needs correcting or I missed a step or something PLEASE let me know! I am excited for my first chemistry experiment outside of school or work.

[Edited on 10-1-2020 by Lytzu]

vibbzlab - 9-1-2020 at 17:34

This is actually a viable experiment.I had tried this experiment once by just crystallising the salt from vinegar. Your yield would not be that great amounts though as vinegar has a low concentration of acetic acid in it. But yes definetly you can get glacial acetic acid by adding concentrated Sulfuric into the salt. I would say but glacial acetic acid since that's not that difficult to get . I wouldn't use up my concentrated Sulfuric acid for acetic acid. But yes if you have no other way or if you want to try it out,then yeah,go ahead. All the best.

DraconicAcid - 9-1-2020 at 18:23

You have a very small amount of sodium carbonate (Na2CO3)? You can buy it dirt cheap as washing soda (although it's probably either the heptahydrate or decahydrate). Or you can buy baking soda (sodium hydrogen carbonate- NaHCO3).

Your sodium carbonate may also be hydrated, and the sodium acetate that you get will probably be the trihydrate.

vibbzlab - 9-1-2020 at 18:34

I think by strong heating them u can get anhydrous form right? Correct me if I am wrong

Lytzu - 9-1-2020 at 18:49

Na2CO3, ah, did I get that formula wrong? Good catch thanks Draconic, fixed.

I will be picking up more washing soda and/or baking soda soon, I just figured I'd use up the tiny bit I have as a test.

I believe the sodium carbonate I have is at least partially hydrated. It may not have come that way but after sitting in a basement for years most of it is in a large relatively hard clump.

If the sodium acetate I obtain is trihydrated how would that affect things? Will the acetic acid created from that be less concentrated because of the water trapped in the sodium acetate crystal? Does this change the sodium acetate itself in any relevant way?

vibbzlab, I had heard that other acids might be able to be used but I'm not certain. So far I've mostly ignored this step in favor of fully understanding the first part of the process (the acetate formation).

XeonTheMGPony - 10-1-2020 at 04:00

to purify: Recrystallize in methanol.

all so a good way to get sodium acetate and ethanol is to get acetone free nail polish remover (Check list of ingredients to be sure) distill out the ethyl acetate then reflux with sodium hydroxide.

to dry the sodium acetate, you can micro wave it till it no longer heats up, be care full to check often, as once all the water is boiled off it will be very hard on the micro wave! Or you fuse it (Heat it to melting point)

[Edited on 10-1-2020 by XeonTheMGPony]

Ubya - 10-1-2020 at 05:10

i did this experiment a few years ago.
i used "vinegar of alcohol", vinegar used for cleaning, not made from grapes so it shouldn't have sugars (white vinegar has sugars). i neutralized 2 liters of vinegar (6% acetic acid if i remember correctly) with sodium bicarbonate. i then boiled down the solution until only a powder remain. the trihydrate melts at 58°C, you want notice when all the water is gone since you'll still have a liquid. i heated until i got a syrup, and then i heated it more until i got the anhydrous form.
i added sulphuric acid and performed a distillation. i got really concentrated acetic acid (by the smell) but it is not glacial as it doesn't solidify in the fridge, i haven't titrated it yet as i don't have a burette.

betamethyl - 10-1-2020 at 16:23

I did this last year too, using a slightly different method. I have my notes written somewhere, but here is a rough outline:

I wanted to produce glacial acetic acid from vinegar, so i planned to: react the vinegar with sodium bicarbonate to produce sodium acetate in my first step; and, to distill with 98% sulfuric acid in the second step. I went to coles and bought around 12 litres of home-brand vinegar.

I didn't know what concentration of acetic acid was in my vinegar. I didn't really have any pH paper, so i would have to estimate the quantity of sodium bicarbonate to add. But if i added too much sodium bicarbonate then i would be left with sodium bicarbonate/carbonate in my acetate solution. This would contaminate my Na acetate with Na bicarbonate, and would create problems during the distillation step with 98% sulfuric acid.

To prevent this, i bought 5 Kg of "Garden lime" from bunnings. This stuff was only around 20% Calcium carbonate/hydroxide as far as i can remember; the rest was mainly sand.

My plan was to react the vinegar with excess calcium carbonate to form calcium acetate, then to react this calcium acetate with sodium carbonate and filter off the calcium carbonate byproduct to get a solution of sodium acetate. Because i couldn't test pH and i didn't know how much acetic acid was in my vinegar, i planned to do this extra step so that i could weigh the calcium acetate product (with minor calcium carbonate impurity due to its low solubility in water) and determine the concentration of acetate in my vinegar more accurately. By producing the calcium salt from the vinegar first, the possibility of having any bicarbonate or carbonate in my RBF when i go to add 98% sulfuric acid is removed (as i had accidentally done this in an earlier experiment). By converting the calcium acetate to sodium acetate before distilling with sulfuric acid, a cleaner, simpler distillation of acetic acid is afforded without calcium sulfate byproduct which can coat the calcium acetate solid in the flask when added to sulfuric acid. When this happens, it kind of ruins the reaction because the mixture is just too viscous to stir magnetically, so you have to pull things apart and poke around with a glass rod while the sulfuric acid is hot and glacial acetic acid vapour is coming out of the flask which will make your eyes water for sure.
In the end, i thought that calcium acetate would be okay to distill with sulfuric acid to form GAA. I found out the hard way that this ended up causing problems.

Calcium acetate:
1. 12 litres of vinegar (approx 5%) was poured into a 40 litre plastic bucket.

2. "Garden lime" was spooned in to the vinegar with a ladle to prevent splashing. The mixture stirred manually. The production of carbon dioxide was mild - probably due to all the sand taking up so much room and 'diluting' things.

3. After leaving the 5 Kg lime in the vinegar overnight, a pinch of bicarb soda was added to the mixture; the white powder bubbled as it dissolved. This indicated that unreacted vinegar was present, so the mixture was poured into a 19L stainless steel metal pot (now in the garbage bin), combined with as much crushed chalk and calcium carbonate-containing items i could find, and boiled over a campfire stove. The mixture was periodically stirred to avoid bumping, but the sand acted like a heat sink and just prevented the liquid from boiling vigorously.

4. After leaving the metal pot on top of the campfire stove all night, a pinch of bicarb soda was added to the liquid the next morning and the mixture did not bubble.

5. Filtering such a large mass of solid would require a series of filtrations and the better half of a day. Instead, i siphoned the supernatant liquid off from the sand and filtered once through kitchen paper towel. I spooned the sand into some old large socks and squeezed the liquid out - this liquid only needed a simple gravity filtration through kitchen paper to clear up.
If the liquids were at all cloudy at this point, it didn't much matter at this stage.

6. The large volume of liquid needed to be removed of all water. Because of the large volumes, i found that heating all 10-15 litres of liquid in one pot was very inefficient as my heater struggled to bring everything to a boil. Instead, i found that using multiple 1-2L pots and using all 4 heating elements on my kitchen stove boiled the water off fastest. The calcium acetate solution (which was brown/yellow) was periodically ladled into these pots until almost all the water was removed. When this occurred, the calcium acetate paste was spooned all into one pot and some water was used to wash the other pots and transfer remaining calcium acetate into the main pot. This main pot was heated until a large amount of solid started precipitating. Soon after, all the water had been removed and the steam adopted a vinegar-like odor. The mixture became homogenous and it appeared that the calcium acetate had melted. I just took the pot off the heat and covered it with a piece of slate to prevent moisture getting in as it cooled. If i was to repeat this, i would recommend pouring the molten calcium acetate onto a metal tray and allowing to cool (preventing moisture), because getting the product out of that metal pot required a hammer and chisel. The yield was around a kilogram i think? maybe 600-800g? i know that i ended up with a yield of acetic acid which worked out to a concentration of 4% acetic acid in my original vinegar.
Make sure you store calcium acetate in a ziplock bag or a screw-lid jar and use soon after manufacture because it was very hygroscopic. I believe sodium acetate behaves much the same.

I also have made sodium acetate. In my first experiment, i combined sodium bicarbonate and vinegar in a pot and boiled to remove the water. It made horrible smelling fumes and went brown. This commonly happened to me over more experiments with sodium bicarbonate, and i never really knew why. I could never remove the organics easily, or prevent degradation. This is why i used calcium carbonate in excess for my experiment with 12 litres of vinegar - i suspected that i was adding too much sodium bicarbonate to the vinegar and it was leading to problems when i heated it to remove the water in later steps.

For distilling, i would recommend one of these 2 methods:
1: Combining sodium acetate (anhydrous) and concentrated sulfuric acid and distilling.
2: If your acetic acid doesn't need to be glacial, you can make lower-concentration solutions by combining an aqueous solution of calcium acetate with dilute sulfuric acid and filtering off the calcium sulfate precipitate.

I have also read things on SciMad about combining oxalic acid with calcium acetate solutions and filtering off the calcium oxalate which is very insoluble, but i have no experience with this.

annaandherdad - 10-1-2020 at 18:14

I bought some washing soda as a cheap source of sodium carbonate. When I dissolve it in water it makes some foam which my reagent sodium carbonate does not, and the solution is not as clear. Therefore I have suspected that my washing soda contains some contaminants.

On the other hand, sodium bicarbonate can be bought in the grocery store and it is fit for human consumption. That does not prove it is pure, but when I bake it in the oven (300 degrees F for an hour) I get a yield that by weight is within 1% of what I would expect for conversion of pure sodium bicarbonate into pure sodium carbonate. (This is for Arm and Hammer brand baking soda, I'm in the US.) Also, the resulting product dissolves in water like my reagent sodium carbonate, making a clear solution without foam.

To summarize, I have suspicions that commercial washing soda is not a good source for high quality sodium carbonate, while baking soda, heated to convert to sodium carbonate, may be.

Cou - 13-1-2020 at 09:38

Nile red made sodium acetate from 5% food vinegar.

symboom - 13-1-2020 at 10:01

What about forming calcium acetate from cheap sodium acetate which then could be used to form acetic acid as calcium acetate reacts with sulfuric acid to form insoluble calcium sulfate and acetic acid not sure if this could also work or just make chloroacetic acid but sodium acetate could react with HCl acid to form insoluble salt which NaCl is insoluble in HCl and acetic acid

karlos³ - 13-1-2020 at 10:11

Quote: Originally posted by XeonTheMGPony

to purify: a all so a good way to get sodium acetate and ethanol is to get acetone free nail polish remover (Check list of ingredients to be sure) distill out the ethyl acetate then reflux with sodium hydroxide.

You probably meant, reflux with NaOH, then distill out the formed ethanol?

Sodium Acetate Experimant 1 Complete

Lytzu - 13-1-2020 at 15:18

Alright so I've just finished my first official home chemistry experiment

. As far as I can tell it's been a success.

As planned in my original post my procedure was as as follows:

9 grams of sodium carbonate was added to a 250mL Erlenmeyer flask. Much of it had to be crushed before hand because it had formed a large clump in the jar.
72g of 5% vinegar were added to the Erlenmeyer flask. A white foam was immediately produced. When fizzing completed there was still sodium carbonate powder on the bottom of the flask. I added an additional 54g(!) of vinegar before it all disappeared. Swirling of the flask was used to aid in this process.
Because I do not have a hot plate or Bunsen burner yet an improvised hot plate was created by placing an old pan on the stove top. The flask was set directly on this surface. The burner was set to medium. Bubbles began to rise in the solution very soon after this. It reached a true boil after maybe 20 minutes. By then it was beginning to take on a yellowish tinge. After about an hour to an hour and a half nearly all the liquid had boiled off. When the sodium acetate began to crystalize it happened very quickly, almost all at once.
The crystalized solution was allowed to cool slightly before I attempted to scrape it into a small jar. Due to lack of scoopulas a long screwdriver was used. I was worried about scrapping the glass so I was only able to recover maybe half of the solids. This weighted in at a surprising 4g. It was still very slightly damp so that is probably the majority of why.

Conclusion: Sodium acetate was produced in this process, although likely mixed with impurities.
Questions:

Why did I have to add extra vinegar? Were my calculatons off? Should I allowed the sodium carbonate to stay in the solution and become part of the resulting powder? Which reactant do I want more off? I would assume I'd rather have extra acetic acid in the solution because that is ultimately what I am trying to form, although this probably boils off.

The foam was the result of CO2 produced from the reaction correct? And when boiling the bubbles were just water vapor, just like when boiling water?

There was one spot on the bottom of the flask where more and larger bubbles were produced and in the beginning "sticking" to the bottom. Why is this? I have seen a similar phenomenon when boiling water.

Oh! Almost forgot. It could just be me working in my kitchen (I know, not the best but its the only available space right now) but the sodium acetate had a very slight sent. Like...biscuits? The fumes from my solution had no scent, unlike betamethyl's
Quote:
i combined sodium bicarbonate and vinegar in a pot and boiled to remove the water. It made horrible smelling fumes and went brown.
though they were using bicarbonate and I was using carbonate. Ideas?
Also,
Quote:
the trihydrate melts at 58°C, you want notice when all the water is gone since you'll still have a liquid.
I don't believe this is what happened with mine, unless it went from a syrup to the anhydrous form very quickly. Is Ubya correct that this is what happens?

Additional questions from the thread not yet answered:

If the sodium acetate I obtain is trihydrated how would that affect things? Will the acetic acid created from that be less concentrated because of the water trapped in the sodium acetate crystal? Does this change the sodium acetate itself in any relevant way? Is Ubya correct that heating it results in an anhydrous crystal?
I believe vibbzlab also had a question about if heating the sodium carbonate(?) could create anhydrous sodium carbonate. I believe he was referring to that chemical.

Notes:The sodium acetate crystals formed 2 sort of layers on the bottom of the flask. On top was slightly brownish and still very slightly damp while below this was a harder white substance. I believe this has to do with impurities rising with the water vapor. I think I saw it mentioned in another thread.

I think that is all. My next experiment will likely be similar, although using sodium bicarbonate instead. After that I may move on to trying this with 85-90% acetic acid stop bath solution. Thanks for the help so far.

XeonTheMGPony - 13-1-2020 at 15:46

Quote: Originally posted by karlos³

Quote: Originally posted by XeonTheMGPony

You probably meant, reflux with NaOH, then distill out the formed ethanol?

That's what I said lol

But you must distill out the ethyl acetate first as there is a bunch of other garbage mixed in usually, once that is distilled out then you reflux over Sodium Hydroxide

Once you have re-fluxed for an hour or better, distill off the fresh freed ethanol, you should have a liquid sludge left.

React to slightly acidic with vinegar (As I hope you used excess base) then evaporate to dryness.

Take this dry mass and dissolve in methanol and filter out any that doesn't, wash with a bit of warm methanol to get any out that was retained in the filter paper.

Recover some methanol by distilling to a gell or just evaporate it all off to recover your pure sodium acetate (Sodium carbonate will remain in the filter paper as a solid.

draculic acid69 - 14-1-2020 at 19:36

Quote: Originally posted by XeonTheMGPony

to purify: Recrystallize in methanol.

all so a good way to get sodium acetate and ethanol is to get acetone free nail polish remover (Check list of ingredients to be sure) distill out the ethyl acetate then reflux with sodium hydroxide.

to dry the sodium acetate, you can micro wave it till it no longer heats up, be care full to check often, as once all the water is boiled off it will be very hard on the micro wave! Or you fuse it (Heat it to melting point)

[Edited on 10-1-2020 by XeonTheMGPony]

Does this type of hydrolysis use an excess of Naoh or equimolar amounts?

XeonTheMGPony - 15-1-2020 at 04:32

I used an excess (Read as: I didn't bother to look it up!), as I wanted it to be dry as possible, then I just used evaporated down vinegar to fully neutralize any remaining hydroxide.

Then I purify it in methanol to get rid of carbonates and gunk, some where on the board I did a solubility amount for it.

After that I got very very good concentrated acetic with minimal crap left in the reaction flask, so it was well worth the effort to recrystallize and dehydrate thoroughly

https://www.youtube.com/watch?v=rq7a5_tvBPM

https://www.youtube.com/watch?v=BzgWq7Dtn3U

[Edited on 15-1-2020 by XeonTheMGPony]

Experiment 2: Sodium Bicarbonate Complete

Lytzu - 28-1-2020 at 21:55

Alright so I've continued with this project but I am wondering if this thread should be moved to beginner as I have so many basic questions. I would like at least some responses related to these questions. I am especially interested in trihydrated vs anhydrated acetic acid. Thanks in advance.

As far as experimentation goes I just finished a run using basically the same method as before but with baking soda instead. My calculations seemed correct this time. I'm still not sure why they were off last time. It doesn't seem to have been the vinegar calculation that was the problem at least, but in case it is...

I determined the molar mass of 5% (by weight) acetic acid vinegar to be 1,200g. My understanding has been that if 5g acetic acid is in 100g vinegar solution and the molar mass of acetic acid is 60g/mol then 1,200g vinegar solution would be needed to obtain those 60 grams. This being at 1 mole. Then I just divided it into whatever fraction I needed.

This time was 1/4 mole. The solution fizzed for a LOT longer than I expected but all the sodium bicarb appeared to react. I stirred it until bubbles no longer formed. Or at least until my patience ran out. About the same time.

On the stove it also took FOREVER to boil off. Eventually I turned up the heat a notch because although it was boiling it wasn't boiling like water anymore, if that makes sense. I've read about the solution "bumping" before, I'm not sure if that's what it was or not. It's difficult to explain.

Also my vinegar was a mix of two different brands this time because I was running out. It seemed less yellow/fewer impurities but that also could have to do with scaling up a couple hundred mL.

The crystallization was also...strange. It began as like a weird crystal foam as it was bubbling. I kept the heat on it long after it crystallized and it became a lot drier than before, as would be expected. I thought about turning the heat up even higher to try to melt it and get anhydrous but it had been quite a few hours already and that wasn't in the plan. Maybe later in the week as a quick thing to do. Need to read more into it. Some input on anhydrous and how all that works would also be appreciated.

Anyway, I also bought some 85-90% indicator stop bath solution for the photo store. Need to do some calculating and reading but that is next on the list of things to try. I am anticipating MUCH less boiling time. Hopefully?

Thanks again, send advice if you've got it.

[Edited on 29-1-2020 by Lytzu]

XeonTheMGPony - 29-1-2020 at 03:36

I strongly urge recrystallizing, as there will be some junk all ways in there that needs to be removed.

B(a)P - 29-1-2020 at 11:02

I am not 100% on what your questions are that you are wanting responses on. To help you understand if your solution was bumping, it is the formation of a pocket of vapour beneath the surface of your liquid that occurs rapidly enough that it vibrates your apparatus. You can use boiling chips or even a stir bar to prevent or at least minimise the effect.

B(a)P - 29-1-2020 at 14:16

I am not 100% on what your questions are that you are wanting responses on. To help you understand if your solution was bumping, it is the formation of a pocket of vapour beneath the surface of your liquid that occurs rapidly enough that it vibrates your apparatus. You can use boiling chips or even a stir bar to prevent or at least minimise the effect.

Lytzu - 29-1-2020 at 17:12

From my previous post. I don't believe any of these have been addressed.

Quote:

Questions:

Why did I have to add extra vinegar? Were my calculatons off? Should I allowed the sodium carbonate to stay in the solution and become part of the resulting powder? Which reactant do I want more off? I would assume I'd rather have extra acetic acid in the solution because that is ultimately what I am trying to form, although this probably boils off.
The foam was the result of CO2 produced from the reaction correct? And when boiling the bubbles were just water vapor, just like when boiling water?
There was one spot on the bottom of the flask where more and larger bubbles were produced and in the beginning "sticking" to the bottom. Why is this? I have seen a similar phenomenon when boiling water.
Oh! Almost forgot. It could just be me working in my kitchen (I know, not the best but its the only available space right now) but the sodium acetate had a very slight sent. Like...biscuits? The fumes from my solution had no scent, unlike betamethyl's

Quote:

i combined sodium bicarbonate and vinegar in a pot and boiled to remove the water. It made horrible smelling fumes and went brown.

though they were using bicarbonate and I was using carbonate. Ideas?
Also,

Quote:

the trihydrate melts at 58°C, you want notice when all the water is gone since you'll still have a liquid.

I don't believe this is what happened with mine, unless it went from a syrup to the anhydrous form very quickly. Is Ubya correct that this is what happens?
Additional questions from the thread not yet answered:
If the sodium acetate I obtain is trihydrated how would that affect things? Will the acetic acid created from that be less concentrated because of the water trapped in the sodium acetate crystal? Does this change the sodium acetate itself in any relevant way? Is Ubya correct that heating it results in an anhydrous crystal?
I believe vibbzlab also had a question about if heating the sodium carbonate(?) could create anhydrous sodium carbonate. I believe he was referring to that chemical.

My assumptions to the answers are, in order of the above list

Something to do with how hydrated the sodium carbonate was
yes
impurity on the glass?
no idea, got a little bit of the same smell on this latest batch, some impurity in the vinegar maybe?
maybe?
different crystal structure in the trihydrated vs anhydrated sodium acetate, less concentrated acetic acid final product
...yes?

B(a)P - 29-1-2020 at 18:31

Quote:

My assumptions to the answers are, in order of the above list

Something to do with how hydrated the sodium carbonate was
yes
impurity on the glass?
no idea, got a little bit of the same smell on this latest batch, some impurity in the vinegar maybe?
maybe?
different crystal structure in the trihydrated vs anhydrated sodium acetate, less concentrated acetic acid final product
...yes?

[/rquote]

My two cents worth, in order of your answers.
Point one, you need to know which hydrate you have before you can get your stoichiometry right. Also it is unlikely the 5% conc is accurate.
Point two, yes about the initial CO2, but CO2 will continue to come out of solution as it gets hotter and more concentrated so not just water vapour.
Point three, could be contamination in the glass or could be an imperfection, scratches make good locations for bubbles of gas to form.
Point 4, I also do not know
Point 5, That certainly is the meeting point, but I have not done the procedure so I couldn't say.
Point 6, if you want to react the substance with something you need to know the what you have before you can work out your stoichiometry. Heating it will work, also you could make a dehydrater if you have time to let it sit.
Point 7, same as point 6

Hope this helps

Lytzu

1-2-2020 at 07:41

Thanks for all the help B(a)p.

Today I'm going to try using the stop bath solution. I also picked up some washing soda (sodium carbonate) so I will try both, but I've got some more questions.

How do I determine how hydrated the sodium carbonate is? Assuming it possibly isn't all the same hydrated compound, how do I make it all be the same? Bake it in the oven? I'd like to get the stoichiometry correct instead of just vaguely guessing at amounts.

In a similar vein, the stop bath solution also isn't an exact concentration of acetic acid. It's the Kodak 85-90% stuff. I won't try to figure it out today but in the future is there a titration or something I can do to determine more precisely the concentration? I will have to look around the forum, because there might already be a thread on that.

Oh, before I forget to ask, Xeon why methanol? This is probably a very stupid question but why can't you just dissolve the acetate back into water and filter that? Is it because the impurities will also dissolve in the water and not be filtered out?

Sorry, I'm very beginner and not understanding half of what you're saying. I've kind of been ignoring it for now until I have a better grasp on things.

I believe you were referring to this method in your first post with the ethyl acetate and sodium hydroxide Xeon? I am planning on trying that eventually. I'm sure it is a better method for obtaining large quantities of acetic acid, but both my comprehension and my technical skill are nowhere near where I would be comfortable trying it yet.

By the way, besides the acetone free nail polish remover M.E.K. Substitute should be a good source of ethyl acetate if you can get it where you are. I already picked some up while I was getting acetone.