Sciencemadness Discussion Board

Acetic acid/ sodium hydroxide

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Hermes_Trismegistus - 30-12-2003 at 01:13

Pardon any foolishness on my part here, but like most science newbies I have more enthusiasm than experience

I've been pouring gallons of aqueous CH3COO- and 2Na+ and OH- down the drains because of a particular reaction I've been running that yields a great deal of it.

How might a man (or woman) seperate out all of the Na as NaOH. and would then there be enough H+ left over to consider the resultant mixture acetic acid?

mother always said "waste not, want not!" and being the dutiful son I am.....:D

Mumbles - 30-12-2003 at 10:40

I'm a bit confused on one point. Is the aqueous acetic acid and sodium hydroxide already mixed, or two seperate solutions?

Two seperate and you're set, but I doubt this is the case from the 3rd paragraph.

It is possible in a bit of a round about way. With those two compounds mixed, you'll get water and Sodium Acetate. A simple acid-base reaction. To get acetic acid, you could boil down the solution to crystalise the Sodium Acetate. Then, dehydrate as much as you can. Add an excess of sulfuric acid, and distill. There will be some water in it'll most likely get the azeotrope, which boils at 70ish.

To get the sodium hydroxide back, convert to sodium metal or sodium oxide, then add to water, and voila you have sodium hydroxide. It'd probaby be cheaper, in terms of energy to just purchase sodium hydroxide.

Saerynide - 30-12-2003 at 15:05

Hardly seems worth wasting H2SO4 to make acetic acid when you can just buy vinegar, ditto with making Na to turn into NaOH. :P

Mumbles - 30-12-2003 at 18:59

Vinegar isn't the easiest thing to concentrate in the world. Water and Acetic Acid form an azeotrope. The azeotrope boils before water. You can't just boil the water and get acetic acid, and most of the acetic acid will boil off with the water. The only truely efficient way, that I know of, to concentrate vinegar is to distill with a fractionating column.

At 5%, you need to distill 5 gallons to get a little under a liter at 100% effiency.

Al Koholic - 30-12-2003 at 22:19

Dude, if you are just worried about keeping the post-drain environment clean then you are going to have to do more, especially if you don't want to "add more chemicals" to the mix. For one, the mixture you are putting down the sink SURELY contains acetone, bleach, and chloroform in addition to the mentioned ions. Well, maybe not so much on the bleach side but you get the picture...its highly contaminated. Secondly, I know that sodium acetate isn't all that toxic for the environment especially considering how diluted its getting if it is going into a city sewer system in which case it is probably being removed at a treatment plant.

If you are dead set on 'getting your clean on' then what I would recommend would be to start some sort of a waste 7 gallon size or so polypropylene. This is typically what I encountered in labs before. Next, now that you have collected all your waste into one area, you need to figure out what to do with it. If you are in a rural area I'm going to guess you have a burning barrel (I know all my neighbors do!!! hahaha). You can stock that with all kinds of rubbish/wood and boil off the water from the solution, evaportating the remaining acetone and chloroform as well. Hopefully for the environments sake, the chloroform doesn't escape the troposphere. Annnnyway, you will then have a nice cake of crystals which I would keep but I suppose you could burn. Convert it to something harmless like sodium carbonate I think. Anyone know what happens when you heat NaAc to decomp? I'd guess it'd give off CO or CO2, H2O and leave a mix of NaCO3 and NaOH. This is basically ashes from burning wood...

Either way you look at it though, the elements came from the earth, got messed with by someone else and put into a rearrangement. Then you got them in this form, rearranged them again and now you will be putting them back into the place they came from where they will eventually rearrange back to more stable forms. Bottom line is, why do you care about the environment because you surely aren't going to disturb it on any but human levels...

Geomancer - 31-12-2003 at 10:04


Bottom line is, why do you care about the environment because you surely aren't going to disturb it on any but human levels...

By definition, the things people care about are on human levels. You should be concerned about waste production if:

  1. You care about peoples' health and comfort.
  2. You care about the structure and diversity of life on Earth.
  3. You care about the elegance of your work.

That being said, the bulk of your contaminants (acetate, sodium, and hydroxide) are relatively harmless. Acetone too. If you're using an excess of hypochlorite, you can release most of the left over chlorine as gas and use it for something else. The dissolved chloroform is somewhat worse. CHCl3 is a common pollutant and many techniques exist to eliminate it, search for "wastewater remediation". You could probably separate it from the solution using activated charcoal. Most of the other products could be destroyed with Fenton type chemistry, but it would be hard to separate them from the acetate.
How much chloroform do you need, anyway? The best solution is simply to recycle your solvent. All these pollutants are common, a little bit down the drain won't hurt much.

BromicAcid - 31-12-2003 at 12:09


I want to know whether I should try to precipitate out the hydroxide through simple concentration, or distill off the acetate, or what?

Yes, I know it ionizes to form the hydroxide ion to a somewhat extensive extent, however it is nowhere near as much as say trisodium phosphate, the pKa is'nt terribly extensive. Acetic acid is a weak acid and sodium hydroxide a strong base, neutralization between the two is essentially complete and although you might have hydroxide ion and sodium ion in soluion you will not be able to precipiate out the sodium hydroxide from this solution. If you really want it might I suggest a ion exchange resin...

If you want acetic acid from the mixture Mumbles has already given the direction in which to go.

Hermes_Trismegistus - 1-1-2004 at 17:22

Originally posted by Mumbles

It is possible in a bit of a round about way. With those two compounds mixed, you'll get water and Sodium Acetate. A simple acid-base reaction. To get acetic acid, you could boil down the solution to crystalise the Sodium Acetate. Then, dehydrate as much as you can. Add an excess of sulfuric acid, and distill. There will be some water in it'll most likely get the azeotrope, which boils at 70ish.


unionised - 3-1-2004 at 08:00

I don't think acetic acid forms an azeotrope with water.
It's just a thought, but if you convert it to the calcium salt and pyrolyse it you get acetone which you may find useful.

Mumbles - 3-1-2004 at 13:35

Acetic Acid forms an azeotrope with water at something like 97 or 98%. It boils at 76.6 degrees. Some references say there isn't, but some say there is. I don't know which is true. I think there is one personally. Otherwise why wouldn't you be able to just heat off the water of vinegar to get glacial acetic acid? If this is possible, I've been missing out for a long time.

guaguanco - 3-1-2004 at 16:03

Originally posted by Hermes_Trismegistus
Originally posted by Mumbles

It is possible in a bit of a round about way. With those two compounds mixed, you'll get water and Sodium Acetate. A simple acid-base reaction. To get acetic acid, you could boil down the solution to crystalise the Sodium Acetate. Then, dehydrate as much as you can. Add an excess of sulfuric acid, and distill. There will be some water in it'll most likely get the azeotrope, which boils at 70ish.

Of course, with all this boiling you are consuming quite a bit of energy. You might reduce the amount of discarded Sodium Acetate (which is about as harmless a chemical as you can find), but you're burning a large amount of energy to do it. No environmental benefit overall.

unionised - 3-1-2004 at 16:04

If that's true then adding a few percent of water to the acid drops the boiling point by about 40 degrees; that's a big change. For what it's worth the CRC handbook explicitly says there isn't an azeotrope.

OTOH getting the stuff to 97% would do for most purposes and you could always freeze it to remove it from the last few percent of water.

acetic acid - water azeotrope

Magpie - 3-1-2004 at 20:35

My 49th ed. (1969) of CRC says that acetic acid and water form a low boiling azeotrope (76.6 deg C) that consists of 97% water. Therefore, the acetic acid cannot be separated from the water as pure acetic acid by fractional distillation, IIRC. I believe this is the same problem that you face with trying to get pure ethanol from ethanol-water mixtures. If this were not true you could obtain the pure acetic acid from vinegar (5% acetic acid) as Mumbles has concluded. A search on Google quickly showed that 30% acid or so could likely be obtained from vinegar by reverse osmosis. But this is not practical for the amateur chemist either due to the specialized membranes and high pressure equipment required.

unionised - 4-1-2004 at 07:50

If acetic acid forms an azeotrope as you say then you will be able to get it pure by simple distillation. (not like alcohol)
Take vinegar (about 95 % water) and distill it. The first thing to come off will be the azeotrope with 97% water. That is a higher water content than the initial mixture. This means that the concentration of acid in the still pot will rise. After all the water has been distilled out as the azeotrope (and taking a fair bit of the acid with it) all that will be left is acetic acid; this can be distilled in turn. OK there is other trash in vinegar but I can't see it making any significant difference.
The reason this works (unlike alcohol) is that you start off with a mixture which loses the unwanted component into the vapour phase. In the same way I could prepare pure water by distilling wine; boil off the alcohol (OK it takes some of the water with it) then distill the water.

I'm just not sure acetic acid does form an azeotrope. I still think that the addition of a few percent of acid dropping the Bpt of the water by nearly 30 degrees is remarkable, so much so that I would expect to have seen a reference to it. (For what it's worth, I guessed the wrong way round for the proportions of acid / water before, it doesn't matter a lot). Also my CRC handbook (1991 student edition) says that the mixture is zeotropic.

fractional distillation of vinegar

Magpie - 4-1-2004 at 22:17


I must agree that you are right about the azeotrope not preventing the aquisition of pure acetic acid in the pot. Thank you for the correction.

Assuming the azeotrope exists, I calculate that there should be about 10 g of pure acetic acid in the pot if one starts with 500 g of vinegar, neglecting consideration of the impurities that surely exist in vinegar, as you say.

As to whether the azeotrope exists or not would also seem to favor your arguments, especially since you have the later reference. If it does not, the yield would be closer to the theoretical maximum of 25 g. If I had my lab set up I would be trying this tommorrow!

Mumbles - 5-1-2004 at 15:32

Are you sure you posted that azeotrope proportion correctly Magpie? Everytime I've seen it, it is 97% Acetic Acid and 3% water, not the other way around.

To get up to glacial from this point, refluxing over CaCl2 or MgSO4 is supposed to work. Those are both anhydrous of course. A tertiary azeotrope, like Benzene should work too.

unionised - 5-1-2004 at 16:06

Does anybody have any evidence for the existance of the azeotrope?
I looked in the library at work and I couldn't find any (most recent CRC book explicitlty excludes it, Kirk Othmer doesn't mention it (but it goes on about the eutectic))

KABOOOM(pyrojustforfun) - 5-1-2004 at 20:00

<blockquote>quote:<hr>Add an excess of sulfuric acid, and distill. There will be some water in it'll most likely get the azeotrope, which boils at 70ish<hr></blockquote>I expected someone mention: no azeotrope when distilled with conc sulfuric acid => PURE acetic acid (unless acetic acid is more eager for the water than sulfuric acid!!!!:o:O:P)

KABOOOM(pyrojustforfun) - 5-1-2004 at 20:01

<blockquote>quote:<hr>Add an excess of sulfuric acid, and distill. There will be some water in it'll most likely get the azeotrope, which boils at 70ish<hr></blockquote>I expected someone mention: no azeotrope when distilled with conc sulfuric acid => PURE acetic acid (unless acetic acid is more eager for the water than sulfuric acid!!!!:o:O:P)

guaguanco - 6-1-2004 at 10:23

Some more info

acetic acid purification

unionised - 6-1-2004 at 11:07

Looks like there isn't an azeotrope if that website is right.

azeotrope verification

Magpie - 6-1-2004 at 15:29


I checked my 49th ed CRC again - it says 97% water.

My 4th ed. of Perry's (1963) and my 10th ed. of Lange's (1967) do not list this azeotrope. In fact, the Lange's specifically says "No azeotrope."

Guaguanco's reference is most interesting. Preparation of glacial acetic acid from vinegar does not apear to be a trivial problem. My text on chemical engineering processes (Shreve's "The Chemical Process Industries", 2nd ed) does not list it, but infers that it is made directly from the oxidation of paraffins. Again - not likely practical for the amateur chemist.

If someone here can come up with a reasonably cheap and easy way to make glacial acetic acid from vinegar they would likely be awarded a Medal of Honor from Vulture (or at least honourable mention).

unionised - 6-1-2004 at 15:56

Add lime
add conc H2SO4
Distill (gently)

Freeze out most of the water
freeze out pure acid.
Both of these might be improved by fermenting a higher alcohol concentration in the first place. That way you have less water to get rid of.
The bacteria will run up to about 10% acid IIRC and you might be able to do better still by adding the lime to remove the acid as the bacteria produce it.

guaguanco - 6-1-2004 at 16:42

Originally posted by Magpie

I checked my 49th ed CRC again - it says 97% water.

The 54th edition (1973-1974) says '97% water'
I trust my CRC. CRC good.:)

[Edited on 7-1-2004 by guaguanco]

unionised - 8-1-2004 at 15:02

My (1991) edition still says it doesn't.

Old CRC good; new CRC better (because they have removed some of the mistakes).

It doesn't matter how many times you check it if it's wrong.
Has anyone out there got a still, a thermometer and some vinegar? That way we can get can experimental verification one way or the other.

guaguanco - 8-1-2004 at 15:14

Originally posted by unionised
My (1991) edition still says it doesn't.

Old CRC good; new CRC better (because they have removed some of the mistakes).

It doesn't matter how many times you check it if it's wrong.
Has anyone out there got a still, a thermometer and some vinegar? That way we can get can experimental verification one way or the other.

Interesting. I have always heard it forms an azeotrope! In fact, I thought that was how 'distilled white vinegar' is formed. Odd that it took so long to overturn; I'd still like to read something that definitively settles the issue...

acetic acid production

Hermes_Trismegistus - 14-1-2004 at 09:33

Acetic acid, CH3COOH, is a colourless, waterlike liquid that has a piercingly sharp, vinegary odour and a burning taste, vinegar is its dilute aqueous sol'n and was used in the earliest languages of antiquity. Theophrastus (372-287 BC) presented a definitive study of vinegar's use in the production of white lead and verdigris manufacture. His work was closely followed by the encyclopediast Pliny(23-79 AD).

The term "Acetic Acid" was first coined by Libavious (1540-1600 BC). Many attempts to prepare "icy" acetic acid by repeated distillation were made, but the subsequent failures prompted Lavoisier to make a distinction between "acetous and acetic acid". Shortly after Lavoisier's death, the connection was proven by Adet and others. Final proof was obtained when Kolbe first prepared acetic acid from its elements in 1847.

Acetic acid (glacial) is commonly prepared from the distillation of wood in the Soviet Union, however only the credits from the distillation byproducts makes this process economic, one can imagine that the byproducts must be so great as to make this an entirely impractical process for the amateur chemist.

The historical preperation of acetic acid by dry distillation of metal acetate salts seems most practical. Many acetate salts are commonly available OTC. The most common of which is sodium acetate. If one wanted to prepare acetic acid from common household vinegar, the logical step would be to prepare an acetate salt of a commonly available powdered metal (like copper), react with the vinegar to form the corresponding acetate, evaporate the excess water, dry the salt, and then dry distill (decompose) to form the concetrated acid. The powdered metal could probably be reused indefinitely.

Much thanks to the Kirk-Othmer encyclopedia of chemical technology, third edition.

Mumbles - 14-1-2004 at 14:57

The heating of the salt can also form Acetone. Calcium Acetate specifically comes to mind. I don't know exactly what happens with +1 metal ions or transition elements. The Hydroxide and Ketene might be possible. Well, the formula works anyway. I've always been under the impression that to go from metal salt to coresponding acid you needed to add another strong acid and distill/precipitate/filter whatever the case may be.

KABOOOM(pyrojustforfun) - 15-1-2004 at 20:43

copper acetate monohydrate melts @ 115°C and decomposes @ 240°C. I couldn't find at which T it loses its water but it certainly much lower than its dec T. maybe by destructive distilling you first collect water and then acetic anhydride which react ...
Cu(OOCCH<sub>3</sub>;)<sub>2</sub>.H<sub>2</sub>O <s>&nbsp;&nbsp;&nbsp;></s> Cu(OOCCH<sub>3</sub>;)<sub>2</sub> + H<sub>2</sub>O
Cu(OOCCH<sub>3</sub>;)<sub>2</sub> <s>&nbsp;&nbsp;&nbsp;></s> CuO + (CH<sub>3</sub>CO)<sub>2</sub>O(:o)
(CH<sub>3</sub>CO)<sub>2</sub>O + H<sub>2</sub>O <s>&nbsp;&nbsp;&nbsp;></s> 2CH<sub>3</sub>COOH
wow if this really makes acetic anhydride it'll be one of the best methods (requires low temps, is simple and recyclable)

Theoretic - 16-1-2004 at 02:56

Maybe it will form acetone and copper carbonate (like the calcium salt) and then the carbonate will decompose. I hope this happens at lower temperatures than with calcium acetate, would be a very nice acetone production!

Acetic Acid

Chemtastic - 19-6-2004 at 16:56

Can vinegar (5% acetic acid) be distilled to 100% glacial acetic acid, or is there an acetic acid-water azeotrope? I haven't found anything on such an azeotrope, so I'd appreciate any information anyone might have.

BromicAcid - 19-6-2004 at 17:08

There may or may not be an azeotrope, it was the subject of discussion in this thread.

[Edited on 6/20/2004 by BromicAcid]

S.C. Wack - 19-6-2004 at 18:59

Chemists who have been dead for almost a hundred years say that acetic acid of any strength can give some glacial acid on distillation. You'll need a lot of it, because some of it will distill before you get to that point. The old references all point to Henry Enfield Roscoe and a J. Chem. Soc. article of his. Note that I said glacial and not 100%.

Given the energy needed, and the ease of other methods, this is not done, and I guess that it never was, industrially. Thus little info.

Finally took a minute (literally) to go to to find the vol/page of the mentioned article: Roscoe, JCS, 15, 270-276 (1862)

[Edited on 21-6-2004 by S.C. Wack]

tom haggen - 19-6-2004 at 19:29

I'm going to try and concentrate acetic acid at very low temperatures and see if that works. Kind of like how some people say they concentrate H2O2. I just haven't gotten around to it yet.

thunderfvck - 19-6-2004 at 22:56

Acetic acid boils pretty close to water, no? Like, a 15 degree difference...I don't know if that will make a difference, but perhaps your boiling water vapour would be tainted with acetic acid...

I always thought that it was better to form sodium acetate and then collect this salt and react it with H2SO4 while heating and collecting the acetic acid that boils off..

There must be better ways.

Even hydrolysing aspirin gives off acetic acid, and maybe that would be cheaper and more efficient for you...

Anyways, haven't put a lot of effort into making this chemical so I suggest you check out that thread..

Esplosivo - 19-6-2004 at 23:27

Glacial acetic acid (which is of very high conc, ie 99.8% - nothing is ever 100% :P) has a boiling point of approx 118 deg celcius. The solution of the ethanoic acid in water as found in vinegar assumes the physical properties of water, such that the boiling point falls to 100 deg celcius. I think the best method if one wants to produce conc. ethanoic acid is either oxidation of conc. ethanol with KMnO4/H+ or else the method mentioned by thunderfvck which I think is best:


I always thought that it was better to form sodium acetate and then collect this salt and react it with H2SO4 while heating and collecting the acetic acid that boils off..

Chemtastic - 24-6-2004 at 11:18

Will only H2SO4 work properly here, or will other acids work too?

2NaC2H3O2 + H2SO4 --> 2HC2H3O2 + Na2SO4, right?

So, NaC2H3O2 + HCl --> HC2H3O2 + NaCl?

On a sidenote:
For the production of sodium acetate, an excess of vinegar is better than an excess of baking soda, right? The vinegar remaining after all baking soda is gone should just boil off...

vulture - 24-6-2004 at 11:49

Please search before posting & continue the discussion in the existing threads.

I'm merging them.

darkflame89 - 25-6-2004 at 01:23

acetic acid and sodium hydroxide could be produced via electrolysis if you had all the time in the world and quick electricity at your disposal..

Just dump the salt into 2 bowl, dissolve them, get a salt brodge to connect the two bowls. Then, insert graphite rod to each of the bowl. You might get acid at the anode, and the hydroxide in question at the cathode. But like i said, you need lots of electricity to use, and you might have to electrolyse a few days..

TheBear - 1-7-2004 at 12:01

I had one of those moments today, maybe it's genious or perhaps it's just stupid:

Anhydrous citric acid melt at 153°C and decomposes below its boiling point at 175°C. Acetic acid boils at 118°C.

Citric acids pKa's:


Acetic acid's pKa: 4.78

What about mixing anhydous NaOOCCH3 and citric acid and heating to ~160°C and then condense the CH3COOH?

Is this method possible? If so: is citric acid bought in food stores anhydous (I'm guessing it's the monohydrate).

[Edited on 1-7-2004 by TheBear]

TheBear - 5-7-2004 at 10:27

Had another thought:

What about using ammonium acetate which has a melting point of 114C? Heating NH4OOCCH3 together with stochiometric amount of citric acid? If noone can see any problems with this procedure I will try to make some glacial acetic acid with it soon.

unionised - 5-7-2004 at 13:13

You might start to produce acetamide.

Mumbles - 11-7-2004 at 19:26

If you use acetic acid you will definatly produce acetamide. What you are describing is almost a direct excerpt from Vogel. Check out page 401. The only thing different is the reaction time. You might not allow for the full conversion with your time.

TheBear - 14-7-2004 at 12:15

Ahh I see.. Well when I want acetamid I will try that.. but for now: glacial acetic acid is the goal!

So I have to get rid of the ammonium ion, what about Cu(OOCCH3)2 which has a melting point of 115C, which is bellow the melting point of citric acid. Or will the CuAc2 start evolving CO(CH3)2? Can't find any information on this point, it was however mentioned somewhere on this board.

(Hope citric acid works for this reaction and won't decompose)

Offtopic: This experiment has been delayed due to construction of a new fume hood.

[Edited on 14-7-2004 by TheBear]

The_Davster - 17-7-2004 at 09:58

from; Textbook of organic chemistry, Noller, 1958
...4-10% of acetic acid which may be recovered by neutralizing with lime and distilling to dryness. The gray acetate of lime so obtained may be converted to glacial acetic acid by concentrated sulfuric acid. In recent years acetic acid has been removed from dilute aqueous solutions by extraction from the vapor state with tar oil(Suida process), and by azeotropic distillationusing ethylene chloride, propyl acetate, or butyl acetate to form a constant boiling mixture with the water(Clarke-Othmer process). It may be recovered as ethyl acetate by esterfying with ethyl alcohol."

[Edited on 17-7-2004 by rogue chemist]

Ca(CH2COOH)2 route to glacial acetic acid

Magpie - 17-7-2004 at 18:58

This looks like a really straigtforward and cheap way to make glacial acid. I can hardly wait to try it. As Mendeleev has stated "sulfuric acid is just so damn useful." It should be the right of every citizen to bear H2SO4.:D

The_Davster - 17-7-2004 at 22:57

No, Magpie, It should be the patriotic Duty of every citizen to bear H2SO4.

unionised - 18-7-2004 at 03:34

Magpie, I'm pleased to see you like the method based on Ca acetate. Did anyone like it back in January?

credit where credit is due

Magpie - 18-7-2004 at 09:26


:P You don't let me get away with anything!

Yes, I do remember that you suggested the acetate method some time ago. I liked it then too. I like your freezing method also.

No Title

Cyrus - 18-7-2004 at 12:34

I tried freezing 5%ethanoic acid- (sounds much better than vinegar) The whole thing froze into a solid mass. Hmm.

I made some sodium acetate (guess how;))and boiled it down to dryness in an iron pot. I will try adding H2SO4 and distilling. My concern is that sodium acetate is efflorescent. Anyone know how much sodium acetate is lost when boiling it down? Anyways, I love the smell of the process, a cross between freshly baked bread and vinegar. :)
Calcium acetate doesn't have the same problem, I might try that.

[Edited on 18-7-2004 by Cyrus]

The_Davster - 18-7-2004 at 12:56

I assume you made your sodium acetate the same way I make mine, baking soda and vinigar:D. I have about 1.5L of this solution to boil down still.:( It is taking a long time because I am using a 600mL beaker for my boiling.

Edit: anyone else notice this thread has three pages but on the "reagents and apparatus aquisition" page only 1 page and 23 posts are apparent.

[Edited on 18-7-2004 by rogue chemist]

Cyrus - 19-7-2004 at 11:52

I noticed it.

I did use baking soda and vinegar.:P

This is slightly off topic, but I think that baking soda and vinegar will make some acetic anhydride. Look at the AA thread. Sodium acetate would be there, and so would CO2, so why not AA if the soln. is cold enough. I noticed a fruity smell when doing this reaction that wasn't sodium bicarbonate or vinegar. Does AA smell fruity?

bobo451 - 19-7-2004 at 12:04

acetic anhydride has nearly the same odor as acetic acid


Cyrus - 19-7-2004 at 16:35

I just mixed about 1 g NaC2H3O2 and 2 g H2SO4 (94%) in a test tube, the sulfuric acid attacked the sodium acetate and the mix turned black. Then it was distilled for about 20 min. Actually it was more like refluxing, because the distillate kept on condensing before it even reached the condenser.:(

A ridiculously small amount of distillate (about 0.25 ml) made it to the test tube at the end of the condenser, which was in an ice bath.

Now the distillate ought to be nearly pure acetic acid, and it ought to be a solid at icy temp. It did smell and taste like acetic acid, but the stuff was a liquid, even at near zero temperatures. :mad:

What am I missing?

Perhaps the sulfuric acid started losing water during the distillation, and the water ended up in the distillate?
Near the end of the process, I noted some SO3 fumes coming off, water might have been there too?

chemoleo - 19-7-2004 at 17:06

Just to tell you- i know the MP of acetic acid is supposedly around 20 deg C. But I cooled down glacial acetic acid (analytical grade) to 4 deg C, and it stilll wouldnt crystllise/solidify. I guess if takes time, even more so if it is impure. Quite similiar to diethanolamine, which supposdely melts at 26 deg C, and yet failed to solidify at 4 deg C for 10 days...

frogfot - 20-7-2004 at 12:49

Cool, I was searching some info about toluene/water azeotropes and this thread came up :)

First of all, H2SO4/acetate salt is very hard to upscale from a testtube. Since the stuff will be solid (as Cyrus already said) and heat transfer will be a problem.
I checked mp of sodium citrate and it's >300*C so, citric acid would give same problem.
Just a suggestion, what if one uses a solvent for this, with high bp... like acetamide... :P it would make reaction mix liquid.

I have prepared 96% acetic acid from 24% like a year ago.. this wasn't that fun..
First, acid was extracted with ethyl acetate on evaporation of solvent, this gave ~60% AA. This was then destilled with ethyl acetate and 96% prod was obtained. Though it seems like alot of AA goes away with ethyl acetate. And some solvent decomposes...
There are some info about ethyl acetate/water azeotrope on the address already posted in this thread. Though they've choosen to use methyl isobutyl ketone, it seems to be more effective (I have no source for it).

Btw, obtained wet ethyl acetate was dried by freezing out the water, very effective and fast (I did this to reuse solvent..)
I have some experimental results somewhere if anyone is interested..

Oh, finally, what about using toluene to remove water? This azeotrope boils at 85*C. A bit higher than mentioned 76*C.. but my 24% AA boils above 100*C IIRC, so there seems to be no azeotrope (can't beliave noone tested this yet, gonna do this tomorrow to be sure..)...

This may be interesting:

Where they use n-butyl acetate/water azeotrope (90,2*C) to remove water from AA. Nothing mentioned about the one with AA/water..

[Edited on 20-7-2004 by frogfot]

[Edited on 20-7-2004 by frogfot]

unionised - 20-7-2004 at 13:58

A good part of the smell of neutralised vinegar is butane dione (AKA biacetyl, AKA diacetyl)

frogfot - 21-7-2004 at 11:15

Oki, tryed to distil 24% acetic acid, it came at about 100*C. Also checked CRC book (don't remember year, but it was new), it said no azeotrope formed between acetic acid and water.
However heres some interesting azeotropes:
water/toluene bp=84,1*C tol. content=86,5%(w)
acetic acid/toluene bp=100,6*C tol. content=71,9%(w)

So, I think toluene would be great to dry AA. I'm currently destilling 100 ml 24% AA with 150 ml toluene, stuff destills at about 85*C, and thats without any defligmator or column.. and pretty fast it goes :) will se how it goes..

Some results from AA drying mentioned above..

I have boiled away about 100 ml liquid, separated clear toluene and returned it back into destilling flask. This was done 4 times, and I got 77 ml water solution which showed to be 15% AA. Next 4,9 ml collected aqueous solution was 20% AA.

Remaining aqueous solution in destilling flask, that was separated after cooling was 15 ml 55% AA (it was turbid, because of some solid precipitate..).

Results is not the best, but remember, this was a simple destillation setup. Think what one could do with a fractionating column or at least vigreoux column.. Any thoats on this? It boiled at 84-85*C for me, while theoretically it should be 84,1*C and I did it pretty fast which is not adviseable when separating liquids by destillation.. IMO yield can be improved.

Btw, would it be promicing to add ethanol into the system? This way we get new azeotrope:
water/ethanol/toluene bp=74,4 water content 12%

So, we get an azeotrope with approx. same water content but it'll boil 9,7*C lower.. ofcaurse there will also form ethanol/toluene bp=76,7*C ..oh, and it would be harder to separate toluene from resulting destillate.. dunno..

[Edited on 21-7-2004 by frogfot]

[Edited on 22-7-2004 by frogfot]

Cyrus - 22-7-2004 at 09:23

Oops, if I said that the contents of the test tube (sodium sulfate, acetic acid, and sulfuric acid) were a solid, I made a mistake. When I added the sodium acetate, there was some solid, but after distilling for a minute or two, all of the solid sodium acetate had disappeared, the contents acted like a liquid, after being cooled down, they acted like a mush.

I still think this process could be scaled up pretty easily.

frogfot - 22-7-2004 at 11:19

Aha, sorry I assumed it was solid. So you had NaHSO4 in the end, cool. Since it melts at 240*C, you're right, it can be easily upscaled in a flatbottomed flask.

Btw, decomposing NaHSO4 would give off water and pyrosulfite, which would in turn decompose to SO3.. as discussed in acetic anhydride thread. So, there could happen a dilution of your product.
I'd also try with larger ammounts, since I think half of your yield was stuck on the walls..

[Edited on 22-7-2004 by frogfot]

Cyrus - 22-7-2004 at 15:53

Um, this may be a dumb question, but if the bp of acetic acid is near 100 deg. C, and the mp of NaHSO4 is more than 200 deg C, how is the NaHSO4 going to melt if there is any acetic acid left?

Could the H2SO4 might affect the phase of the NaHSO4?

[Edited on 22-7-2004 by Cyrus]

frogfot - 23-7-2004 at 06:57

AA will leave and it will gradually melt.. first by dissolving in AA, and then on its own..
Important part is that it'll melt in the bottom of flask giving good heat conductivity for the rest of mass. As example, this was actually the main problem for me when making acetone from Ca acetate.. stuff was in powder form.. and even if reaction required 400*C my 1500*C burner couldn't drive reaction to completeon..

Btw, today I checked newer book on azeotropes (year 2002) and it said water/AA gives an azeotrope with bp of 77*C with 3% AA content.. :( (Oh, IIRC someone mentioned this already)
So, this could be checked using a column..

[Edited on 23-7-2004 by frogfot]

Cyrus - 23-7-2004 at 10:39

I tried the reaction again with larger amounts- 17 g sodium acetate, an excess of sulfuric acid, when added the sulfuric acid attacked the sodium acetate again, leaving a black gunk/liquid.

I am assuming this residue is carbon, how can this be prevented? I don't like losing all of my acetic acid yields just to make carbon!

unionised - 24-7-2004 at 06:08

The black gunk is probably from impurities (sugars etc) left over from the vinegar being attacked by th sulphuric acid.

kclo4 - 28-1-2005 at 22:30

woudent Aluminum hydroxide and sodium acetate work (possiply work to make other acids to)

[Edited on 30-1-2005 by kclo4]

acetic acid from vinegar

Magpie - 1-2-2005 at 20:09

I have been attempting to obtain glacial acetic acid from vinegar (5% acetic acid).
I first tried a simple distillation of 250 mL. Everything came over at 100 C and I left 12 mL in the pot. The distillate and the pot residue smelled about like the starting vinegar.

Then I tried a fractional distillation using a column packed with a stainless steel scouring pad. Same results although this time I took the pot to very near dryness. Again the temperature at the still head stayed at 100 C. Based on these results it looks like there is indeed an azeotrope at about 97% water. However, if there is I would say its temperature is very close to 100 C, not the 76 C reported in my CRC! I will go to the library to see if I can get some help on this issue.

I also prepared some calcium acetate using 250 mL vinegar reacted with a stoichiometric amount of hydrated lime [Ca(OH)2]. Yield of Ca(Ac)2 is 79%. I will try to make glacial acid out of this using H2SO4.

Has anyone else been doing some experimentation in this area?

vinegar distillation

Magpie - 2-2-2005 at 12:46

Vinegar at 5wt% acetic acid is 1.6 mole% in acid. According to the vapor-liquid equilibrium data I found at the library the liquid and vapor composition at this liquid concentration are very nearly equal. This I believe makes concentration of the acid by distillation impossible or at least impractical.

I also checked "Azeotropic Data" by Gmehling et al (1994). There were 52 citations! Pressures ranged from 1.07 kPa to 5864.7 kPa with 31 citations at 101.7 kPa (atmospheric pressure). All of the references indicated "none" for existence of an azeotrope.

Needless to say I'm giving up on distillation of vinegar.

Magpie - 3-2-2005 at 16:02

Attempted to make glacial acetic acid out of the Ca(Ac)2 I made a few days ago. Decided to use hydrochloric acid instead of sulfuric after reading Muspratt in the MSDB library (I love that reference). Set for distillation using a fractionating column and stainless steel packing (scouring pad from Fred Meyer). Noticed that the column was tending to plug and a dark green (Malachite) colored liquid flowed down the column when I cut the heat. Fearing a plugged column with consequent flask or colum burst, I aborted. Examination of the packing showed that it had been nearly eaten away in the lower part of the column. I saved the pot liquid and will try again when I get some glass beads.

Magpie - 4-2-2005 at 19:31

I found some glass beads (hollow cylinders about 2mm dia x 7mm high) at the local bead shop. Repacked my Hempel fractionating column with the beads and distilled my crude acetic acid. All the water came off at 100 C then the nasty mix started foaming and looked like it might plug the fiberglass support plug in the bottom of the column. So I aborted. Also at that time there was a strong smell of acetic acid in the lab! What's left in the pot is very crude glacial acetic acid (~ 10 mL). The Fe, Cr, and Ni acetates generated from the old stainless steel packing are contributing the black/dark green color I believe. The acid quickly froze at room temperature. I may try to clean this up and redistill. I'll use a simple distillation as all the water is gone. :D

There is a solid mass of CaCl2 mixed with my acetic acid in the bottom of the flask. I think it's time to scrap this batch. I may try another larger batch next week.

[Edited on 5-2-2005 by Magpie]


Magpie - 5-2-2005 at 08:39

Re: glacial acetic acid preparation

There's a big difference in following some school recipe and striking out on your own into a long abandoned method. For one thing you have to be much more on your guard to protect your safety and that of your equipment.

I think removing the water from my reaction mix by fractional distillation was a mistake. Once that water was gone there wasn't much liquid left in the flask but a whole lot of CaCl2. If I would have kept the heat on another minute I could likely have cracked my expensive Kontes 250 mL RBF with ground glass joint. That would have really pissed me off. I've also told my wife that if she ever sees me in the lab without my safety glasses on to kick my ass!

frogfot - 5-2-2005 at 09:51

What if one dries dilute acetic acid in excicator above CaCl2? The air inside will be saturated by AA and water, while only the latter will stick to the drying agent..
Will take ages, but this should theoretically work..

Magpie - 5-2-2005 at 11:27

Frogfot what you suggest might work. I don't know if CaCl2 is more hydgroscopic than acetic acid, however.

My current plan is to proceed as before only use a double batch (500 mL) of vinegar charge. Once I get the Ca(Ac)2 powder I will react it with 31.45% HCl (muriatic acid). I will then codistill off the water and acetic acid using a simple distillation. Muspratt says this distillate will be about 45% acetic acid IIRC. Then I will fractionally distill this distillate. That way I'll get rid of the the CaCl2 and be working with a clear product.

Icarus - 5-2-2005 at 19:57

Why not mix your calcium acetate with sulfur and do a dry distillation?

Magpie - 6-2-2005 at 11:45

Icarus says:


Why not mix your calcium acetate with sulfur and do a dry distillation?

I don't understand how this would produce acetic acid. Please provide chemical equation.

unionised - 6-2-2005 at 11:58

It would make sense if he meant sulphuric (acid).

Icarus - 6-2-2005 at 19:54

I don't understand how this would produce acetic acid. Please provide chemical equation.

Apologies, I was trying to throw up some ideas from left field and I thought maybe some dry distillation salt combination could be utilised.

I was thinking about the dry distillation of Calium Phenylacetate and Calcium Acetate with sulfur.

Magpie - 6-2-2005 at 21:10

Unionized: Yes sulfuric acid should work as indicated by Muspratt.

At first I was going to use sulfuric acid as I didn't have any hydrochloric. But after closer reading of Muspratt I decided to try muriatic first - so a trip to Home Depot for muriatic. I felt that I was less likely to have flask fouling problems with the CaCl2 byproduct than a CaSO4 byproduct. Because I distilled too near dryness cleaning up my RBF still took a little work even with the water soluble CaCl2.

I was initially worried about HCl gas in my distillate until I convinced myself that HCl doesn't exist in the reaction mix due to Ka = 1.7 x 10^-5 for HAc. I believe that fact is what makes this procedure viable.

frogfot - 7-2-2005 at 09:46

So theoretically, using 37% HCl and an acetate one would get 49% AA. Sounds nice.

Gonna test H2SO4/Ca(AcO)2 with denaturated alcohol as solvent, this should work..

EDIT: Oh, I think I'm bugging.. using alcohol as mentioned above would probably give quite big amounts of ethyl acetate.. since not all sulfuric will react at once..

EDIT2: Oki, did an experiment with ethyl acetate as solvent. As far as I know nothing will happen to the solvent. Of caurse it can decompose, but it shouldn't, due to pretty high AA conc.

I mixed 100 ml EtOAc with 10 ml 96% H2SO4, cooled to room temp and added 28,49 g Ca(AcO)2 in portions while stirring. This gave a thick white goo, which got a sticking AA smell above smell of EtOAc (this doesn't say anything).

Now if my sences didn't exagerate this, I've noticed couple of things:
After adding first portion and shaking, there appeared a turbidity, composed of flaky solids (prolly CaSO4).
Another observation is that added Ca(AcO)2 particles had somekind of transparent shell (around particles). I'd say this could be the Ca(AcO)2 reacting with acids (H2SO4, HSO4-, AA). The dissociation of a polar salt into quite nonpolar solvent may give this funky shell.

Gonna reflux this tomorrow.. One thing that's good about using EtOAc is that once it's distilled off, it will take all the water as an azeotrope. Unfortunately some of the AA will come over too, but the solvent may be reused (after freezing out the water, which I've already tested in another exp).

Now theoretically this should give 21,63 g AA, quite impractical due to big volumes of reaction mix (~170 ml). But by now, it seems to be more practical then do extractive distilation of 24% AA by EtOAc(it works to give 96% AA, but it's tedious, more time and solvent is needed while solvent is consumed to noticeable extent)

[Edited on 7-2-2005 by frogfot]

[Edited on 7-2-2005 by frogfot]

[Edited on 7-2-2005 by frogfot]

neutrino - 7-2-2005 at 14:26

Remember that mixing conc. Sulfuric acid and ethanol will give you a good deal of ether. Don’t distill your product to dryness because of the inevitable peroxides formed .

Mumbles - 7-2-2005 at 14:34

I think you would be distilling the acetic acid off before the ether would go. Even at room temp it will form ethyl hydrogen sulfate. This combines with a molecule of ethanol to give ether and the sulfuric acid back at highish temps(130ish I believe). It seems that this compound would react with any acetic acid to form ethyl acetate anyway.

[Edited on 2-7-2005 by Mumbles]

Magpie - 7-2-2005 at 21:53

Made a new batch of Ca(Ac)2 today using a 1L charge of vinegar. It took 2.5 hours to evaporate off the water using a stirrer-hotplate at full heat.

Magpie - 8-2-2005 at 16:55

Weighed my oven dried Ca(Ac)2 and calculated an 88% yield based on the assumption that my vinegar was 5% acetic acid.

Reacted the Ca(Ac)2 powder with a stoichiometric amount of muriatic acid. Then did a simple distillation. It was very interesting. The temperature came up to about 105C fairly soon and then stayed at nearly 108C for most of the rest of the distillation. Toward the end the pot got foamy and the distillate temperature dropped to 104C (I don't understand this). Anyway my equilibrium data from the library for the acetic acid-water system says that at 108 degC the distillate should be about 83 wt% acetic acid. I have an estimated 60 mL of clear distillate now. Tommorrow I will do a fractioinal distillation in an attempt to obtain glacial acetic acid. BTW the pot remains were dark brown and were essentially all CaCl2 crystals. I did not take it to dryness (I'm a chicken).

As an aside I'm thinking that the use of sulfuric acid instead of muriatic acid may have a definite advantage. That is, the CaSO4 precipitate formed after the reaction with Ca(Ac)2 could be removed right away by filtration. Then you could go directly to fractional distillation for the glacial acetic acid. I may try this if I'm not too sick of the smell of vinegar by then. :)

neutrino - 8-2-2005 at 17:22

Calcium sulfate precipitates like this are very hard to filter out and tend to form a gel-like mass very quickly. Unless you have a good vacuum source, I wouldn’t recommend it.

Magpie - 8-2-2005 at 19:25

Neutrino: When filtrations get tough I go to diatomaceous earth.

unionised - 9-2-2005 at 14:00

Have you got the concentration above the eutectic point? If so then you might want to think about freezing out acetic acid.

Magpie - 9-2-2005 at 16:18

Today I did a fractional distillation of my water/acetic acid from yesterday. Again there was a surprise. My first problem was bumping so I stopped, cooled down, and added some pottery shards. Started back up with nice even boiling. But then my column was always flooded (5/8" ID) and there was boiling liquid all the way up the column. The temperature at the still head was 102C and stayed there. More liquid was pumped over than distilled over. But I continued as the temperature wasn't too high. I stopped when the pot had 19 mL remaining. This may not be pure acetic acid but my bet is that it is pretty high test (>90%). I placed it in my embossed "acetic acid" bottle and set it in the garage to see if it freezes. When I get some AgNO3 I will test it qualitatively for Cl-. So my overall yield is 19 mL "glacial acetic acid" for a 1L vinegar charge. This is (19)(1.06)/50 x 100% = 40% yield. I think I could do better the next time. Especially once I get the kinks out of my fractional distillation packing and support.

Unionized: I'm sure I'm rich enough to do a freeze separation. What is the eutectic concentration BTW?

Magpie - 9-2-2005 at 22:03

I'm afraid that my previous estimate of acid strength is incorrect. I have the solution sitting outdoors and it appears to be just now forming some crystals. Temperature is just below O deg C. Based on a qualitative binary phase diagram for water-acetic acid the concentration can't be more than about 50%. :(

[Edited on 10-2-2005 by Magpie]

bio2 - 9-2-2005 at 22:56

Unless a person is trying to recycle acetic acid for "fun"; given the price ($108 for 22L ACS glacial) it just ain't worth it. Hell I pay almost that much for high grade water. Yea, I know it's too much. Drums are way cheaper.

Yea, I thought I was smarter than them too and just had to try it although I was told what everyone already knows, FORGET IT.
I use 6 liters per reaction of GAA so naturally I just had to try it myself., of course finding out that they were correct. After 3 or 4 days of fractionating through a 600mm column, painfully slowly, large reflux ratio, etc. etc out of about 10 liters of 50% I ended up with maybe 750ml
of about 98% acid , not even glacial. Probably spent more on electricity than the stuff is worth. Vinegar to glacial? LOL good luck and have a fun time learning to distill.

The freezing method ain't so great either unless your after 100ml and you can fit in the freezer while separating it.

Azeotropic water removal works much better but still ain't worth it without a dedicated facility. Try PhEt or PhMe2 as they take over 30% water with them..

Magpie - 10-2-2005 at 09:29

Bio2 thanks for your valuable input. You have confirmed the difficulties I'm experiencing. But for me this never was about economics or effort. This is about fun, the challenge, and developing my skills using a relatively non-hazardous preparation. :D

My earlier report of crystals at 0 deg C was false (wishful thinking) so I put my acid in the freezer overnight which is at -17 C. No crystals formed there either. So at least I know it is not pure water! Based on my rough binary phase diagram my best estimate now is 18-38 mole% acetic acid. The eutectic composition looks like about 30 mole% acetic acid on my rough binary phase diagram.

I have a primary acid standard, potassium acid phthalate, given to me gratis a number of years ago. I'm going to make an NaOH standard solution and then titrate my acetic acid to determine its strength.

[Edited on 11-2-2005 by Magpie]

titration results

Magpie - 10-2-2005 at 20:08

Titrations went well using potassium acid phthalate, Red Devil Lye (NaOH), and phenol red indicator. BTW Red Devil Lye appears to be a good quality source of NaOH for the home chemist.

Using 0.17N NaOH my titration determined the acetic acid to be 12.5N, which is 70 wt% or 41 mole%. At -17 C in my freezer it must have been just on the verge of forming crystals of acetic acid.

JohnWW - 11-2-2005 at 00:04

"Lye" is the term commonly applied to solution of potassium hydroxide, KOH, rather than NaOH. It was originally made by the lixivation of of the ashes of burnt vegetable materials, which contain potassium, and used mostly to make lye soap. So either the makers of Red Devil Lye are using misleading advertizing if their product is in fact NaOH, or you are mistaken about what it contains.


ballzofsteel - 11-2-2005 at 03:07

Hows about dry distillation of an acetate salt with sodium bisulfate.
Has this been mentioned?

Magpie - 11-2-2005 at 09:13

JohnWW: The Red Devil Lye label says "contains sodium hydroxide." I have always heard commercial grade sodium hydroxide referred to as "lye."

Ballz: You pose an interesting route to acetic acid. Ka for HAc = 1.7 x 10^-5. Ka for HSO4- is 1.2 x 10^-2. Therefore the Ac- wants that hydrogen 1000 times worse than the SO4--. It seems like it should work. I would add a little water though just to facilitate things but I don't really know if it would be necessary. Do you or anyone else have experience with this route?


ballzofsteel - 11-2-2005 at 15:10

No, no personal experience with producing acetic acid this way.
I think I remember reading it somewhere??
Sodium bisulfate is my chemical of the week is all.Many valuable applications for this neat and readily available pool chem.
A great de-hydrating agent or replacement for H2SO4!
I know for sure it works with making formic from a formate,also pyruvic acid from tartrate,hcl gas from salt and the list goes on.

Commercially available,it is hydrated so you could use as is, if water is what you want.I dont think it would be needed though.It has a rather low melting point.

neutrino - 11-2-2005 at 16:01

Originally posted by ballzofsteel
I remember reading it somewhere??

Possibly, you’re confusing this with the method where an acetate is refluxed with sodium pyrosulfate (produced by heating the bisulfate), giving you acetic anhydride. There is a thread on this at RS.


ballzofsteel - 11-2-2005 at 16:39

No,I dont think so.

But that sound a hell of a lot better if it can be done!
Do you have a link to this RC thread?

S.C. Wack - 11-2-2005 at 16:43

No, ballzofsteel is correct, its JohnWW who's off. Lye is not and never has been a solution of KOH.

Icarus - 11-2-2005 at 17:34

This might have some interesting information if someone can find it.

Brown, W.D. (1963). Economics of Recovering Acetic Acid. Chem. Eng. Prog. .

Unfortunately, my local library is poorly stocked.

[Edited on 12-2-2005 by Icarus]

JohnWW - 11-2-2005 at 19:09

(Originally posted by S.C. Wack)
No, ballzofsteel is correct, its JohnWW who's off. Lye is not and never has been a solution of KOH.

Wrong! It is Ballzofsteel and SCWack who are off. I entered the single word "lye" in Google and got 493,000 results. The relevant results were very much in favor of "lye" meaning KOH solution, not NaOH, most especially that made the old-fashioned way, by lixivating wood ash, which contains at least 10 times more K than Na, and mostly used to make "soft" soap, see , , , , , which are among the highest-ranking results.
Some of the relevant results - the less authoritative ones - allowed that it may refer to either KOH or NaOH solution indiscriminately, and only a minority (mostly the least authoritative) applied it exclusively to NaOH or did not specify which.

Wrong again?

ballzofsteel - 11-2-2005 at 19:25

Hold your horses thar big fella!

Now,Im quite accustomed to being not 100% right all the time(well a lot of the time).But please wait until Ive made a statement about a certain topic before you go trashing me all over,saying Im WRONG and that.Youre giving me a complex.
Seriously though,I dont mind.Ill just concider it pre-emptive nay-saying for future Wrongs:P

I always thought Lye was NaOH.There you go.

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