Sciencemadness Discussion Board

Acetic acid/ sodium hydroxide

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S.C. Wack - 11-2-2005 at 20:52


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,

So lye gets a lot of hits. So?
Lye sodium gives twice as many hits as lye potassium. This should tell you something right there.
What relevant hits are those? You cant be talking about the links that you provided. Nowhere do they say that lye is only KOH, and excludes everything else.
Yeah, that brainydictionary is very authoritative. Grasping at straws.
Try reading the old books. Nowhere do they say lye, they say potash lye or soda lye. They say caustic potash and caustic soda when talking about the hydroxides.

If the modern definition of lye has been skewed, it has skewed in favor of NaOH, for whatever reason. Yet this does not mean that NaOH means lye either, although no one is mistaken about what the Red Devil product contains.

You seem to have grasped that the word originally meant the leachate of ashes, in other words ---- potassium carbonate and other stuff, depending on the ashes. Not the hydroxide.

They way you crudely try to twist things in your last post is even more laughable than most of your posts, and as usual it has little to do with the subject at hand.

Originally posted by JohnWW
SC Wack, you are just a troll! You have given yourself away, by your utter irrationality.

Originally posted by JohnWW
"Lye" is the term commonly applied to solution of potassium hydroxide, KOH, rather than NaOH.

to which i said: bull, now and always.

to which JohnWW says: UTFG

to which I said: lye "potassium hydroxide" -"sodium hydroxide" = 2440 hits
lye "sodium hydroxide" -"potassium hydroxide"= 17,700 hits. Yes, it is obvious that the Google hits are very much in favor of KOH, not NaOH. To JohnWW at least.

he refuses to accept what it says

which is, KOH is not and never has been the definition of lye. It has been included in the wide spectrum of lye, and so has NaOH, and NaOH is very commonly called lye. I'm the one being irrational? I'm the troll?

[Edited on 12-2-2005 by S.C. Wack]

JohnWW - 11-2-2005 at 22:44

SC Wack, you are just a troll! You have given yourself away, by your utter irrationality.

frogfot - 12-2-2005 at 06:46

Oki, the first H2SO4/calcium acetate/ethyl acetate experiment failed. Neutrino, you were right about filtering.. it was a real pain to filter CaSO4 with vacuum.. and on drying the solids was nothing like CaSO4, it was composed of soft crystalline mass. So, I guess reaction is long from compleate.. and I've abandoned the filtrate. I have also recovered zero solvent.. all was adsorbed on CaSO4 (it separated on addition into water)

Now some happy news. Got another idea that I recently tested, and it seems to work. This time I wanned to use sulfuric acid with an acetate that will form a soluble salt in alcohol. Reaction was performed in cold to prevent esterification.

I took 9,42 g sodium acetate (partially crystalline, so it should have some crystall water..), added it to 30 ml 92% ethanol and while cooling the slurry in ice bath, I added 3,1 ml 96% H2SO4 dropwise, while stirring.

On checking of pH in solution, it remained around 5 at all times, which means reaction is fast. This is basically because Na2SO4 is slightly soluble in alcohol, so it doesn't build up a coating around acetate salt.

Then I made attempt to filter the slurry, but this took alot of time. So, I transfered all mess into a beaker, heated and cooled it slowly hoping that the small Na2SO4 crystalls will recrystalize to bigger. After, filtration went a bit faster but it was still slow..
After washing solids with ethanol, I distilled off fractions of 79-95*C and 95-124*C (5,73 g). There remained couple milliliters of brownish liquid.

There was no indication of ethyl acetate formation.

Then I titrated all products:

Fraction 79-95*C contained 1,1g AA
Fraction 95-124*C contained 2,68 g AA
Na2SO4(dissolved in water) contained 0,21 g AA

Theoretically I could get maximum 6,896 g AA, so I got 39% yield in form of 5,73 g 47% AA.
At first this seems to be crap, but, there are many buts...

First, this was a very small batch. Second, my ethanol (92%) and sodium acetate contained quite alot of water. This could be done better if I used more concentrated denaturated alcohol and predried the acetate. Third, quite alot of AA destilled with alcohol. Though this shouldn't be problem since it can be reused in next batch (ethanol doesn't even have to be dried).

One serious problem is the size of Na2SO4 crystalls.. How would one make them bigger? The only idea I can come up with, is to let it sit for days.. maby they will regrow.. or?
Bigger crystalls would simplify the filtering and most probably decrease the adsorbtion of AA on crystalls surface.

Time to make more sodium acetate..

Magpie - 12-2-2005 at 15:43

Have you tried laying down a 2-3mm bed of diatomaceous earth (kieselguhr) on your filter paper using a water slurry prior to filtering your Na2SO4 slurry?

[Edited on 12-2-2005 by Magpie]

frogfot - 13-2-2005 at 01:02

Nope, I couldn't actually find it in my previous attempts. Would usual clay do the trick?
But then again, Na2SO4 particles would still adsorb alot of AA...

Magpie - 13-2-2005 at 11:43

My guess is that clay would definitely not work. Sand would be much better. The reason diatomaceous earth is so good is that it's made of of silaceous intricate skeltons of very small sea creatures (diatoms). This provides a very open structure for filtering. I visited my friendly swimming pool supplier. He happened to have an open 50 lb bag of Celite. He gave me a coffee can full for nothing as I was buying some other chemicals. :D

FrankRizzo - 14-2-2005 at 14:20

SEM image of Diatoms:

Quince - 11-6-2005 at 23:39

Originally posted by frogfot
Now some happy news. Got another idea that I recently tested, and it seems to work. This time I wanned to use sulfuric acid with an acetate that will form a soluble salt in alcohol. Reaction was performed in cold to prevent esterification.

I tried this, making the sodium acetate from vinegar and baking soda, and using isopropanol instead of ethanol (I have none of the latter). It didn't work, with the distillate (all of which came over at 79*C) having no acidity at all, and the brown shit left in the flask smelling worse than the inside of the devil's asshole.

I need it to make lead acetate, as I don't have any calcium sources to make calcium acetate->copper acetate->lead acetate as discussed in another thread.

[Edited on 12-6-2005 by Quince]

Glacial(?) acetic acid distillation

The_Davster - 11-11-2005 at 22:52

A pile of sodium acetate produced by boiling down baking soda and vinegar was heated over an alcohol burner untill it was dry. It had to be stirred often or it would carbonize.
Based on a procedure from an old chem lab manual from the uni library, 20g of the now anhydrous sodium acetate was added to a 250mL RBF. 12-13ml of 98%H2SO4 was measured out separatly. A 19/22 distillation setup was assembled and the H2SO4 was poured ontop of the sodium acetate. The rxn was exothermic from the start and white fumes were given off and some of the sodium acetate darkened. The mix was distilled over an alcohol burner. White fumes were given off the entire time the distillate came over(pic 1 and 2 in attachment). I have no idea what these fumes are, they seem to either dissipate or condense in the recieving beaker. The material left in the distilling flask was crushed with a stirring rod and washed out as I did not want it so solidify completly in the distilling flask and be never able to remove it. It seemed rather carbonized. Yield of suposedly glacial acetic acid was 8.1mL.

[Edited on 12-11-2005 by rogue chemist]

distillation.JPG - 54kB

Magpie - 11-11-2005 at 23:17

Rogue I tried a very similar procedure only used vinegar and hydrated lime [Ca(OH)2] to form calcium acetate. I dried this in my drying oven. Then reacted it with HCl. I thought about using H2SO4 as the formation of insoluble CaSO4 should drive the reaction. But I chickened out as I thought I might have to scrape this out of my RBF. I then set up a fractional distillation column to get glacial acetic acid. I got some strong acid (60% IIRC) but not glacial. I concluded that I did not have fine enough control on the temperature gradient on my distillation column. So I want to try this again when I can insulate and control the column better.

Your situation seems strange to me with those white vapors and carbonized gunk. I wondering if the H2SO4 is being too agressive here.

The_Davster - 11-11-2005 at 23:36

Well, I do not really expect it to be glacial as the procedure called for changing the recieving flask when the temp reached 116C, but my thermometer ends at 115(I need a better one)
I hope those fumes are just uncondensed acetic acid (thats what they smelled like with some weird smell mixed in) as a fair ammount of these were released into my lab. I don't know about the H2SO4 being too agressive as the procedure can have 7.5% oleum used instead of the conc. sulfuric acid. Perhaps it is even normal, but as it is a lab guide it wants students to find this out for themselves.

mrjeffy321 - 12-11-2005 at 13:58

When distilling the Sodium Acetate and Sulfuric acid mixture to make Acedic acid, do you keep going all the way to dryness or just short of dryness?

How do you measure then final concentration of the condensed acedic acid solution?

Magpie - 12-11-2005 at 14:33

When using a fractionating column the temperature at the still head should be the normal boiling point of acetic acid. When it changes to 100C you're getting water.

Titrate your acid or perhaps measure its density & use a table to find concentration.

mrjeffy321 - 12-11-2005 at 16:46

what about this, skip the distilation process and then cool the solution to crystalize some of the Sodium ___ out.

Say with this reaction using Sodium Acetate and 10 Molar Hydrochloric acid.

NaC2H3O2 + HCl --> HC2H3O3 + NaCl

so say you start out with 1 mole of Sodium Acetate (82.04 g) and add 100 mL of 10 Molar HCl to react it completely to form Acedic acid and Sodium Chlorde.
You will be left with a solution of Acedic acid and NaCl dissolved in about 100 mL.
This makes for a rather salty 27% acedic acid solution, and actually, with NaCl solubilty at about 36 grams / 100 mL of water, not all of the NaCl will dissolve. You can remove some more of the NaCl by cooling it, further raising the concentration of the acedic acid.
Given it isnt pure and it isnt anywhere near the concentration some people are aiming at, but it saves the distillation step.

neutrino - 12-11-2005 at 17:40

>When using a fractionating column the temperature at the still head should be the normal boiling point of acetic acid. When it changes to 100C you're getting water.

It would be the other way around. Acetic acid has a normal boiling point of 118*C, so it would come over last.

Magpie - 12-11-2005 at 19:06

Of course you are right. I just couldn't remember which had the lower boiling point.

jimmyboy - 28-5-2006 at 13:57

well its pretty clear that using sulphuric acid on an acetate just results in a black carbon mess -- the two i would like to try sometime would be oxidation of ethanol (some everclear would work) or dry distillation of copper acetate -- the first one seems the most promising - mixing some alcohol with excess permanganate and adding it to sulfuric would work - i was wondering if manganese dioxide would work just as well - much cheaper but not nearly as potent - just add more? -- anyone tried these methods?

ordenblitz - 28-5-2006 at 17:40

I have distilled glacial acetic acid many times from the reaction of sodium acetate and sulfuric acid. It is quite easy infact. I can not figure why you are having such problems.

12AX7 - 28-5-2006 at 18:05

NaCH3COO + H2SO4 == NaHSO4 + CH3COOH, or is it more like
2NaCH3COO.3H2O + H2SO4(aq?) == Na2SO4 + CH3COOH ?


Fleaker - 29-5-2006 at 17:47

My friend tried the ethanol to ethanoic acid reaction with a strong oxidizer and 1 drop of sulfuric acid. He did not get any appreciable vinegar smell, rather the smell was quite fruity, almost like the smell of fresh apples (I too smelled it). That was using potassium permanganate and heating it to 50*C and refluxing in a test tube. There was oxidation and reduction because the solution lost its purple color and a sludge of MnO2 was seen on the bottom of the tube. Based on his results which I saw for myself, I do not think you can get ethanoic acid that way. You're welcome to try for yourself jimmyboy.

jimmyboy - 30-5-2006 at 09:57

hmm - well i read all the posts in this thread and alot of people are coming up with black mess in their reactions - possibly unclean acetate from the vinegar/sodium bicarb reaction i would guess - a "fruity smell" you say? hmm definitely acetaldehyde --- was their any special setup to your reaction orden? you said you have done it a number of times -- or just dump the sulfuric in the acetate and distill..

ordenblitz - 30-5-2006 at 16:12

It's been a while since I have done it so I am working from memory.
They first time I did this I used exactly what you are using, NaHCO3 + Vinegar produced sodium acetate. I think mine had a slight excess of bicarbonate but it didn't matter since I used an excess of H2SO4.

I remember that the reaction started very fast giving me little time to assemble the condenser to the flask. The next time... I pre-chilled the acid before adding it to the flask. While this reaction proceeds similarly to making HNO3 from nitrates and sulfuric, the heat required is less so you should proceed slowly and judiciously with the burner.

[Edited on 31-5-2006 by ordenblitz]

Magpie - 29-9-2006 at 18:02

I have reduced some 5% strength vinegar to 127g of damp sodium acetate. I then reacted this with 25mL of sulfuric acid (17.5M), cooling as required. There was a lot of Na2SO4 being formed so added 27mL of H2O just so I could keep it stirred on my magnetic stirrer. Then I removed the Na2SO4 crystals by decantation/filtration. I now have about 125mL of solution that I estimate is 50%H2O/50% glacial acetic acid.

For ordenblitz: How do I get the GAA isolated now. By fractional distillation, or what? How exactly did you get the GAA?

ordenblitz - 1-10-2006 at 16:05

I took the extra time to not only boil off the excess water in the sodium acetate but also cook it down to absolute dryness with heat before making the GAA. It was then a simple matter of adding the H2SO4 to the acetate in a boiling flask and distilling the product with a bit of additional heat.
I never made it the way you are proposing and I cant say if its easy to remove the water after. But I can tell you that removing the water from the acetate is easy.

Magpie - 1-10-2006 at 19:16

Thank you , ordenblitz. I'm not a fan of trying to separate water from acetic acid. I will be giving your method a try. ;)

Magpie - 4-10-2006 at 15:58

I have finished reducing 1 gallon (3875mL) of "triple distilled white vinegar," which is 5% acetic acid, to a much more concentrated form: 87%. It is not yet glacial (99+%) acetic acid but it's definitely getting there.

I made this in 3 batches. The 1st was sort of a hodge-podge of CaAc + NaAc. The NaAc was taken back from a diluted HAc. So this batch (30mL) wasn't too efficient.

The 2nd batch was made from NaAc taken to dryness at about 140C. This yielded about 50 mL.

The 3rd batch was the same as the 2nd only I fused (melted) the NaAc first, then let it solidify, then reground it to a powder. This last step, I think, was a waste of time as it didn't change anything. This also yielded about 50 mL. I combined all distillates for a total 130 mL for a yield of (0.87)(130)/[(3875)(0.05)]x100% = 58%. It would have been better if I hadn't messed around with the 1st batch.

Each batch was done with about 80g of anhydrous NaAc and 60mL of concentrated sulfuric acid (Rooto) in a 500 mL RBF assembled to a simple distillation setup.

In all cases I placed the RBF with powdered NaAc in an ice-bath while adding pre-chilled (w/ice-water) acid. There is a lot of heat generated during this step. Also there is a lot of smoke generated. I don't know what this smoke is. Once this subsides I heated the mix with a bunsen burner as required to get a steady generation of HAc collected as a distillate.

I also noted what others have: generation of a lot of carbon black. Also there are smoky odors like that of a wood fire. Perhaps this is some of the "empyreum" spoken of by the early chemists?

This preparation was a lot of work. It has to be a labor of love as you can buy glacial acetic for reasonable prices. At least I can say that this method beats my previous preparation with calcium acetate and HCl which yielded 65% acid.

[Edited on 6-10-2006 by Magpie]

Magpie - 6-10-2006 at 14:23

Not satisfied with 87% acetic acid I continued my quest for homemade glacial acetic acid. I tried to freeze it to glacial plus eutectic solution by taking it down to -26C. It froze alright but the eutectic solution was occluded in the acid and therefore was not separable.

So I went back to trying fractional distillation. I packed my little 8" (20cm) column with broken glass pieces from a smashed food jar. I collected three clear distillate cuts: the 1st was 20mL at 101-104C, the 2nd was 45mL at 104-110C, and the 3rd was 50mL at 110-116C. I left a few mL of brown solution in the pot. So, it is tough to get glacial acid. If I would have had a longer column and glass raschig rings (or other optimized packing) I might have done better.

It amazes me that by concentrating something I put on my salad I can make this really wicked (and useful) reagent. :D

[Edited on 7-10-2006 by Magpie]

not_important - 6-10-2006 at 21:00

Freezing must be done at a controlled rate. For some materials a very slow decrease in the temperature will result in the excess componant freezing out in a fairly compact layer, but not always.

The alternative is to freeze to a slush, and vacuum filter or centrifuge to separate the frozen from the eutectic.

Even after freezing distillation is a good idea, as it cleans the acetic acid up. Even if you distilled it at the start, from adding acid to acetate, if there was more than a few percent of water then other organics will come over in the distillate.

Magpie - 7-10-2006 at 19:32

@not_important: Thanks for the information on freezing. I was wondering if there was something I was missing.

Yes, the distillation was good just to clean up the acetic acid. There was a fair amount of carbonaceous junk left in the pot. No wonder the commercial white vinegar is "triple distilled."

I made another attempt to enrich the 2nd cut by fractional distillation but was only mildly successful. My final product is 65 mL at 92.5% acetic acid, which I plan to use for an esterification. I've run out of ideas on how to get it any richer.

not_important - 8-10-2006 at 00:42

Could try dehydration with neutral salts, silica gel, molecular sieves, and so on. Also, I believe that fused (anhydrous) sodium acetate will grab onto the water in the acid. But 90+ percent is fine for most esterfications.

When you distill a mixed solution of the lower alphatic carboxylic acids, a mixture comes over. It will be enriched in some of the acids, but by no means pure ly one acid. But for food use, a small amount of those acids besides acetic, and the other organics, just make a slightly richer taste. It's only we fussy chemists how are bother by it.

If you make sodium or calcium acetate as a step in the process, extracting the dry salt with on orgaic solvent might remove some of the larger organics. As an example, boiling down sodium acetate from vinegar and sodium hydroxide/carbonate/bicarbonate will usually give an off-white to tan product. Stirring thiis with "100%" isopropyl alcohol removes much of the coloured material.

16MillionEyes - 21-5-2007 at 10:28

I see most of the effort has been done through distillation. Has anyone tried using god ol' home-plausible process of freezing the vinegar and then collect some of this liquid as it thaws?
I've done it and a weekend with an irritated respiratory tract proved me that this easy process gives off concentrated acetic acid. The only problem is, how can I accurately measure the concentration given based on the rate at which acetic acid thaws compared to that of water?

DerAlte - 28-8-2007 at 21:37

Summary of above thread

(1)The mysterious azeotrope of the acetic acid/water system is as elusive as the blue-arsed fly; it does not exist. The system is zeotropic.
(2)You can increase the % acetic acid in a solution by distilling and get 55% or better with effort (several distillations)
(3) "Pure" or glacial: good luck. Try freezing a well concentrated solution?

For reference, graphs of vapor (partial) pressure vs. temp are attached to help (all data plotted from CRC 1995-1996 ed.)

Lower trace, acetic acid, middle water, top sum of the two. 101 Kpa ~ 1atm. =760 mm Hg. The top trace is not highly meaningful since the presence of the one solvent in the other reduces the vapor pressure of that component.

The fact that the system is zeotropic tells us that the slope of the boiling temperature of a mixture plotted against composition (0-100% acid) has no maximum nor minimum but must have a positive slope throughout. Raoult’s Law never holds except for low dilution factors but assuming a straight line tells you very roughly where a given composition boils.

This graph shows the ratios of the partial pressures. Remember that these graphs refer to the vapor in contact with the pure liquid only. But they may help.



[Edited on 28-8-2007 by DerAlte]

[Edited on 29-8-2007 by DerAlte]

unionised - 29-8-2007 at 10:39

is a whole lot more common than any acetic acid/ water azeotrope.
I think you missed out the use of involatile salts- like calcium or sodium acetate and sulphuric acid as a method too. It's a whole lot easier to remove water from sodium acetate than from acetic acid.
Otherwise a useful summary of a remarkably long thread.

DerAlte - 29-8-2007 at 11:21

@ unionised - LOL! Yup, missed that bit. Better not make an already long thread longer! Regards, Der Alte

merlic79 - 22-8-2008 at 19:25

Why don't you guys just use phosphoric acid? It is readily available in 85% concentrations. As well it is food grade. I mixed it with some dried NaAc and then did a simple VERY pungent. I plan to take this acid and distill with H2SO4 to get pure GAA.

S.C. Wack - 5-7-2009 at 15:03

Vacuum crystallization:

Note the water remaining when vacuum is not used.

starman - 5-7-2009 at 15:38

Quote: Originally posted by S.C. Wack  
Vacuum crystallization:

Note the water remaining when vacuum is not used.

Gee,simply vacuum stripping the product really.I wouldn't have thought such common technique patentable.I suppose the confined temperature range got it over the line.

entropy51 - 5-7-2009 at 17:02

I recently tried the procedure in this patent, starting with about 90% acetic acid and had no luck with it.

I'd be very excited to hear that someone else succeeded!:D

freeze crystallisation of acetic acid

BenZeen - 18-3-2010 at 02:19

Ok, i have used the FSE and couldn't find much so here goes...
The freezing point of vinegar is -2°c at a conc of ~5% acetic acid, and I have seen a photo thread on another website of a person using freeze crystallisation to concentrate acetic acid from wine vinegar ( and I wonder if it would it be possible to use this technique to concentrate acetic acid from white vinegar (about 5% conc), or if anyone has done it already? Thanks;)

Magpie - 18-3-2010 at 08:25

S.C. Wack - 26-10-2013 at 16:22

This attachment is from 1918 and the serial book Vinegar Bulletin, by Paul Hassack. It's the interesting part of THE CONCENTRATION OF SPIRIT VINEGAR, and OTHER CONCENTRATION METHODS OF FERMENTED VINEGAR pp. 210-211.


Attachment: Vinegar_Bulletin.pdf (110kB)
This file has been downloaded 749 times

[Edited on 27-10-2013 by S.C. Wack]

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