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Author: Subject: Acetic acid/ sodium hydroxide
Icarus
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[*] posted on 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.
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Magpie
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[*] posted on 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.




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[*] posted on 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]
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neutrino
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[*] posted on 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 .
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[*] posted on 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]
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Magpie
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[*] posted on 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.



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[*] posted on 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. :)




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neutrino
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[*] posted on 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.
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[*] posted on 8-2-2005 at 19:25


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



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[*] posted on 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.
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[*] posted on 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?




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[*] posted on 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]




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[*] posted on 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..
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[*] posted on 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]




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[*] posted on 10-2-2005 at 20:08
titration results


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.




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[*] posted on 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.
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[*] posted on 11-2-2005 at 03:07
aa


Hows about dry distillation of an acetate salt with sodium bisulfate.
Has this been mentioned?
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[*] posted on 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?




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[*] posted on 11-2-2005 at 15:10
sorry


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.
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[*] posted on 11-2-2005 at 16:01


Quote:
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.
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[*] posted on 11-2-2005 at 16:39
Realy?


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?
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[*] posted on 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.
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[*] posted on 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]
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JohnWW
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[*] posted on 11-2-2005 at 19:09


Quote:
(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 http://journeytoforever.org/biodiesel_ashlye.html , http://www.brainydictionary.com/words/ly/lye186463.html , http://www.suncitysoap.com/oldsoap.html , http://www.octavia.net/9thclife/Lye.htm , http://www.roostx.com/lye_soap/ , 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.
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[*] posted on 11-2-2005 at 19:25
Wrong again?


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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