Sciencemadness Discussion Board

Acid Base Chemistry (H2SO4 + Na2OOCCH3)

rrkss - 25-12-2009 at 19:52

I'm in the process of preparing Glacial Acetic Acid from Sodium Acetate and Sulfuric Acid and just want to make sure I did my arithmetic right.

My goal is to obtain 50 mL of close to glacial acetic acid.

Glacial Acetic Acid has a molarity of 17.4 M so 50 mL will have 0.87 mol of Acid in it.

The equation of the reaction is as follows:

H2SO4 + 2NaOOCCH3 --> 2HOOCCH3 + Na2SO4

since its a 2 to 1 ratio of Acetic Acid/H2SO4 I will need to add 0.435 mol of H2SO4 to 0.87 mol of Sodium Acetate.

98% H2SO4 has a molarity of 18.4 so according to my math, I need to add approximately 24 mL of sulfuric acid (98%) to make 50 mL of glacial acetic acid.

Then after distilling I use anhydrous copper sulfate to dry the acid.

Is my thought process correct or am I missing something. I appreciate the effort.

bbartlog - 25-12-2009 at 20:57

Your math looks right to me, but it seems like it would be easy enough to just use a slight excess of sulfuric acid so as to make sure you've turned all the sodium acetate into acetic acid.
There's also a whole thread devoted to the preparation of glacial acetic acid, which would be a good thing to look at just to be aware of potential pitfalls.




sonogashira - 26-12-2009 at 03:01

This is from page 66 of Norris - Experimental Organic Chemistry (from forum library I think...)
Most preparations like this are covered in old books so no-need to start with calculating moles etc.; unless it is your interest of cause!

89. Preparation of Glacial Acetic Acid (SECTION 103).—Melt
cautiously in an iron dish about 50 grams of anhydrous sodium
acetate. Grind the salt to a coarse powder in a mortar; weigh
40 grams of the salt and place it in a 200-cc. distilling flask. Add
cautiously through a funnel, keeping the flask cold by immersion
in water, 25 cc. of concentrated sulphuric acid. Place a thermometer
in the flask in order to determine the temperature of
the vapor. Connect with a condenser and receiver, and distill
off the acetic acid. Weigh the acid obtained. Calculate the
theoretical amount which can be obtained from 40 grams of
sodium acetate, and the percentage yield of the experiment.
CAUTION.—Glacial acetic acid causes painful blisters when left
in contact with the skin.

rrkss - 26-12-2009 at 08:44

Thanks both of you!

entropy51 - 26-12-2009 at 09:18

I use about 1.2 mole H2SO4 per mole of anhydrous NaOAc, which is a little larger than Norris' ratios.

Don't expect 100% yield. This is one of those reactions that goes better on paper than in the flask.:(

Be sure to dry the NaOAc if you are starting with the trihydrate.

Do look at some of the other threads on this topic. Panache has posted his methods, which seem to yield glacial acid.

rrkss - 27-12-2009 at 08:45

I'm gonna try that. My plan was to boil down a liter of vinegar after hitting it with baking soda and then leaving the stove on for an hour on low to dry out the acetate salt. Putting that in a round bottom and reacting it with H2SO4 and distilling off the product.

rrkss - 28-12-2009 at 19:38

Well I finished the reaction. It took me 42 mL of concentrated sulfuric acid to completely react with my sodium acetate. After the distillation, the reaction flask was completely charred and full of what essentially was soot. I was not able to remove all the charring by washing. I assume that this was from the impurities in the distilled vinegar.

I've just finished adding anhydrous copper sulphate which turned blue upon addition telling me that I have water in my product. Now I have to wait for it to dry on the copper sulphate overnight, followed by me pipetting the product to a new dry round bottom and distilling. Hopefully that gives me the pure glacial acetic acid I seek.

entropy51 - 29-12-2009 at 09:23

Quote:
After the distillation, the reaction flask was completely charred and full of what essentially was soot. I was not able to remove all the charring by washing. I assume that this was from the impurities in the distilled vinegar.
The charring can be avoided by

(1) Let the NaOAc crystallize out of the neutralized vinegar. Don't boil it to dryness or the impurities stay with the NaOAc. Filter off the trihydrate and dry it without burning, which requires care. Panache says microwaving is the way to go, but I've not tried that yet.

(2) Use enough H2SO4 to wet the NaOAc and don't overheat it while distilling. You can even distill under vacuum at a lower temperature. The residue left in my distillation flask dissolves in hot water, no scrubbing needed.

rrkss - 29-12-2009 at 20:38

Thanks for the tips.

After all my work, my acetic acid is freezing at 9 degrees celsius. Anybody know how much water is still in my sample because this is less than the 16 degrees of glacial.

entropy51 - 30-12-2009 at 08:24

Quote: Originally posted by rrkss  
Thanks for the tips.

After all my work, my acetic acid is freezing at 9 degrees celsius. Anybody know how much water is still in my sample because this is less than the 16 degrees of glacial.


Please Google before asking! You would have found this which says 8 C is 95% and 95% is typically what my stuff comes up as by titration. Personally I don't find solid drying agents up to the task of taking out the last few percent of H2O. I recommend titrating it instead of trying to measure a freezing point.

rrkss - 30-12-2009 at 12:24

Thanks for the link, could not find the info via google but probably did not phrase it correctly. Its strange that I get around 95% acid when I used 98% H2SO4 and anhydrous Acetate Salt. Could be moisture from the atmosphere or maybe even residual moisture the acetates which were not removed.

entropy51 - 30-12-2009 at 12:44

Quote:
Its strange that I get around 95% acid


Looking back at my notes, I was getting 93 to 96% when I made my NaOAc using vinegar + Na2CO3 and evaporating to dryness. Then I switched to NaOH + vinegar and starting crystallizing the NaOAc, getting nice white crystals which I then dried first in a vacuum dessicator then by carefully melting and heating until it solidified again. Using this process, my concentration rose to 99% by titration; that's the concentration of the crude distillate, which usually had a BP over 115 C.

My half-arsed theory is that hot H2SO4 oxidizes the impurities in the NaOAc to H2O + CO2, but I can't prove it.

Panache has given a prep in another thread in which he uses NaOAc + NaHSO4, with enough H2SO4 to act as solvent IIRC. He says he gets glacial AcOH. I plan to try that soon. If I were you, I would purify the NaOAc and try Panache's method if you need stronger acid. Depending on your intended use, you may not need 99%.

rrkss - 30-12-2009 at 13:23

Thanks for the very detailed response. Your theory makes sense.

I am using the acid to prepare ethyl acetate. Since the fisher esterification reaction produces ester + H2O from acid + ethanol in a 1 to 1 ratio, I wanted to minimize the amount of H2O in the reactants as possible to push the equilibrium as far towards the products as possible. Don't need absolute glacial Acetic Acid but the less water I have, the better.

Same applies with the Ethanol but I've got pretty much Anhydrous Ethanol so my concern was the acid. Did not want to pay hazmat fees to buy the little bit of Glacial Acetic Acid I was going to be using since I don't normally need the Acetic Acid.

entropy51 - 30-12-2009 at 13:41

Quote:
I am using the acid to prepare ethyl acetate. Since the fisher esterification reaction produces ester + H2O from acid + ethanol in a 1 to 1 ratio, I wanted to minimize the amount of H2O in the reactants as possible to push the equilibrium as far towards the products as possible. Don't need absolute glacial Acetic Acid but the less water I have, the better.
Good Lord No! You can make EtOAc in aqueous reactions! S.C. Wack posted the prep from Norris' organic lab manual (IIRC) on this very forum. (search for it) I think Norris is in the forum library. What you have is fine for making EtOAc. Go for it!

Don't ask me why all that "driving the equilibrium" stuff is not strictly true for EtOAc, but it ain't. Maybe Nicodem knows.

BTW I make EtOAc by adding 50% excess EtOH to NaOAc, adding enough water to make it stirrable, then adding H2SO4. It makes AcOH in situ. Reflux for a while and distill off the EtOAc. Extract repeatedly with saturated CaCl2 solution, dry on MgSO4 and then distill using a packed column.

rrkss - 30-12-2009 at 14:52

Wow, I wish I had known that sooner, took me a hell of a lot of work to make acetic acid.

One question, why do you use the alcohol in excess instead of the acid in excess. Wouldn't it be easier to destroy the excess acid with bicarbonate solution forming insoluble acetate salt (insoluble in the ester) than to have to wash with CaCl2 to get rid of the alcohol. I guess either way you end up with the same result just curious to why you picked the ethanol as the reagent of excess.

One more question, why would you fractionally distill after drying with MgSO4. Wouldn't a simple distillation work?

Thanks for the help.

entropy51 - 30-12-2009 at 15:03

IIRC, the excess of EtOH is from Norris. I think in either case you have some EtOH in the distillate since there is a ternary azeotrope EtOAC/EtOH/H2O.

I use a packed column because some azeotrope comes over before the pure EtOAc. A lot of the older references recommend drying on K2CO3, but that didn't work for me.

EtOAc is easy to make, somewhat more difficult to purify if you need it sans EtOH and H2O. Using strong AcOH is probably better, but I think your 95% should be fine. I don't claim to have a foolproof procedure, by the way, but I usually end up with a decent yield of pretty good EtOAc. I don't know if it's dry enough to react with Na metal, if you're headed in that direction.

rrkss - 3-1-2010 at 09:49

Well I mixed 33 mL of acetic acid with 23 mL of Ethanol, added 0.4 mL of 98% H2SO4 and got a 43% yield of product after 2 hours of reflux. My yield was surprisingly low but I assume it is because I did not reflux long enough or maybe I lost some product when I was washing with bicarbonate solution of nuetralize the acid (EtOAc's soluble in water to an extent).

entropy51 - 3-1-2010 at 10:55

The books usually say 90 some %, but mine runs closer to 70%. It is easy to lose during the workup. Easy to make, hard to purify and dry.

sonogashira - 7-1-2010 at 03:21

^Maybe it could help to add salt to the water-bicarbonate wash do you think? It may help to dry the ethyl acetate also if there is enough?

rrkss - 7-1-2010 at 23:44

Well saturating with salt will definately increase the polarity of the aqueous layer and thus reduce the solubility of the ethyl acetate. Might be a good way to increase my yield by reducing product loss during the workup. Thanks for the advice.

One question, for entropy. What solvent do you use to recrystalize your NaOAc in? I need more Acetic Acid and want it as anhydrous as I can get so I need to remove most of the organic junk from my NaOAc from vinegar.

[Edited on 8-1-2010 by rrkss]

sonogashira - 10-1-2010 at 03:30

Can't you get it on ebay? (Normally it is sold either trihydrate or anhydrous - usually for "hot ice"). Much better if you want larger amount I would think? And you can just heat to remove the 'trihydrate' part. :)

rrkss - 13-1-2010 at 18:44

Entropy sent me a wonderful method which he uses. I just recrystalize the product and wash it with cold distilled water followed by vacuum filtration. This gives me product in the purity I need without the ebay shipping charges. I can make 200 grams of this stuff with about an hour's work.

rrkss - 11-2-2010 at 12:02

Well did another synthesis of acetic acid using this method. This time I used anhydrous NaOAc that I purchased online. It made acetic acid but the solution reaked of Sulfur Dioxide. Something that did not happen with my homemade stuff.

Upon addition of CuSO4, What appeared to be copper metal precipitated out and the entire solution turned a yellow brown color from the previous colorless. The SO2 smell was so strong that I did not even smell acetic acid this time. After vacuum filtration the strong acetic acid smell was there. Now its time for distillation to turn the pale blue solution into clear acetic acid.