Sciencemadness Discussion Board

oxidation of diols

Magpie - 18-9-2009 at 20:26

oxidation of ethylene glycol to oxalic acid

HO-CH2CH2-OH + 8HNO3 --> 8NO2 + HOOC-COOH + 6H2O

Preliminary
Yesterday I recovered about 75 mL of ethylene glycol from antifreeze by fractional distillation. Column packing was a ss scrub pad. This antifreeze contained, besides EG, diethylene glycol, water, dipotassium phosphate, corrosion inhibiters, defoamer, silicates, and dye.

The distillation went very well with a little water coming off first right at 100C, then the EG came off cleanly at 196-197C. I started with 125 mL and left about 25 mL in the pot. It's not like I have any shortage of antifreeze.

Oxalic Acid Procedure
The following procedure is based on "Oxidation with Nitric Acid of Aliphatic Alcohols and Diols to Carboxylic Acids" from the Russian Journal of Organic Chemistry, 2006. My thanks to forum member sonogashira for this paper.

Today 5.6 mL (0.1 mole) of EG was placed in a beaker and 33 mL (0.5 mole) of 68% nitric acid (the azeotrope) was added and stirred. Temperature only rose to 30C from room temperature. This was then placed in a 100 mL RBF suspended in a silicone oil bath under temperature control at 88C. For some minutes nothing was happening although I could see a little brown coloration, indicating production of NO2. My PID controller wasn't properly tuned and the temperature rose to 97C. Still no action. Then, all of a sudden there was a runaway with copious amounts of NO2 generated. I quickly turned on the hood fan which immediately vacated the brown fumes filling the hood space.
Things did settle down without any boilover. The reactants were kept stirred in the bath for a total of 4.5 hours.

The products were set aside to cool to room temperature whereupon crystals (needles) of putative oxalic acid formed. These were Buchner funnel dried with no wash. After air drying 2.3g of crystals were obtained for a yield of 25%, assuming the crystals are not hydrated.

Discussion
This synthesis was done only to determine the viability of the Russian procedure for making dicarboxylic acids from diols. The expected yield was 69%, so that is a little disappointing. But there was no char, and I did use 68% nitric acid vs the prescribed 80% nitric acid. There may be better procedures for making dicarboxylic acids but the viability of this one is established for me. The workup is extremely simple.

oxalic acid prep.jpg - 93kB

[Edited on 19-9-2009 by Magpie]

oxalic acid crystals.jpg - 52kB

[Edited on 19-9-2009 by Magpie]

[Edited on 19-9-2009 by Magpie]

not_important - 18-9-2009 at 20:41

Have you considered the "nickle peroxide" oxidation using a nickle salt with NaOCl? It might not work well with ethylene glycol/oxalic acid, but for longer terminal diols I think it would be viable. The workup might be a little more complex, depending on the water solubility of the acid and its sodium salt, but still shouldn't be too bad.

As for your experiment, indeed you may have lost some oxalic acid by solution in the more dilute HNO3. Isolating the product as calcium oxalate might have been a better indicator of yield.

Magpie - 18-9-2009 at 20:50

Quote:

Have you considered the "nickle peroxide" oxidation using a nickle salt with NaOCl?


No, I've never heard of that method. The only oxidants that I've seen commonly used for oxidation of alcohols are KMnO4, K2Cr2O7/H2SO4, and nitric acid. Potassium persulfate probably also would work.

The reason I chose nitric acid to begin with is that I found some procedures in Cumming for diols.

Looking harder at my 1st post I see that theoretically the mole ratio of nitric acid to EG is 8:1. But I only charged at 5:1. I just mimicked the Russian procedure.

[Edited on 19-9-2009 by Magpie]

kclo4 - 18-9-2009 at 21:24

Same reaction, might be of interest: http://www.thevespiary.org/Douchermann/orgysynpages/oxalicac...

not_important - 18-9-2009 at 23:52

Well, I can't find the thread it was discussed in, so here's one of the articles attached. Gives a 68% yield of adipic acid from the corresponding diol. It chops 1,2-diols, opening rings or giving the two acids from the fragmented chain.

Also see this PDF http://www.sciencemadness.org/talk/files.php?pid=93387&a...
for using preformed "nickle peroxide" for similar purposes. It can oxidise benzyl alcohol in benzene to the aldehyde with yields in the 80-90 percent range. It seems to be a little milder and easier to control/stop than NaOCl+Ni.

There's also some literature on using electrolytic oxidation at nickle oxide electrodes.


Attachment: Nickel-hypochlorite oxidizer.pdf (186kB)
This file has been downloaded 2997 times


behemoth - 19-9-2009 at 04:14

Quote: Originally posted by Magpie  

The following procedure is based on "Oxidation with Nitric Acid of Aliphatic Alcohols and Diols to Carboxylic Acids" from the Russian Journal of Organic Chemistry, 2006.


I have found it as:

"Oxidation with nitric acid of aliphatic alcohols and diols to carboxylic acids"

Authors: Svetlakov, N.; Nikitin, V.; Nikolaeva, E.

Source: Russian Journal of Organic Chemistry, Volume 43, Number 5, May 2007 , pp. 773-774(2)

Publisher: MAIK Nauka/Interperiodica

Would someone post this reference please?

Magpie - 19-9-2009 at 07:19

This paper was recently posted in the references thread. If you don't have access to this ask a moderator for same.

Rich_Insane - 19-9-2009 at 07:45

Forgive me if I'm wrong, but wouldn't the use of dichromate via chromic acid work to oxidize diols? As far as I know, alcohols can be oxidized to carboxylic acids by this method.

Magpie - 19-9-2009 at 07:52


This has already been mentioned. ;)

Quote:

The only oxidants that I've seen commonly used for oxidation of alcohols are KMnO4, K2Cr2O7/H2SO4, and nitric acid. Potassium persulfate probably also would work.

S.C. Wack - 19-9-2009 at 09:19

I'd guess that adding the nitric acid dropwise to hot glycol would be a better procedure than adding all of it and then heating.

Alkaline permanganate works for oxalic acid, but is obviously slightly less simple in the isolation.

entropy51 - 19-9-2009 at 09:21

Magpie, there is a procedure for the HNO3 oxidation of glycol in the June 1939 edition of J Chem Ed, page 285, "Some Improved Procedures for the Elementary Organic Laboratory" by Keith M. Seymour. The file is attached.

This prep uses 6N HNO3 instead of concentrated and says the diluted acid gives a higher yield. It also says that a runaway is essentially to be expected. He uses 200g 6N HNO3 for 12.4 g of glycol.

It reports a yield of 55 - 67%, sometimes as high as 80%. I'm fairly certain that your crystals are the hydrate.

[Edited on 19-9-2009 by entropy51]

Attachment: organic lab improvements 1.pdf (172kB)
This file has been downloaded 786 times


Magpie - 19-9-2009 at 09:31

Thank you entropy, I can likely find that journal at my local library.

I agree that my crystals are probably oxalic acid dihydrate. I'm just too lazy to dry them and determine a mp. That would be the safe thing to do, however.

It's interesting that they claim 6N HNO3 is best. The Russians make the same claim for 75-80% HNO3. :D

Edit: Thanks for the reference link. I didn't know that they did anything serious at Reed College. :o

[Edited on 20-9-2009 by Magpie]

[Edited on 20-9-2009 by Magpie]

Magpie - 24-9-2009 at 08:18

Yesterday I began experimentation with the preparation of malonic acid from 1,3-propanediol using nitric acid as oxidant. I mixed 0.2 mole (14.4mL) of PD with 1 mole of 6N nitric acid (169 mL) in a 500mL RBF with reflux condenser attached. I turned the hood fan on and pulled down the window. I began gentle heating with a bunsen burner until the reaction took hold. What followed was a viscious runaway, with reactants (hot nitric acid) spraying out the top of the reflux column. All I could do is duck back around the corner and let it go. What saved me from burns is the hood window.

This is, of course, both embarassing and sobering. I've experience oxidizing alcohols and should have been more cautious. I think the relatively moderate reaction with ethylene glycol had lulled me into a false sense of security.

S.C. Wack says:
Quote:

I'd guess that adding the nitric acid dropwise to hot glycol would be a better procedure than adding all of it and then heating.


This is the right approach when doing this type of reaction on an experimental basis. After hosing down everything in the hood with water I prepared to try it again. This time I would go half scale and add the diol with a pressure equalizing funnel.

I first added 3 drops of diol to the acid before any heating. Then I gently heated until that small amount of diol began reacting. Then removed the heat and added the diol 5 drops at a time as the reaction subsided. This kept everything under good control. After all the diol was added the RBF was placed in a silicone oil bath set at 105C and heated for another 1/2 hour.

I can see now why only a 5:1 mole ratio of nitric acid to diol is needed. The nitric is mostly reduced to NO, not NO2. This is most graphically seen at the top of the reflux column where the clear gas rising in the column is converted to NO2 as it contacts the O2 in the atmosphere. So, this is probably a more correct stoichiometry:

3HO-(CH2)3-OH + 8HNO3-->3HOOC-CH2-COOH + 8NO + 10H2O

I now have the putative malonic acid products sitting in an evaporating dish. I don't want to go through an ether extraction. For one thing I don't see how it can be very effective when the solubility of malonic acid in water is 1g/0.65 mL water. I am planning on doing another batch where I will convert the malonic acid to diethyl malonate. This may be the most effective way to recover the acid itself, also.

Any suggestions for workup of the malonic acid are welcomed.

malonic acid rx.jpg - 85kB

[Edited on 24-9-2009 by Magpie]

not_important - 24-9-2009 at 19:19

Note that the Org Syn entry does an ether extraction from a concentrated solution with a lot of CaCl2 in it, and they use a continuous extractor.

As for workup, unfortunately the best workup seems to be along the lines of the O.S. route. As you've a solution of the free acid, treat the solution with saturated solution of calcium acetate, filter off the calcium malonate and wash it, then follow O.S. for the free acid.

Alternatively just evaporate the solution under reduced pressure to remove excess HNO3, perhaps redissolve in water and repeat the evaporation.


Magpie - 24-9-2009 at 19:31

re: not_important

Thanks for the good suggestions.

Quote:

Alternatively just evaporate the solution under reduced pressure to remove excess HNO3...


I just finished doing this using a steam bath, but no reduced pressure. As the few mL left is cooling some crystals are starting to form. During the evaporation a lot of small bubbles were generated along with a fair amount of NOx. I hope those bubbles were NOx and not CO2. At 135C this is supposed to happen:

HOOC-CH2-COOH ---> CH3COOH + CO2

How important do you think it is to evaporate at reduced pressure?

not_important - 24-9-2009 at 22:07

Quote: Originally posted by Magpie  
re: not_important

... During the evaporation a lot of small bubbles were generated along with a fair amount of NOx. I hope those bubbles were NOx and not CO2. At 135C this is supposed to happen:

HOOC-CH2-COOH ---> CH3COOH + CO2

How important do you think it is to evaporate at reduced pressure?


The azeotrope with a concentration of 68% HNO3 has a boiling temperature of 120.5 °C

Check out the decarboxylation mechanism, it's promoted by acids. Guess your yield will give a clue if it's important or not.


Magpie - 25-9-2009 at 13:23

I decided to keep my malonic acid product on the steam bath until no more NO2 evolved. As the liquid volume decreased more NO2 evolved. I added a few mLs of water twice. After that I just heated it until NO2 stopped evolving and it was nearly dry. At this point it was rehydrating as fast from the steam bath vapors as it was losing liquid from the heating. I then placed it in my drying oven, set at 70C, for 2 hours. I then weighed it and took a melting point. If the mp was not satisfactory I was going to try washing it with very cold DCM. But the mp, or decomposing point, was 137C vs my literature value of 130-135C. The weight is 4.8 g for a yield of 46%.

I'm very satisfied with this method for making malonic acid, especially in comparison to the classical method that uses KCN. Further procedure refinement would no doubt result in increased yield.

Thanks to Pantone159 for the idea of using 1,3-propanediol as precursor, way back when. Thanks also to entropy51 for the reference from J. Chem. Educ.

malonic acid crystals.jpg - 66kB

[Edited on 25-9-2009 by Magpie]

entropy51 - 25-9-2009 at 13:45

Quote:
But the mp, or decomposing point, was 137C vs my literature value of 130-135C

Magpie, is there any possibility that you heated your sample too rapidly and overshot the MP?

Magpie - 25-9-2009 at 14:07

I don't think so. I came up on it fairly slowly. My 1968-9 CRC gives a mp of 135.6C.

pantone159 - 25-9-2009 at 17:45

Nice to hear that this worked (and that you didn't get hurt!)

I have a reference that describes some purification procedures for malonic acid, if anyone is interested, it was for a situation where extremely high purity was required. (The BZ reaction is demanding.)

Following from:
Noszticzius, McCormick, Swinney. Effect of Trace impurities
on a bifurcation structure in the Belousov-Zhabotinskii
reaction and preparation of high-purity malonic acid.
Journal of Physical Chemistry, Vol 91, No 19, 1987, p5129-5134.

Quote:

Methods for Purification of Malonic Acid.

Recrystallization from Acetone-Chloroform Mixture.
First 100g of malonic acid was dissolved in 115 mL of acetone
on a hot plate, and the hot solution was filtered and
cooled to 20-25 C in an ice bath. Then 175 ml of chloroform
was added with continuous stirring and cooling. When the
temperature dropped again to 20-25 C, the crystals were
filtered amd dried at laboratory temperature.
Yield: 85 g.

Recrystallization from Ethyl Acetate.
This procedure is recommended as a second intermediate step
only for malonic acid samples already purified by the
acetone-chloroform method. Also, after a recrystallization
from ethyl acetate, its traces have to be removed by a
final recrystallization from acetone-chloroform.
First 100 g of malonic acid was dissolved in 300 ml of
ethyl acetate on a hot plate. Then the hot solution was
filtered and cooled in an ice bath to 20-25 C, and the
resulting crystals were filtered and dried at laboratory
temperature. Yield: 65 g.

Recrystallization from Acetone-Nitric Acid Mixture.
This procedure is recommended as an alternative second
step only after recrystallization from an acetone-chloroform
mixture. Unpurified malonic acid samples may contain
contaminants which can react with nitric acid. First 20 g
of malonic acid was dissolved in 40 ml of concentrated
HNO3 at 40 C. Then another 80 g of malonic acid was
dissolved in 95 ml of acetone on a hot plate. Both solutions
were filtered. The acetone solution was cooled to 60 C
and 10 ml of the nitric acid solution was added with
continuous stirring and cooling in an ice bath. When the
temperature dropped to 40 C, another 20 ml portion of the
nitric acid solution was added, and at 25 C all the
remaining amount was added. By adding NaCl to the ice bath
the solution was cooled further to about -15 C. Then the
crystals were filtered and dried under the hood at laboratory
temperature. Yield: 85 g.


Magpie - 25-9-2009 at 20:23

I also thank BromicAcid. Apparently he was the first to mention the use of 1,3-propanediol for making malonic acid.

I certainly miss his contributions and prolific experimental work. But I have a feeling he'll be back...someday. ;)

[Edited on 26-9-2009 by Magpie]

UnintentionalChaos - 25-9-2009 at 20:32

Quote: Originally posted by Magpie  
I also thank BromicAcid. Apparently he was the first to mention the use of 1,3-propanediol for making malonic acid.

I certainly miss his contributions and prolific experimental work. But I have a feeling he'll be back...someday. ;)

[Edited on 26-9-2009 by Magpie]


Where did you get the 1,3-propanediol anyway? This is a nice prep and I'd be interested in trying it.

Sedit - 25-9-2009 at 20:33

What he get 18 to life?
LMAO sorry I couldn't resist. Its friday and late. I sure hope Im not correct because that wouldn't be cool at all and make my crappy jokes just ignorant. I always liked bromics post though:D.

[Edited on 26-9-2009 by Sedit]

Magpie - 26-9-2009 at 06:35


Quote:

Where did you get the 1,3-propanediol anyway?


I aquired this through a private source. I understand it is (or was) also available from ChemSavers.

[Edited on 26-9-2009 by Magpie]

Sedit - 19-6-2010 at 23:04

I see mentioned in this thread the suggestion of adding the HNO3 to the hot glycol, could you explain what advantage this would have? I would think the direct opposite and adding the alcohol to the hot nitric would promot better results since you would keep an excess of oxidant handy at all points in the reaction.

BTW Magpie you suggest that you have never heard of the Ni peroxide oxidation but if you read the paper provided again you will see that Ni peroxide, as well as other methods, where mentions in the first paragraph of the references paper.

I will soon be running this reaction large scale and I scaled up a bit more tonight on EtOH to synthesis AcOH. Since my HNO3 is formed from KNO3 and H2SO4 im considering mixing the H2SO4 plus EtOH and dripping this into a warm KNO3 solution. Small scale experimentation suggest this may be a clear winner but I still have alot more work to do and since this reaction always seems on the brink of runaway it worrys me scaling up.


You will noticethat the shorter the chain on the R-OH the higher the yeilds so I think I will be able to make GAA in decent yeilds pretty cheep since I confirmed glacial status on a test tube run this morning.

What is the mechanism of this reaction? Unless I over looked it I do not see it suggested in the provided reference.

S.C. Wack - 19-6-2010 at 23:55

Quote: Originally posted by Sedit  
I see mentioned in this thread the suggestion of adding the HNO3 to the hot glycol, could you explain what advantage this would have?


Uh, you know...that part about the runaway? A certain amount of fuming is reasonable but... And maybe although HNO3/V2O5 gives oxalic acid with sugar and cellulose, it seems unlikely to have unlimited resistance to hot nitric acid. I don't know what the actual result would be since I haven't done that, but it's an avenue I'd pursue if I had that runaway.

Sedit - 20-6-2010 at 06:12

Ok I got it now thanks...
After much thought last night, on the semi scaled up version I ran, and using the means of HNO3 generation that I do(H2SO4+KNO3), I have concluded that adding the acid to the alcohol to be the best means as well.

Runaways aside the alcohol precipitates dissolved sulfate generating a slush and copious amounts of NOx fumes more dense then I have ever personally seen. I would like to stir the reaction as it proceeds and the precipitated sulfates would make that all but impossible. Im going to also investigate some electrochemical means of HNO3 synthesis today since that could possibly make this a very cheep route to AcOH if I can find a means to an end.

Sedit - 23-6-2010 at 12:27

I just got a brief bit of time to skim the papers kindly provided by solo and there appears to be a means of adding the Alcohol to the acid while still avoiding runaways. It appears the initiation of the reaction comes from the presences of NO2 ions and the addition of NaNO2 in proper concentrations avoids the induction period before the reaction begins meaning that as the alcohol is stirred into the acid it will be turned into the carboxylic acid avoiding the runaway reaction.

This will allow me to keep a huge excess of oxident present without the prior fears of a runaway. The most interesting part is that the oxidation of EtOH and IpOH both result in the desired product of AcOH.

[Edited on 23-6-2010 by Sedit]

Sedit - 24-6-2010 at 18:28

OK I know im talking to myself here but I would love to hear some input on my new take on this reaction. I know im deviating abit from the thread topic but the reaction remains the same.

Sodium nitrate is soluble in EtOH, something dont have the full solubility data on just yet.

What about adding H2SO4 to a solution of NaNO3 in EtOH. HNO3 is formed on the spot and it could quell the dreaded runaway by generating localized heating leading to "completion" of the reaction before it has a chance to take off. This is been suggested in small scale already but I only have clean KNO3 right now and its solubility in alcohol isn't that great yet its showing success better then other means.

Can anyone forsee problems doing it this way that im missing? Im thinking with stong stirring and ice cooling this should proceed smooth as a babys behind followed with a post reaction heating to push it to total completion.

not_important - 24-6-2010 at 19:09

From

A Dictionary of Chemical Solubilities, Inorganic 2nd Ed. Comey

Available at the Internet Archive and Google Books

(and yes, the calcium nitrate file is just CaNO3... because Ca(NO3)2... has some potential filesystem squibble.


I don't see your reasoning matching what would happen, I think that a runaway is even more probable. Also, given the near anhydrous conditions I'd expect EtONO2 to form.

If nothing else I'd dissolve the H2SO4 in some EtOH first, so as to avoid thermal input from it reacting with the alcohol in the nitrate mix. More heating drives the reaction faster, that's what the runaway is about, the heat of the reactions speeding up the reactions even more.

(remember - babies bottoms may be smooth, but consider the high volume of you-know-what that comes out of them :-)


Personally, I don't see the reason to waste expensive or increasingly difficult to obtain oxidisers on simple lower alcohols; I'm a bit puzzled as to why you've chosen this route.



NaNO3-EtOH.png - 71kB CaNO3-EtOH.png - 17kB NH4NO3-EtOH.png - 76kB

[Edited on 25-6-2010 by not_important]

Lambda-Eyde - 24-6-2010 at 19:38

http://www.ab.ust.hk/hseo/tips/ls/ls005.htm

http://users.wpi.edu/~rajat/MSDS%20Ethanol.pdf (Read section 5)

Sounds like a bad idea to me.

Sedit - 24-6-2010 at 20:16

Quote: Originally posted by not_important  
From
I don't see your reasoning matching what would happen, I think that a runaway is even more probable. Also, given the near anhydrous conditions I'd expect EtONO2 to form.

If nothing else I'd dissolve the H2SO4 in some EtOH first, so as to avoid thermal input from it reacting with the alcohol in the nitrate mix. More heating drives the reaction faster, that's what the runaway is about, the heat of the reactions speeding up the reactions even more.

Personally, I don't see the reason to waste expensive or increasingly difficult to obtain oxidisers on simple lower alcohols; I'm a bit puzzled as to why you've chosen this route.

[Edited on 25-6-2010 by not_important]

Thank you for the solubility data I owe you one.


To answer last question KNO3 and H2SO4 are accessible and cheep for me and given the prices discussed and researched on large amounts of AcOH I am confident I can produce this chemical cheeper then buying it from a supplier perhaps around half price or less.

Dissolution of the H2SO4 into EtOH has crossed my mind as well simular to how I performed the EtBr synthesis sometime back. This is the next step in my small scale before moving on BTW and has been on the checklist for sometime.

Now back to the most important question.

""I don't see your reasoning matching what would happen, I think that a runaway is even more probable. Also, given the near anhydrous conditions I'd expect EtONO2 to form."""

Reasoning is as such....
Small scale experimentation has shown what I have assumed. Heating occures in a local area causing the "runaway" in a very local area of the drip of acid quickly. Stirring cools this back to "baseline" with ease since the higher temperatures are only in a local area. The runaway so to speak has shown considerable larger concentrations of AcOH over acetaldahyde then allowing the reactants to mix then react "slowly" over time.

The Ester IS my main concern since it will result in a large loss of yeild and adding H2O is no issue since I have decided the best workup might be to distill, dry, and extract with DCM. I have had success of extracting (aq) solutions of AcOH so this part is a no brainer for me once the AcOH is synthesized.

I desire this reaction because after years of research this appears to be the cleanest and cheepest means of producing large amounts of GAA economically. The NaOAc + H2SO4 method sucks and produces impure product. This is showing MUCH more potential then that by far.




Lambda, Nitration is much more hazardest then this reaction yet everyone seems to have no issue running that daily. This is not going to sit on the shelf for anytime so the dangers associated are null.

Lambda-Eyde - 24-6-2010 at 20:35

Quote: Originally posted by Sedit  

Lambda, Nitration is much more hazardest then this reaction yet everyone seems to have no issue running that daily. This is not going to sit on the shelf for anytime so the dangers associated are null.


I misinterpreted your post. My bad. My understanding was that you didn't want the nitric acid and the ethanol to react.

not_important - 24-6-2010 at 20:58

When you say
Quote:
The NaOAc + H2SO4 method sucks and produces impure product.
do you mean the vinegar as the source of acetate, or any source of NaAc?


Sedit - 24-6-2010 at 21:05

Vinegar of course. Im sure Pure NaOAc would produce a better product but no matter how clean I try to get it the results still appear tainted. Perhaps its due to decomposition products or poor recrystalization methods(something I pride myself in BTW) but in the end its not a good means to large GAA source. 50-100ml at best is the limit of that reaction before it starts to present issues and im sure there are others here that will confirm this for me.

not_important - 24-6-2010 at 22:07

OK

A) I found that using s slow cooker to evaporate the neutralised vinegar until you've got essentially liquid NaOAc.3H2O, then slowly pouring that into chill i-Pr alcohol with good stirring, gives a pretty clean precipitate of the acetate, partially dehydrated to the monohydrate. Washing the ppt with several small amounts of further chilled IPA removes much more of the colouring; after a single wash with IPA, adding a bit of water the the acetate and going through the melt-evaporate-pour and washing the precipitate twice gave me a white product without much loss (the trihydrate's solubility in EtOH is around 50g/l, in IPA it's lower) A dichloromethane extraction of the powdered solid acetate also might give good results. In either case sugars in the vinegar would remain with the organic solvent insoluble portion, but as they're non-volatile the distillation of the AcOH should take care of them.

I'd also consider the oxidation of EtOH by NaOCl catalysed by nickel salts (which form NiOOH/NiO2). Filtering removes the nickel, evaporation leaves NaCl and NaOAc.3H2O

In either case I believe it is important to avoid localised overheating when adding the strong acid to the NaOAc. Dispersing the acetate in acetic acid and slowing adding H2SO4 seems to work, a similar approach is used for formic acid when the dehydration to carbon monoxide can be a real problem.


Sedit - 28-6-2010 at 06:02

I found that using an old stainless steel deep fryer that I modified for steam generation to be the best means of concentrating the NaOAc since it has ability to boil off gallons of liquid in a matter of minutes instead of hours. I will have to give IpOH a shot if I ever go that means again but even though im a huge kitchen chemist even this seemed a little excessive for 500 grams or so of NaOAc. Ill take electrochemical means that produces NaOAc from EtOH over evaporation of that much liquid again but thats a method of the acetic acid threads.


I had a chance to attempt the HNO3 oxidation at a one more scale of acid which I sort of botched but showed some interesting results to me anyway. I used 1:1 KNO3-H2SO4 with 20ml of H2O to increase the fluidness of the acid when I could have gotten away with 1:0.5 nitrate to acid ratio but was not thinking. However after reading abstract after abstract over the past few days I decided I was going to start with the addition mol of EtOH the paper recommends of .2 mol or something on the order of around 10 grams of EtOH and slowly add more until NO2 stops being released. After getting all the way to 50 grams of EtOH and still showing signes of NO2 I stopped because this does not make much sense to me.

The reaction contains the initial NO2 ions needed for initiation of the reaction from over heating as the H2SO4 is added presumably. Drop by drop I started to add the EtOH to warm acid and stand back to observe. At first I thought stirring would be good but this causes heavy foaming and I found if I just leave it as is it starts a relatively mild cycle that was not hard to control with addition rate at all. I would still take much caution for the first few grams of addition since it does show potential for going out of hand. A cold water bath was used for cooling and I think Ice would be much better since alot of heat is generated at the start of this reaction and with NO2 already present it does not need an initial heat source to get it going.

One drip of EtOH is enough to cause all of the NO2 to vanish instantly from the flask leading me to think the NO2 ester plays an important role in the oxidation of some kind.

Even after breaching 1 mol of EtOH there is still signs of NO2 but much calmer then before and the entire reaction is now fluid with the Potassium bisulfate almost completely dissolved. This would precipitate at the start and I attribute this to the sulfates solubility in AcOH and insolubility in EtOH where as it could also be in solution from H2O formed in the reaction, that is yet to be seen.

The smell of AcOH(Dont sniff the FLASK drip it on blotter paper) is very very strong now after a slow warm water bath that lasted a couple days since I was allowing it to react then add a little more and so on....

Its non flammable by any means at this point.

Why would this reaction still show signs of NO2 generation even after 1 mol of EtOH? I hypothesized that I may be able to do about .5 mol but not 1. Could the acidic nature of KHSO3 be affecting the reaction somehow?

I wish I had a GC or some means of analysing stuff like the big boys do but its hard here down on the bottom. Still betcha I have more fun doing what I do for free then they do for money:D

I obviously have alot of work to do before I could give a definitive yeild. I am going to add a few drops to Sodium bicarbonate and see how wet it stays after its done reacting. NaOAc should be relatively dry where as H2O and EtOH would leave it damp afterwards. And it will allow me to smell for the presence of EtOH, AcO, and any esters that may form. Freezing the reaction mixture would be a good idea as well if it can indeed freeze. I want to have a good idea whats in the mixture before I distill it because all this talk of EtOH and Nitric acid forming explosive mixtures has got me a little uneasy.

Sedit - 28-6-2010 at 18:05

Dripping it onto Sodium bicarbonate produced alot of CO2 some of which could be from HNO2 or HNO3 but in the end when it was neutral there was evidence of alot of presumably H2O remaining, since it contained very little if any of a smell, which is pretty much unacceptable to me. I was unaware this reaction produced a mol of H2O as it proceeded and this is counter productive to what I wanted.

The liquid remaining also had a remaining smell of something I can not explain that smelled sort of like burnt AcOH but I can not quite explain it. Surely not the sweet smell of an ester like I would expect. It was only faint though so whatever it was there is not alot of it or it has a weak reaction to the nose. Going to try to freeze it and see what I get from that test. I hoped the sulfate would suck most H2O out of the reaction but I don't believe it did. I still non the lest did not detect even a hint of EtOH or AcO smell in the remaining liquid.

Still I know AcOH is produced in a decent quantity(Maybe almost quantatively but I will need to remove all other acids before titrating to determine the yeild) and I will try to work it up and try to determine yeilds for the sake of science but in the end I think im going to start attempting some vapor phase reactions to yeild AcOH in the near future but that is another topic for GAA threads and not this one.

blogfast25 - 24-7-2010 at 09:03

Where did all the stuff go???

I tried to repeat Julius B. Cohen's method for oxidising sucrose to oxalic acid with conc. nitric acid and V2O5 but with remarkably little success so far.

I don't have any conc. HNO3, only 38 w% (6N), so had to adapt the recipe. I ended up with a hybrid between Cohen's and Magpie's (first post in this thread), figuring that sucrose could very loosely be described as a hexamer of ethylene glycol.

Cohen calls for 20 g sugar (sucrose), 140 ml conc. HNO3 and 0.1 g of V2O5.

I cooked up 10 g, sugar, 140 ml 38 % HNO3 and 0.05 g of V2O5, by adding the sugar to a hot (steam bath) mixture of the nitric and the catalyst. Based on Magpie's first attempt, this is somewhat sub-stoichiometric in HNO3.

Ready for a run away, instead the solution turned green (V(H2O) +3) first and slowly started fuming NO2. That fuming build up until thick oodles of brown smoke bellowed out of the flask but without any run away: the solution was boiling away nicely.

After an hour or so on the steam bath the reaction started slowing down and the boiling stopped. I transferred the liquid to a Pyrex jug and started to heat with a small direct flame, for about another hour, topping up the liquid with water so that the level didn't fall below 100 ml. At the end the fuming had all but stopped and I reduced the green solution to about 40 ml, cooled and even chilled it but no crystals of HOOC-COOH where forth coming.

The next day the solution was carefully reduced to almost nothing and a reddish precipitate formed: hydrated V2O5? But little else was found. I filtered the solution, the filtrate ran green. But adding concentrated Ca2+ yielded no calcium oxalate...

I tied again, this time 5 g sugar, 115 ml 38 % HNO3 + pinch of V2O5. Assuming sucrose is about (EG)6, this is then stoichiometric in HNO3. The procedure above was used.

This time, after careful and complete evaporation, an amber-reddish crystalline mass was obtained probably less than one gram though. This dissolved effortlessly in water, turning clear green in the process. Added anh. CaCl2, this dissolved effortlessly, without any calcium oxalate forming!

Am I overheating this stuff and oxidising also the oxalic acid, or what is going on here??? :o

One more attempt will be made with 5 g sucrose, 115 ml 38 % HNO3, no V2O5 and cooling on a cold water after the reaction starts...

Paddywhacker - 24-7-2010 at 16:34

If the oxidation is going too far, then maybe more sugar or less acid would do the trick. Also, adding the acid mixture to boiling sugar solution will stop the acid being in excess.

blogfast25 - 25-7-2010 at 09:31

Quote: Originally posted by Paddywhacker  
If the oxidation is going too far, then maybe more sugar or less acid
would do the trick. Also, adding the acid mixture to boiling sugar solution will stop the acid being in excess.


Cohen actually recommends adding the sugar to the acid. I can't really see it make much difference but I could be wrong on that. I believe an excess of nitric is better than not enough but it's more an article of faith than anything else...

Third attempt: this time a mixture of 5 g sugar and 115 ml of 38 % HNO3 (no V2O5) was gradually heated on the steam bath. It took decidedly longer for the reaction to start but once started, it's indistinguishable from a reaction that includes the catalyst. Interesting to note is that the solution also turned green, as in the presence of V2O5.

Taking the bottle off the steam bath and the reaction stopped. I see little point in doing that, so continued with the steam bath for about 2 hours, until NO2 elution had subsided much. That was then left to stand, quasi-stoppered, overnight.

The flask, with camera shake, just after taking off the steam bath. It was then returned to it:



The 100 or so ml of liquid left was transferred to Pyrex receptacle and simmered under partial reflux. The oxidation reaction started up again and continued for about an hour or so, with the solution also picking up a brownish colour. At about 50 ml it was transferred to a smaller Pyrex receptacle and gently boiled in to almost nothing. It dried, showing plenty material, including some black stuff. It seems that right to the end some oxidisable saccharide remains, which then carbonises at the end.

The Pyrex jug, after gentle boiling (with camera shake): the head space is still full of NO2:



After the smoke has cleared: whitish/yellow mass with a black spot. White material can also be seen on the inverted glass lid (right):



To the solid material was then added 10 ml of water: most of it dissolves but not all. It was heated on steam bath for some time, then filtered. The filtrate is clear but brownish. On the filter: a white powder. The powder was isolated. It doesn't dissolve in water but it does dissolve in an excess of strong HCl: an oxalate? It will be neutralised with NH4OH and checked for oxalate.

The filtrate was gently evaporated on steam bath until almost nothing, then left to dry: needle like structures so typical of oxalic acid developed (hard to see in the photo below). Unfortunately the brown muck did not separate out much and the amount of material is very small. Rough yield to be determined yet but it seems to me that the oxidation of ethylene glycol with diluted nitric (Magpie and other sources) works well, but not the oxidation of sucrose for which much stronger nitric is needed for more quantitative yields...



I wonder if improvement could be obtained by carrying out the whole operation on steam bath, from reaction to evaporation of water...


[Edited on 25-7-2010 by blogfast25]

blogfast25 - 29-7-2010 at 07:50

Well, a fourth attempt did work: this time 200 ml HNO3 38 % and 8 g of sugar, a ratio I believe uses a slight excess of HNO3.

Most of the reaction was carried out on steam bath, then paused overnight and finished off by gentle simmering. The last 100 ml or so were evaporated on steam bath. NOx continued to elute up to the last moment.

5 - 6 g of crude oxalic acid, slightly wet and very slightly yellowish:



The crude has been worked up with three precise recrystallisations.

So it's possible to make oxalic acid even from sugar with fairly dilute HNO3. But it will work faster and better with azeotropic or higher HNO3...


un0me2 - 30-7-2010 at 05:30

You can make formaldehyde from it too, at least according to these Authors (they use Lead Dioxide and acetic acid, forming the tetraacetate in-situ:)

Formatik - 30-7-2010 at 20:22

Quote: Originally posted by blogfast25  
The powder was isolated. It doesn't dissolve in water but it does dissolve in an excess of strong HCl: an oxalate? It will be neutralised with NH4OH and checked for oxalate.


In Beilstein it's said that oxalic acid stubbornly retains alkalis, method for purification given there recrystallizes from boiling 10-15% HCl. Then washing with a little cold water and recrystallizing from alcohol.

[Edited on 31-7-2010 by Formatik]

un0me2 - 30-7-2010 at 21:58

Funny, oxalic acid is one of the few I can purchase no questions asked. They use it for a huge number of reasons, look for it via google "oxalic acid msds" normally works for me.

blogfast25 - 31-7-2010 at 13:08

Quote: Originally posted by Formatik  

In Beilstein it's said that oxalic acid stubbornly retains alkalis, method for purification given there recrystallizes from boiling 10-15% HCl. Then washing with a little cold water and recrystallizing from alcohol.

[Edited on 31-7-2010 by Formatik]

Nice tip. But I don't see where the alkali would come from in my case...

Quote: Originally posted by un0me2  
Funny, oxalic acid is one of the few I can purchase no questions asked. They use it for a huge number of reasons, look for it via google "oxalic acid msds" normally works for me.


I actually now found it on eBay...

blogfast25 - 1-8-2010 at 11:49

Having gone a little 'calorimeter mad' of recently I've attempted to measure the approx. enthalpy of reaction for the oxidation of ethanol (a model for mono alcohols) in watery medium with dichromate (K2Cr2O7) and sulphuric acid. We know these oxidation reactions are highly exothermic: KMnO4 and glycerine mixtures for instance heat up enough to ignite a Classic Thermite mix.

I came across a classroom recipe for the demonstration of the oxidation reaction of ethanol with dichromate and adapted it very slightly to:

4 ml of 50 % H2SO4
25 ml of water
3 g of K2Cr2O7

to which after mixing, 4 ml of alcohol is then added. This reagent mixture is non-stoichiometric: molar ratio of alcohol to dichromate being about 6:1, instead of the stoichio 3:1. This should guarantee that all dichromate (about 0.01 mol) is reacted and about 0.03 mol of ethanoic acid is formed.

A dry run showed that despite not all the dichromate dissolving in the H2SO4 solution, the reaction starts promptly and runs to completion with all the dichromate being used up and a hot, non-turbid, deep violet (Cr3+) mixture of excess alcohol, excess acid, ethanoic acid, and potassium and Cr [+III] sulphates results in a matter of seconds. End temperature was about 70 C, a ΔT of about 50 C: for 0.03 mol ethanoic acid formed that's very promising.

I then repeated the experiment au calorimeter with precisely measured quantities, including a precisely weighed amount of dry K2Cr2O7. I need to specify that for 'ethanol', read 'denaturated alcohol', so a mix of mostly ethanol and some methanol was used. For the calculation of reaction heat, it was assumed that all liquids had the heat capacity of water. To make this more realistic, a 100 ml water of precisely known temperature was added after the reaction had completed in the calorimeter and the final enthalpy content was then determined after stabilisation..

Thus an estimated enthalpy of reaction for the oxidation of ethanol with K2Cr2O7 of -245 kJ / mol of acid formed was obtained.

That's quite a whopper for a reaction that involves no lattice energies whatsoever!

Of course this is only one data point and the oxidation heat of an alcohol is likely to depend somewhat on the R group (in R-CH2-OH) as well as the oxidant used but if we assume the oxidation of 1,2 propane diol to malonic acid with nitric acid as attempted here:

http://www.sciencemadness.org/talk/viewthread.php?tid=12832#...

by Magpie to have a comparable enthalpy per mol of oxidised alcohol group, then the oxidation of 0.2 mol 1,2 PD x 2 mol -OH/mol of 1,2 PD x 245/mol of -OH group = 98 kJ, or 98,000 J.

Assume the 169 ml of 6 M nitric to have the same heat capacity as 169 ml of water, then the temperature increase upon full oxidation of the 1,2 PD to malonic acid is expected to be ΔT ≈ 98,000 / (4.1813 x 169) =138 C !! No wonder the experiment resulted in painting Magpie's fume hood a new shade of 'nitric acid'!

So, boys and girls, be careful with these oxidations of diols and triols and make sure you've plenty heat sinks/cooling in place to prevent run aways!

Update - calculation error alert!

Doh! Where I described the stoichoimeteric ratio of alcohol/dichromate to be 3, that should really have been 1.5 as the oxidation of ethanol to ethanoic acid involves 4 e- (and not 2 e- as previously assumed for some reason):

CH3-CH2-OH + 5 H2O ---> CH3-COOH + 4 H3O+ + 4 e-

The estimated enthalpy of reaction for the oxidation of ethanol with K2Cr2O7 then becomes - 490 kJ / mol of acid formed (and not -245) and the estimated temperature rise in Magpie's run away, 276 C.

My bad.


[Edited on 2-8-2010 by blogfast25]

blogfast25 - 2-8-2010 at 08:01

The value of - 490 kJ/mol sounding so high, I decided to replicate the experiment above, this time using KMnO4 as oxidiser. A dry run showed that the reduction goes from Mn [+VII] to Mn [+IV], as brown MnO2 precipitates. The assumed stoichiometry here is:

4 MnO4- + 3 C2H6O (ethanol) + 4 H3O+ ---> 4 MnO2 + 3 C2H4O2 (ethanoic acid) + 9 H2O

The recipe used in calorimeter was aimed at oxidising 0.015 mol alcohol to acid, using a limiting amount of precisely weighed KMnO4 and an excess of alcohol and acid:

4 ml 50 % H2SO4
25 ml water
3.16 g KMnO4

to which after mixing and temperature measurement, 4 ml of alcohol (denaturated) was added. After reaction another 100 ml of water at known temperature was added and final enthalpy content of the flask determined.

This yielded a reaction enthalpy of - 490 kJ/mol of formed acid. Precisely the same value as above. I think this shows the nature of the oxidant plays only a minor role, as long as it's powerful enough, of course...

draculic acid69 - 15-11-2022 at 02:19

Quote: Originally posted by not_important  
OK

A) I found that using s slow cooker to evaporate the neutralised vinegar until you've got essentially liquid NaOAc.3H2O, then slowly pouring that into chill i-Pr alcohol with good stirring, gives a pretty clean precipitate of the acetate, partially dehydrated to the monohydrate. Washing the ppt with several small amounts of further chilled IPA removes much more of the colouring; after a single wash with IPA, adding a bit of water the the acetate and going through the melt-evaporate-pour and washing the precipitate twice gave me a white product without much loss (the trihydrate's solubility in EtOH is around 50g/l, in IPA it's lower) A dichloromethane extraction of the powdered solid acetate also might give good results. In either case sugars in the vinegar would remain with the organic solvent insoluble portion, but as they're non-volatile the distillation of the AcOH should take care of them.

I'd also consider the oxidation of EtOH by NaOCl catalysed by nickel salts (which form NiOOH/NiO2). Filtering removes the nickel, evaporation leaves NaCl and NaOAc.3H2O

In either case I believe it is important to avoid localised overheating when adding the strong acid to the NaOAc. Dispersing the acetate in acetic acid and slowing adding H2SO4 seems to work, a similar approach is used for formic acid when the dehydration to carbon monoxide can be a real problem.



Surely sugars are only present in malt,brown and balsamic vinegars, shouldn't all white vinegar be distilled?