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[*] posted on 20-5-2002 at 18:03
Acetic anhydride preparation


I'm theorizing that concentrated acetic acid could be dehydrated to acetic anhydride by mixing with concentrated sulfuric acid; then heating, and condensing the vapors, yielding acetic anhydride. Any comments or additional ideas?



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[*] posted on 20-5-2002 at 18:04


If the production were this easy, I think it would already be documented somewhere. But maybe not. Have you looked at Rhodium's materials on this topic? All of their methods require chlorides of phosphorus or sulfur, not so nice to work with and certainly not available as consumer products.

www.rhodium.ws/chemistry/anhydrides.html
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[*] posted on 20-5-2002 at 18:04


It turns out that there's another synthesis route to acetic anhydride that requires more equipment but uses more accessible chemicals. Acetone is pyrolyzed with a catalyst in an electrically heated glass tube and the product is dissolved in glacial acetic acid to produce acetic anhydride. Search on the Hive for more info.
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[*] posted on 20-5-2002 at 18:04


Thanks for reminding me of that... I believe that is the ketene process.



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[*] posted on 20-5-2002 at 18:05


I am sorry to dissapoint you, but it is not possible to dehydrate acetic acid to Ac2O.

Some time ago I was also working on the synthesis of this compound, but when I got from a friend about 1,5 L Ac2O, I stopped with that .

I strongly discourage you to do something with the route using ketene. Ketene is very dangerous. The setup is very difficult to realise. The procedure gives very low yields.
The whole process is only used in industry, and you can not always transform such syntheses to lab-scale.

I've did thought out another route using acetaldehyde which could be realised fairly easy. But you do have to make/buy acetaldehyde first for that.

Acetaldehyde can be made by leading acetylene through a solution of a Hg2+ salt and then condensing it.
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[*] posted on 20-5-2002 at 18:06


Not that I've actually done it, but I have good reason to believe that the ketene process CAN be performed on a lab scale. There was a thread on the Hive that included photographs of a home setup that had actually been used as part of drug manufacture in the early 1980s (now beyond statute of limitations). The person who used it was using thorium oxide as a catalyst, per instructions found in (I believe) Organic Syntheses (may have been a journal article; can't recall.) He claimed (IIRC) to be able to produce about a liter of acetic anhydride per day. There was an extended discussion on catalyst preparation and alternatives to the exotic and radioactive thorium catalyst. Certainly, I believe it may be hazardous, but not so hazardous as to disregard entirely. I believe the greater hazard is with potential legal trouble.
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[*] posted on 20-5-2002 at 18:07


No, but only if there are no other alternatives. And there ARE in this case.
I will give some more details on the route via acetaldehyde some other time.

But if you're using it as a nitration catalyst there are several other possibilities to use, but also not very easy to get propably..

Further you must also not underestimate the power of simple mixed acid nitration. The concentration of the nitronium ion is such is mixtures is actually quite large.

Other acid-catalyzed nitrations can be done using mixtures of HNO3 with HF, HClO4, BF3, TFA, TFAA etc.

And of course nitronium salts are also a interesting group of nitration chemicals! NO2BF4, NO2SF6, NO2PF6.

And probably the most powerful nitration mixture is Magic Acid. This is a mixture of FSO3H, SbF5 and HNO3.
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[*] posted on 20-5-2002 at 18:08


es, I remember a thread that I started at The Forum a while back, "Holy Grail Oxidizer"? That included instructions on how to prepare [NF4+][ClO4-]. This was promptly dismissed by nbk2000, who then gave me a lesson on the basic fundamentals of chemistry. *sarcastic tone of voice* "Chemistry isn't just 1+1. Quit talking out of your ass". Perhaps sometime soon I shall repost that, here.

I believe it is rather easy for the home chemist to prepare CH3CHO. I haven't had the chance to test to see if it had actually worked, but recently I attempted to prepare CH3CHO from CH3CH2OH via the dehydrogenation route. Alcohols can be dehydrogenated by passing the hot vapor over hot copper metal, in the absence of reactive gasses such as oxygen. This would be the reaction for the formation of the CH3CHO:

(hot CH3CH2OH vapor passed over hot copper metal in oxygen-free environment)
CH3CH2OH --} CH3CHO + H2

If a small amount of oxygen gas contaminent is present, the results will not be disasterous, because the oxygen will first react with the hydrogen produced.

How I set up the experiment:
I took a 250mL flask, and filled it with 50mL of CH3CH2OH. I placed a rubber stopper in it, and connected a thin copper metal tube to the opening in the stopper (tight fit). I set the flask on a metal screen on top of a propane burner. I then bent the copper tube around so that it would come in close contact with the flames emmitted from the burner. The tube then continued on, bent down so that condensed vapors would drip into another flask (did not use a stopper at that end because then pressure might build inside of the two flasks and metal tubing, causing a stopper to shoot off). I of course had the metal tube, at the condensation end, packed with ice/salt filled baggies. I condensed a fair amount of liquid (don't remember just how much).




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[*] posted on 20-5-2002 at 18:09


Hmm. This is another one of those industrial processes which have been transformed to home experiments.

But I must say that I find this rather stupid: heating with burners when working with the highly flammable and quickly vaporizing acetaldehyde. Even if you try to cover things up a little you can get quite dangerous situations. I have had some nasty experiences with a comparable experiment in the past so that's explains the nature of this reaction of me.

And I also doubt on the effectiveness of this experiment. How long was your copper tube? You'll need quite a long contact area.

BTW, I heard it is also possible to partial reduce ethanol to ethanal with K2Cr2O7. I find this also rather strange, since most simple alcohols get immediately reduced to the acid when you try to do this.
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[*] posted on 20-5-2002 at 18:09


I haven't worried about my experiment, considering that what acetaldehyde I have is either collected in a flask which is in an ice bath; or is being condensed in an environment free of oxygen. I have also not worried because I have been doing this outside (I live in a windy area), and because I am nowhere near the experiment while ethanol is still present in the flask that is being heated. I haven't even been able to ignite the vapors of boiling isopropanone that are jetting out of the neck of a flask (in the windy conditions outside). Please elaborate on your accident. :-)

The part of the copper tube that is being heated is about eight inches long. The length of the entire tube is a number that I am uncertain of, considering its now-curved nature. The acetaldehyde is immediately condensed after it passes through the heated part of the copper tube. The copper tube is around 2-3mm in diameter. I put eight kinks in the part of the copper tube that is being heated; at these kinks the width of the opening is about 1mm. I figured this would cause a lot of turbulence as well as slowing down the passage of the ethanol vapors, causing far more of it to be converted to acetaldehyde.

Even if this experiment failed to convert all of the ethanol to acetaldehyde, I could just run the ethanol / acetaldehyde mix through the experiment again, to dehydrogenate more of the ethanol into acetaldehyde.

As for oxidizing alcohols to aldehydes / ketones, I think it would be worth considering using hypochlorites to reduce an alcohol to its respective ketone / aldehyde. The problem with using permanganates or other oxidizing agents (this is just a theory) is probably that the permangante ion will react with just one alcohol molecule, reducing it fully, instead of reacting with multiple alcohols, reducing them partially, down to their respective ketone / aldehyde. A hypochlorite, however, would not have that problem.

CH3CH2OH + NaOCl --} CH3CHO + NaCl + H2O

The acetaldehyde could then be distilled. Water could be removed by adding the correct amount of MgSO4 (not the hydrate, MgSO4*7H2O!), which over an hour or two will absorb the water. Again, that mixture would then be distilled (but actively monitored to avoid heating the MgSO4*xH2O formed which would cause the water to be released again, defeating the purpose of that process).




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[*] posted on 20-5-2002 at 18:10


Although I agree with you that it's quite certain there is atleast some CH3CHO in your product. But I think you have to realise more that often A LOT of side reactions are also possible that usually aren't listed in the standard literature. This story also goes for example on your theory of those carboxyl amines.

You would be maybe amazed how uncertain professional chemists often are about products obtained via (suspected) quite standard reaction mechanisms. Practically *every* compound that is synthesised is run through IR, NMR, and HPLC.

Our little 'research group' also walks to this problem. We only have IR and TLC capabilities and althoug that's already better then most hobbyists have, we really need for example at least a HPLC and a DSC.

For the possibility of using hypochlorites for that oxidation I also have my doubts. Hypochlorites are used in some other oxidation reactions, yes. But I'm really doubting what that OH group is going to do when ClO- approaches, several other possibilies also possible. And ClO- isn't a very powerful oxidator also. But I gues it would be quite easy to try.

But it's already proven that you can also do it using dichromates you why even bother further. Well, only if you can't get dichromates maybe.

BTW, how did the obtained product smelled?
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[*] posted on 20-5-2002 at 18:12


Hello Pyroweb euh Requiem II! Aren't you related to that guy on the weapon and explosive forum tsv or tvs + number?
I also think to remember you are part of that famous research group ;-) High energetic material & explosives.When I saw pichon, I thought it might be you, but when I saw that someone has provided you 1,5 L Ac2O (since it was me, there is no doubt :-)).
Why 3 different names?

First of, yes CH3-CH=O is easy to do from acidic KMnO4/ K2Cr2O7 by dropping (in a closed vessel with a cold trap) drop by drop ethanol on the heated previously mentionned oxydiser/acid solution; of course don't do this with 96-98% H2SO4 and dry oxydiser, otherwise flash kabooom from Cr2O6 or Mn2O7! No here we are talking about dilluted solutions of acid and mediumly concentrated solution of oxydiser.
True that primary alcools are oxydised to aldehyde first and then to carboxy acid. Here is the reason why this has to be performed at high T (actually 10°C over the boiling point of the aldehyde and lower than the bp of the alcool).
For the low molecular aldehydes, it is often the case that bp is lower than the parent alcool and always than the parent acid; over its boiling point, most of the aldehyde is then volatilised and gets out of reach of oxydiser!
A tiny portion of it remains and react further reason why you mustn't be too concentrate in oxydiser (would you trust KMnO4/acetic acid- it has been involved in many lab accidents). The cold trap is there to collect the volatile aldehyde! This is one of the safest way to get pure aldehydes!

Now ketene process is absolutely hard and waytoo dangerous to do in a lab.

There are many ways to make Ac2O, but all of them involve toxic halides!
The general idea is to get CH3-CO-Cl; now use your brains how to get it from acid halides!
Most of acid halides are easily done, but I'm a bit tired today, maybe another day!
(Tips: S2Cl2, SCl2, SOCl2, SO2Cl2, COCl2, PCl3, PCl5, SCl4, SCl6, POCl3, ....)

Also to madscientist:
CH3-CH2OH + NaOCl --) no acetaldehyde, but a complex mixture of crotonisation of CH3-CH=O (see aldehydes in basic media) into CH3-CH=CH-CH=O and the like; and... chloroform; characteristic reaction of CH3-CO- groups (aceton, acetaldehyde, acetic acid, acetophenone, ...see haloform reaction in basic media); resulting in the splitting of the CH3 and the formation of an acid.
CH3-CH=O + NaOCl -OH(-)-) CCl3(-) + HO-CH=O
CCl3(-) + H2O --) CHCl3 + OH(-)

P.S.: NaOCl is always basic media!

PH Z
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thumbup.gif posted on 26-5-2002 at 21:18


I'm not sure, but I think that it's possible that I prepared acetic anhydride by accident today. I mixed 164g CH3COONa (contaminated with a small amount of NaHCO3, around 1-2g) with 75g 94% H2SO4. I then poured that into a flask; began heating it, and condensing the vapors (typical distillation). I noticed a very strange, sickly-sweet odor; very difficult to describe. I got a whiff of a very small amount of it, causing me to choke for a few moments. The condensed liquid (which I got 22mL of, if I remember correctly) was still liquid, showing no signs of imminent freezing, at -10C. Now, if that had been acetic acid, it would have frozen at a far higher temperature than that... acetic anhydride, on the other hand, would not freeze until the temperature was FAR lower. I'm postulating that the following reactions were occuring, forming at least a fair-sized quantity of acetic anhydride (there is probably a significant amount of acetic acid remaining). Keep in mind that there was a slight stoichemical excess of sulfuric acid.

2CH3COONa + H2SO4 --> 2CH3COOH + Na2SO4
7CH3COOH + H2SO4 --> 7CH2CO + H2SO4*7H2O

And of course, the following occurs:

CH2CO + CH3COOH --> CH3CO(O)OCCH3

Tomorrow I'll try droppering a small amount of the distilled liquid onto an aluminum plate; if there is no visible reaction, then it is definitely high-purity acetic anhydride. Otherwise, it contains at least a medium amount of acetic acid.




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[*] posted on 27-5-2002 at 08:26


I placed a few drops of the distilled liquid on a piece of aluminum foil. No visible bubbling, or audible bubbling, resulted. It has a pH of 1, though.



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[*] posted on 27-5-2002 at 12:15
Hmmm...


I wouldn't be so quick to jump to the conclusion that this is acetic anhydride. I do not observe a rapid reaction with my own glacial acetic acid and aluminum foil. Many favorable reactions of aluminum can be difficult to start because of the oxide coating. Try mixing some of the liquid with water and stirring. Does it form two phases? I think acetic anhydride should be immiscible with water; it will of course be hydrating, but the reaction is slow enough that you should have time to identify a separate phase if you really have the AA.
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[*] posted on 8-7-2002 at 01:23


Aluminium foil usually has a resistant oxide layer, better try zinc powder.
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[*] posted on 9-8-2002 at 08:21


Okay, maybe a brute force way to produce anhydride but it's maybe worth a try.

How about mixing glacial acetic acid with sulfuric acid and magnesium sulfate (anhydrous!). Magnesium sulfate is a powerful dehydration agent.

Here's a hypothetical reaction:

2CH3COOH + H2SO4 + MgSO4
-> CH3(CO)O(CO)CH3 + H2SO4 + MgSO4.H2O

This should be done in a distilling setup and the temperature shouldn't go above 60C to prevent the MgSO4.H2O from dehydrating again.

I doubt if it will work, but who knows?
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[*] posted on 9-8-2002 at 08:26
Ac2O


Ac2O is more hydrophilic than salt hydrates, so the reverse reaciton would predominate. Has anyone ruled out P2O5 dehydration? I dont know too many things that cant be dehydrated by P2O5....

CH3COOH can be dehydrated to acetic anhydride, with ketene, CH2=C=O.

Ketene reacts with acetic acid to dehydrate it straight to Ac2O. Ketene is quite noxious though.

If someone wants to try aluminum, try an aluminum amalgam with HgCl2. This is very commonly used to strip the oxide coating off the aluminum, usually to use the aluminum metal as a gentle reducing agent.

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[*] posted on 9-8-2002 at 11:53


Oh well. I think P2O5 should work, I think I read that somewhere. Too bad it is also a controlled substance.

Hang on a sec. See the thread for preparation of elemental phosphorus. I mentioned something about calciumphosphide. If this is burnt, it will produce P2O5. Now I just have to find a way to collect it.
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[*] posted on 9-8-2002 at 11:54


I can't edit...

Well, if it is more hygroscopic than salt hydrates, that's why I added the H2SO4, this might fix the water. Just trying to combine forces, he.....;)
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[*] posted on 9-8-2002 at 11:54
P2O5


Nobody has ruled out dehydration with P2O5, but you should keep in mind that P2O5 isn't exactly growing on trees around here. Hmm, that reminds me: I've always heard P2O5 referred to as a "powerful" dehydrating agent. Is there anything *more* powerful than P2O5?
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[*] posted on 9-8-2002 at 12:28
Sure


Extremely waterscavenging species do exist, but they are never used for this purpose.

Take alkyllithiums for example, or grignard reagents. They are hard and cumbersome to prepare, and are very useful indeed, and they are extremely reactive toward water and similar loose protons, but no one would ever make these for this purpose.

P2O5 can be bought from chem supply, but it isn't very cheap. I think P2O5 is the strongest "common" dehydrating agent.

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[*] posted on 10-9-2002 at 05:19
magnesium perchlorate for dehydration?


what about using magnesium perchlorate?
its used in place of P2O5( but much easier to prepare...

(ref.: Gmelin, Syst.-Nr. 27, Mg, Tl. B, 1937, S. 154–158 ï Hager 5, 643 ï Hommel Nr. 284.)
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[*] posted on 23-9-2002 at 02:23


Good idea. Only drawback is that perchlorate ions are maybe considered too precious for this purpose?

Also, I've been thinking. Would it be possible to hook up two ethanol molecules by the O by splitting H2O and then oxidizing it to anhydride? Hmm just realize ethers are too unreactive for that...
Mkay, how about hooking up acetaldehyde?




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[*] posted on 23-9-2002 at 10:07


Anodic oxidation of MgCl2 at low current density.
Then carefully heating the resulting
Mg(ClO4)2*6H2O to give Mg(ClO4)2.
Just dont know wether magnesium perchlorate would decompose, or even explode...

Another thing....in the acetaldehyde pathway for Ac2O, how do they manage in industry to get Ac2O and not CH3COOH when oxidizing acetaldehyde???

HLR




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