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Author: Subject: HOOC-COOH + SOCl2 -> ClOC-COCl ?
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[*] posted on 4-9-2008 at 01:46


SUBSTITUTED ALIPHATIC MONOBASIC ACIDS
Acetic acid can be chlorinated by gaseous chlorine in the presence of red
phosphorus as catalyst to yield successively mono-, di-, and tri-chloroacetic
acid ; the reaction proceeds better in bright sunlight. If the chlorination is
stopped when approximately one molecule of chlorine per molecule of acetic
acid is absorbed the main product is monochloroacetic acid :
CH3COOH + Cl2 --> CH2ClCOOH + HC1

The conversion of an aliphatic carboxylic acid into the a-bromo- (or achloro)
acid by treatment with bromine (or chlorine) in the presence of a
catalytic amount of phosphorus tribromide (or trichloride) or of red phosphorus
is known as the Hell-Volhard-Zelinsky reaction. The procedure
probably involves the intermediate formation of the acyl halide, since it is
known that halogens react more rapidly with acyl halides than with the acids
themselves:
3RCH2COOH + PX3 > 3RCH2COX + H3P03
RCH2COX + X2 • RCHXCOX + HX
RCHXCOX + RCH2COOH • RCHXCOOH -f RCH2COX

VOGEL, Practical Organic Chmistry


So preparing the Trichloroacetyl chloride wouldn´t be the Problem in this proces, the real problem is the purification ot the Trichloroacetyl choride, because it decomposes as Sauron already told at it´s boiling point.

For the practical preparation of the Trichloroacetyl chloride I´d consider using the reaconed mass if red P which would be necessary to form 1 mole Acetyl chloride out of one mole Acetic acid, so the mix just has to be chlorinated further untill no Cl2 is absorbed anymore, prefrably in the sunlight after a while.
But after that the first problem would be to get the Phosphorus acid out of the desired product. Simple distillation can´t employed, maybe vacuum distillation can do the job.


Theres another thing, Ethylen glycol should be easily and cheap available as an anti-ice solution for cars for example, I looked it up in ebay and other shops but can´t find any of it.
woelen, do you know a cheap supplier for it?
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[*] posted on 4-9-2008 at 02:46


I'm sorry, but it seems you have not been paying attention to some other threads.

Trichloroacetyl chloride is UNSTABLE. As it is formed, it tends to fall apart to the trichloromethyl radical and CO and phosgene (CO + CL2 -> phosgene. See the H.C.Brown article on preparation of volatile acid chlorides by use of benzoyl chloride. About 40% of the expected 90% yield of trichloroacetyl chloride ended up as CCl4 or CHCl3.

See also the review I posted on chlorination of ethanol to chloral. If dichloroacetaldehyde is chlorinated with UV the trichloroacetaldehyde goes to trichloroacetyl chloride then falls apart as described above. This is so efficient that I plan to use it as a method of preparing CCl4.

You will find putting on the third Cl into acetic acid slow and wasteful of chlorine. That is also true of the chlorination of ethanol/acetaldehyde.

And you will find the preparation of trichloroacetyl chloride from trichloroacetic acid, low yielding by any method I know of.

Strictly speaking the HVZ reaction refers to the alpha-monohalogenation of a carboxylic acid catalyzed by a halogen carrier such as PX3 (X=Cl, Br) or the elements of PX3 (red P in presence of halogen. While the classic work on mechanism states that dihalo and trihalo products can be prepared by this method, the author cites no reference and in the literature, none could be found by either Nicodem or myself. HVZ is not done under UV or sunlight. The mechanism is that the PX3 forms some RCOX and that is easily a-halogenated by an enol process.

The photochemical chlorination of acetic acid is a free radical chlorination. If you try to apply it to higher carboxylixc acids you will get complex mixtures. In the special case of acetic acid, where the alpha carbon is the only carbon with hydrogens on it, it works fine. You could do same reaction in the dark with SOCl2 as Cl radical source and a radical initiator like benzoyl peroxide or AIBN at the usual temperatures for such initiators.




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[*] posted on 4-9-2008 at 02:54


Quote:
Originally posted by Per
Theres another thing, Ethylene glycol should be easily and cheaply available as an anti-ice solution for cars for example, I looked it up in ebay and other shops but can´t find any of it.
woelen, do you know a cheap supplier for it?


Pure ethylene glycol can be bought as antifreeze from garages.
You have to check the labels to find out if the brand that you are buying is 100% ethylene glycol and it has a blue dye added so that solutions of it look undrinkable.
Distillation yields the pure water white glycol.
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[*] posted on 4-9-2008 at 03:03


Ethylene glycol is very very hygroscopic and you will only liberate it from water by use of special techniques exclusing all atmospheric moisture. It also forms beaucoup azeotropes with water IIRC.

Most antifreeze in US these days is propylene glycol which is nontoxic. EG if you drink it, will kill you (40cc is LD50) and is is relatively prominent statistically in accidental poisonings and the other kinds (homicides.)

------------------

I guess what I am saying boild down to: trichloroacetic acid is a pain to make and converting it to the acyl chloride lossy (except with TCT).

So if you can buy trichloroacetic acid why not just buy trichloroacetyl chloride if you want to try out the patent procedure? It is rather nasty stuff after all.

And if you can get (or make) TCT then just make oxalyl chloride from oxalic acid (anhydrous) with 54% yield and be done with it without all the other bother?

[Edited on 4-9-2008 by Sauron]




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[*] posted on 4-9-2008 at 03:19


I am sure that in the great nanny state of the US of A ethylene glycol antifreeze is well nigh unobtainable.

However in the small soggy islands off the coast of Europe it is still available as a specialist antifreeze.
I have distilled it and the bp was spot on so I would guess that they are buying the commercial technical product, bunging in the blue dye as per some regulation and repacking it.
It is also available as a 99+% product from several small companies that specialise in cleaning solvents.
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[*] posted on 4-9-2008 at 03:40


I guess in the small soggy islands you refer to, the government is still desirous of decreasing the surplus population, as Dr Swift described in his "Modest Proposal".

I just buy anhydrous EG reagent grade and open a fresh bottle in glove box under dry conditions. And I try not to imbide.

I looked up that patent, and indeed I have seen it before. It calls for esterifying EG with two mols of trichloroacetyl chloride, forming

Cl3CCOOCH2CH2OOCCCl3

and then photochemically chlorinating that to replace all four hydrogens so we now have

Cl3COOC(Cl2)C(Cl2)COOCCl3

and decomposing that to (COOH)2, and trichloroacetyl chloride.

Oxalyl chloride boils low so distilling it out of the acyl chloride would not be a problem.

I am still rereading the patent but I really find it hard to suspend my disbelief concerning the stability of trichloroacetyl chloride under reaction conditions that include UV photolysis, given what we know about the trichloromethyl radical.

Also look at some of the details. The purity of the Cl2 and the low oxygen content. Is this achievable with lab generated chlorine? All the specs for low metals content. Ugh. The large excess of the acyl chloride in the first (esterification) step and utilized again in the third step as solvent. Bottom line is lots of trichloroacetyl chloride to prepare a little oxalyl chloride. I question the economy, even though the trichloroacetyl chloride is regenerated. Note the high boiling residue of partially chlorinated glycol trichloroacetates after after removal of oxalyl chloride and trichloroacetyl chloride. So recovery of the acyl chloride is not quantitative, not nearly.

This process appears to be costly and fiddly. Still it might be interesting to mess with if one is set up for chlorine generation, UV photochemical reactions, and handling/distillation of very aggressive acid chlorides.

I seriously doubt anyone actually makes oxalyl chloride this way commercially.

[Edited on 4-9-2008 by Sauron]




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[*] posted on 4-9-2008 at 04:10


There are a lot of propylene glycol antifreezes on the market round here and I would guess that the ethylene glycol one is only used by a small fraction of consumers now.

Anyway I am offering the information in the spirit of the boards which is to obtain your chemicals by OTC routes etc.

Ethylene glycol antifreeze is available in at least parts of the States.
Probably not California but almost certainly Alaska. Or as it was quaintly referred to last night on television Russian America, which it was until the Russians sold it the Americans. A dreadful shame as the Cold War could have been so much more interesting!

[Edited on 9-4-2008 by Polverone]
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[*] posted on 4-9-2008 at 04:35


The southernmost Russian (tsarist of course) settlement in North America was not in Alaska but in northern California near Sacramento. And the purchase of what at the time was jeeringly described as Seward's icebox, was unpopular at the time and of course now Alaska is by a wide margin the largest state in the 50 states. Texas looks larger on Mercator-projection maps because it is much closer to the equator. But if you look up the actual areas of the two, you will be surprised and Texicans will get their comeuppance. They are in second place. They are still largest of the 48 contiguous states, if they find that a comfort.



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[*] posted on 4-9-2008 at 05:39


Quote:
Originally posted by Sauron
Ethylene glycol is very very hygroscopic and you will only liberate it from water by use of special techniques exclusing all atmospheric moisture. It also forms beaucoup azeotropes with water IIRC.

Most antifreeze in US these days is propylene glycol which is nontoxic.]


It is dried the same way as anything else without difficulty.
It does not form an azeotrope with water.
I'd guess that 95% of the antifreeze on the shelf is ethylene glycol. By guess I mean that it might be 94 or 96%. I've heard that the product sold in California has something in it to make it bitter, to humans at least.

Extremely simple to isolate and purify it from antifreeze.
No idea what all this has to do with the subject.




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[*] posted on 4-9-2008 at 06:42


If by "the subject" you mean the thread topic, reaction of oxalic acid with thionyl chloride - then it has nothing whatsoever to do with it. That reaction does not work, and I posted the literature to prove it.

If you mean the extended topic including this patent, which has come up in other threads before, then EG has some significance. If anyone thinks they are going to meet the purity requirements of the patent with EG salvaged from OTC antifreeze mixtures, then have at it. I prefer to work with the inexpensive already dried reagent grade EG readily available to me.

But thanks for the correction about the azeotropes, or non-azeotropes. EG forms many many azeotropes but not with water.




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[*] posted on 4-9-2008 at 07:01


I think the major problem is not the ethylene glycol but obtaining the trichloroacetyl chloride and the chlorine.
I think realistically you would need a cylinder of the gas.

And if you can buy trichloroacetyl chloride you can certainly buy oxalyl chloride. It is not cheap but most people won't need a lot anyway.

[Edited on 4-9-2008 by ScienceSquirrel]
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[*] posted on 4-9-2008 at 07:31


Oxalyl chloride costs me $650-$700 per Kg.

The ex works from Acros is well over $300. By the time it gets here and the agent takes his bite, after duty and VAT, it';s as above. So, preps for it are always of interest. I only know of two. The one from PCl5 is mostly unavailable these days, and anyway is 50% yield and uses a lot of pentachloride. The TCT method is only slightly better yield, but the reagent is cheap and for some more accesible, and more efficient.

Oxalyl chloride has general utility and some unique reactivity. It would be more widely used if it were not so dear. As for not being someone everyone needs much of, that's a criticism with little meaning as it could be levelled against many or most useful reagents, and its validity depends entirely on what work is at hand, doesn't it?

The patent deserves closer attention. Trichloroacetyl chloride is not that impossible to make or buy. One just has to know the pitfalls.

-------------------------

What is really worth a look is the earlier Depont patent cited in the patent above, I have attached it below. It describes the catalyzed thermal decomposition of various esters of tetrachloroethylene glycol to give oxalyl chloride and the corresponding acyl chloride. In the simplest case tetrachloroethylene carbonate is decomposed to oxalyl chloride and phosgene (the acyl chloride of carbonic acid). The tetrachloroethylene carbonate is prepared per another Depont patent by photochemical chlorination of ethylene carbonate in CCl4. Ethylene carbonate is cheap.

More interesting is tetrachloroethylene oxalate, prepared by transesterification of ethylene glycol with diethyl oxalate followed by photochemical chlorination. The decomposition of this ester gives only oxalyl chloride.



[Edited on 5-9-2008 by Sauron]

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[*] posted on 4-9-2008 at 10:45


And here is the patent from same Dupont chemists detailing the preps of the requisite tetrachloroethylene carbonate and tetrachloroethylene oxalate.

Take note that unlike the case of tetrachloroethylene carbonate, the decomposition of tetrachloroethylene oxalate produces no phosgene byproduct, only oxalyl chloride.

I really do find the "inferior" Dupont process, using ethylene oxalate low-MW polymer, chlorinating this, and decomposing the tetrachloroethylene oxalate in chlorobenzene, much more economical and accesible than the "improved" French method - at least on a bench scale. NO trichloroacetyl chloride called for. The ethylene oxalate might even be commercially available, anyway it is easily made from diethyl oxalate and ethylene glycol by transesterification.

I have requested the original Carothers paper cited in the patent. Carothers of Dupont, "The Great Synthesist."

From one mol diethyl oxalate and 1 mol ethylene glycol there is obtained about 110 g of crude polymeric ethylene oxalate.

From this is obtained by chlorination under UV over 32 hours, 256 g tetrachloroethylene oxalate.

Decomposing this catalytically in chlorobenzene using Darco G-80 (a Norit activated carbon product) almost 100 g oxalyl chloride is isolated by distillation.

The obvious weakness of this process is that it is expensive of chlorine; the above example starting on a 1 mol basis requires 480 g Cl2 over 32 hours. However the French "improvement" is just as inefficient in its use of Cl2. This is to be expected as the reaction is after all a chloro-oxidation!

[Edited on 5-9-2008 by Sauron]

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[*] posted on 5-9-2008 at 09:13


To recap, the route using the trichloroacetyl ester of ethylene glycol is objectionable because the requisite acyl chloride is expensive, hard to prepare and unstable.

Trichloroacetyl chloride in fact costs more than oxalyl chloride.

Falling back on the Dupont method using ethylene oxalate from diethyl oxalate and ethylene glycol works fine but just as with the French patent, is expensive of chlorine.

1.5 Kg diethyl oxalate and 900 g EG will get you 1 Kg oxalyl chloride but you will need to generate and use 5 Kg Cl2 (70 mols) along the way.

That's a lot of Cl2, no matter how you generate it.

Attached is Carothers' paper on ethylene oxalate from 1930.

[Edited on 6-9-2008 by Sauron]

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[*] posted on 5-9-2008 at 19:18


The far easier way to prepare oxalyl chloride is to use cyanuric chloride on anhydrous oxalic acid, in dry acetone at RT with stirring for 3-4 hours.

K. Venkatamaran & D. R. Wagle
Tet. Lett., No. 32, 3037-3040 (1979)

Yield 52%. Which is no worse than the prep from unobtainium PCl5. See attached below. This is original from Tet Lett, not the incomplete version from Rhodium and Erowid.

For those who cannot buy Cc (aka TCT and NOT to be confused with TCCA) you can buy or prepare methyl thiocyanate, then chlorinate that and get a good yield of cyanuric chloride from first phase of chlorination. About 65-70 g per 100 ml of the thiocyanate. Also continued chlorination of the mother liquor mixture will get you a bunch of CCl4 along with SCl2 and S2Cl2.

See threads on this and the J.Chem.Soc. paper by James I posted there.

Methyl thiocyanate is not particularly cheap and is made from CS2 ultimately, also not cheap. Or you can use MeI or dimethyl sulfate if you prefer. Better have a good hood. Same applies of course to oxalyl chloride and for that matter, trichloroacetyl chloride.



[Edited on 6-9-2008 by Sauron]

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[*] posted on 6-9-2008 at 04:34


So this process seems far more economically than the french process and it would be the first time I´ll trying out an transesterification.

Chlorine could be conveniently prepared by dropping hypochloric acid into calcium hypochlorite, this was once available on ebay as a fast swimming pool chlorinator but now it isn´t any more, TCCA is still available.

The Ethylene glycol makes me more concerns, may I´ll find a bottle of it at a local supplier or may they could order it, my car just get´s along with Ethylene glycol and nothing else.

All in all very interesting synthesis, may also for the preparation of other acid chlorides.
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[*] posted on 6-9-2008 at 05:26


I' suggest using a RB or pear shaped flask with a large center neck like 45/50 and a reducing adapter to take a Claisen head, rather than a Claisen flask. (For the transesterification.) First ethanol comes over then a middle fraction that is a mixed ester and it's the white tacky polymeric residue that is left in the pot that is the ethylene oxalate. Removing that from a formal Claisen flask sounds like a chore.

Why not generate Cl2 from HCl and TCCA? Should be available at same place calcium hypochlorite is for same appln: swimming pools.

Anyway I'm glad you like it.




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[*] posted on 6-9-2008 at 08:47


Why making the work of removing the ethylene oxalate from the flask, just adding the solvent, CCl4 or CCl2, vigorous stirring for a few minutes and start chlorination should also be possible.

The same for the tetrachloroethylene oxalate, and are there any chances that the decomposition of the tetrachloroethylene oxalate takes place without Darco G-80, could anything more available be used or does it decomposes at a high temperature without a solvent, then with the fear that the oxalyl chloride decomposes also?
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[*] posted on 6-9-2008 at 19:37


The AC was only one of numerous catalysts possible. I am sure that triethylamine hydrochloride ("chlorohydrate" as the French called it in their patent) works fine, in a trivial amount.

Yes, solvent will help removal of the oxalate but a Claisen flask (unitary as opposed to assembled from standard joints) is not what you want, and not ideal for the subsequent steps either.




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[*] posted on 7-9-2008 at 01:33


Triethylamine hydrochloride is readily available for me, so this step isn´t a problem anymore, thx.

Now I understand how the industry can make oxalyle chloride relatively cheap, if they had to make it out of PCl5 or TCT it would be much more expensive.
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[*] posted on 7-9-2008 at 04:13


TCT is cheap. It's made industrially from HCN and HCL -> ClCN then trimerize to TCT. But for insight into how oxalyl chloride is made industrially see Ullmann's and Kirk-Othmer. I do not believe these processes involving tetrachloroethylene glycol esters are commercial.



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[*] posted on 3-4-2010 at 13:15


i mixed SOCl2 and Oxalic acid.there was not any trace of oxalyl chloride.
i mixed SO2Cl2 and SODIUM Oxalate.there was not any trace of oxalyl chloride although one of my friends fold me he made it in this way in diethyl ether as solvent.i did not use any solvent but just dry.PCl5 gave a good result.
i read in Organikum book that it will result in (COCl)2 if anhydrous ZnCl2 catalist is used.

[Edited on 3-4-2010 by halogenstruck]
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