Sciencemadness Discussion Board


Per - 22-7-2007 at 05:21

I think it doesn´t work but I ask anyway.

Is it possible to prepare Oxalyl chloride with Thionyl chloride and Oxalic acid in excess?

Heared that it doesn´t work because half-chlorides but I don´t know anything about half-chlorides.

If it works, then as follows:


HOOC-COCl + SOCl2 -> ClOC-COCl + SO2 + HCl

Does that work?

Sauron - 22-7-2007 at 05:44


The only two reagents that do work are PCCl5 and TCT (seperately, of course.)

I would say there is a distant possibility that benzoyl chloride or phthaloyl chloride might chlorinate oxalic acid (preferably anhydrous) and so I think those are worth a try.

TCT chlorinates all the dicarboxylic acids except malonic. For oxalic acid the yield is 52% after 3-4 hours stirring at RT in a i:1:1 molar ratio of oxalic acid:TCT:TEA dissolved in just enough dry acetone to give a clear soln.

Per - 23-7-2007 at 04:51

Ok, thanks.

Is then at least the theory about the half-chlorides right?

With the Oxalic chloride I wanted to synthesice DNPO, is there any hope of success by just using Oxalic acid and the alcohol 2,4-Dinitrophenol?
The yields mustn´t be high, may the water could removed by aceotrope distillation with CHCl3 (provided that it could be used as solvent) I think, but just if it works.

Sauron - 23-7-2007 at 07:18

See the old Roger Adams articles in JACS c.1915-1920 (use the ACS pubs search engine.) He originally got involved with oxalyl chloride by using it as a reagent to derivitize phenols, so I bet the anser to your probelm is in one or another of his articles there.

Along the same lines is hexanitro-oxanilide. From trinitroaniline. See Davis' book, chapter on utilization of coal tar.

Nicodem - 23-7-2007 at 09:46

Originally posted by Per
I think it doesn´t work but I ask anyway.

Is it possible to prepare Oxalyl chloride with Thionyl chloride and Oxalic acid in excess?

I can't believe one can have such short memory! You have already been answered this same question in the thread you yourself started:

Per - 23-7-2007 at 11:56

Couldn´t agree, in this thread was just discussed if the method with the Ac2O works and then PCl3, later TCT, couldn´t find there any information if SOCl2 works.

Roger Adams, hmm, could just find "Organic Reaktions Volume" 1-3, search engine, just don´t know what the ACS pubs search engine is, I know just the search of the board and this doesn´t help me.

Davis´ book, do you mean "Hans Eduard Fierz-David and Louis Blangey" or "David Shirley", and yes, searching is not my force, up to now.
Hexanitro-oxanilide seems to be interesting, all I found about it until now is that it´s an exotic oxplosive and hard to get.

And there´s just one more little idea about the prep. of DNPO:

Would Sodium oxalate and 1-Chloro-2,4-Dinitrophenol work, should produce NaCl and DNPO but seems to be too easy to work.

garage chemist - 23-7-2007 at 12:28

Phenyl esters can not be prepared from the phenol and the acid, you can see that in the aspirin synthesis. Only anhydrides or acyl chlorides work, and since oxalic acid does not have an anhydride, only oxalyl chloride can work.

Put your energy towards making PCl5 and anhydrous oxalic acid instead of following routes to nowhere.
A preparation of anhydrous oxalic acid is found here:
The CCl4 can probably be replaced by chloroform, cyclohexane or another solvent with similar azeotrope properties as CCl4.

And PCl5 is made via my method with red P in chloroform. Dont forget to remove the residual chloroform completely in vacuum at 100°C, otherwise it will be impossible to separate from the oxalyl chloride.

Sauron - 23-7-2007 at 16:04

@Per, By "Davis book" I mean The Chemistry of Powder & Explosives" by T.L.Davis, an oldie but goodie, available free from forum library.

The ACS search engine can be found at the website of JACS. I will look up the particular artticle (I probably have it) and post it here.

Nicodem - 24-7-2007 at 02:07

Originally posted by Per
Couldn´t agree, in this thread was just discussed if the method with the Ac2O works and then PCl3, later TCT, couldn´t find there any information if SOCl2 works.

There can only be two simple reasons for your inability to find the answer in that thread. You either have not read it again or you simply do not know that SOCl2 is called thionyl chloride. In short, you are either too lazy or too ignorant. Both these personal characteristics are to be avoided if you want to work with thionyl chloride, oxalyl chloride and similar nasty stuff that can leave you seriously injured.

Sauron - 24-7-2007 at 17:01


Roger Adams, JACS 37 2716 (1915)

describes phenyl esters of oxalic acid prepared by use of oxalyl chloride in presence of pyridine and at low temperature.

Yields described as quantitative.

He gives an example of di-o-nitrophenyl oxalate, which could not be prepared by previous methods.

Adams' method is to chill pyridine to o C, add the oxalyl chloride slowly w/stirring, then using a spatula, carefully break up the yellow lumps of adduct formed. Two equivalents of the phenol are then added and the mixture kept at 0 C for two hours. It is then poured into ice and conc hydrochloric acid. The pyridine dissolves while the pehnyl ester ppts out.

[Edited on 25-7-2007 by Sauron]

Attachment: ja02177a017.pdf (333kB)
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Per - 25-7-2007 at 10:59

For prep. OXALIC ACID (ANHYDROUS) I would of course use the method with the laboratory oven, had tried it former with 150°C and it works.
A heat magnetic stirrer is at least enough to do this properly.

The ACS search engine can be found at the website of JACS. I will look up the particular artticle (I probably have it) and post it here.

That would be great, thanks.

The chemistry of power and explosives:

The nitroanilines are more difficult to nitrate than
aniline because of their inhibiting nitro group, and more easy to
nitrate than nitrobenzene because of their promoting amino group.
In o- and p-nitroaniline the amino and nitro groups agree in
activating the same positions, and both substances yield 2,4,6-
trinitroaniline when they arc nitrated. In m-nitroaniline, the
nitro group "activates" the 5-position, while the amino group
activates the 2-, 4-, and 6-positions. Nitration takes place under
the influence of the ortho-para -orienting amino group, and 2,3,4,6-
tetranitroaniline results.

Trinitroaniline (picramide)
2,4,6-Trinitroaniline, orange-red crystals from alcohol, m.p.
186°, has but little interest as an explosive for the reason that
other more powerful and more valuable explosives may be prepared
from the same raw materials. It may be prepared by nitrating
aniline in glacial acetic acid solution or by the use of mixed
nitric-sulfuric acid in which no large excess of sulfuric acid is
present. The presence of nitrous acid must be avoided, as this
attacks the amino group, replaces it by hydroxyl, and results in
the formation of picric acid. The nitration of aniline in the presence
of a large amount of concentrated sulfuric acid yields
m-nitroaniline73 and later the nitro compounds which are derived
from it.

Have you references about hexanitro-oxanilide,
is the chemoluminescence with this ester better than with DNPO?

And thanks for the attachment, good to have such an instruction.

Sauron - 25-7-2007 at 21:09

Acros sells anhydrous oxalic acid, btw.

Fleaker - 25-7-2007 at 22:18

Ought to be cheaper and less hassle to get from a photography supplier. I think there's a link in the "Readily Available Chemicals" thread in Reagents and Apparatus Acquisition. HTH

freachem - 26-7-2007 at 04:22


Garage chemist, please direct me your method of making PCl5 with red phosphor and chloroform. I searched this forum but found nothing.

Thank you

Sauron - 26-7-2007 at 06:26

Chloroform is only a medium, for elemental chlorination of suspended red P with lots and lots of Cl2.

2 P + 5 Cl2 = 2 PCl5

As you can see the CHCl3 is not involved in the reaction.

If you work out the stoichiometry you will see that need rather a lot of chlorine. DRY chlorine.

Per - 26-7-2007 at 08:32

Here´s the instruction, unfortunately in german:,124916.html#124916

Getting Oxalic acid dihydrate isn´t a problem, could be as mentioned easily obtained from a beekeeper which uses it to prevent his bees from pest.

[Bearbeitet am 26-7-2007 von Per]
There was a problem with the link.

[Bearbeitet am 26-7-2007 von Per]

Per - 3-10-2007 at 06:49

It took some time but now
I have phosphorous pentachloride and anhydrous oxalic acid.

PCl5 was prepared from red phosphorous and dry chlorine without a solvent, just leading the chlorine over the phosphorous, of course very slow. It resublimed at the wall of the Erlenmeyer flask and could be easily scraped off.

But now I´m not sure what´s the best way for preparing oxalic chloride by this chlorinating agent.
Chloroform would be resistant but couldn’t used as solvent because of its similar boiling point.
So the question is if just simple mixing PCl5 and anhydrous oxalic acid is the best way for preparing oxalic chloride, they are both volatile and the products are liquids, Cl-CO-CO-Cl and POCl3.

Eclectic - 3-10-2007 at 06:56

Just mix the reactants dry in the proper proportion and immediately transfer to your distillation setup. When I did this 30 years ago, I tried mixing in a blender, which turned out not to be a good idea because the heat of mixing started the reaction going a bit before I was ready for it.:o

[Edited on 10-3-2007 by Eclectic]

Sauron - 3-10-2007 at 06:59

This reaction is done neat, in absence of solvent.

Here's Sartori:

Much better than the garbled version on rhodium.

Note that it takes 400 g PCl5 to chlorinate <100 g oxalic acid and the yield is 50%. This is not efficient, and that is why I like TCT. However, one uses what one can get.

Further note that POCl3 is by product, don't toss that away!

And protect yourself from both products!

[Edited on 3-10-2007 by Sauron]

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Eclectic - 3-10-2007 at 07:15

Actually, HCl gas is evolved, not chlorine, at least I didn't see any when I did the prep., which I found by following the references in the Merck Index. (way before Rhodium's site)

Sauron - 3-10-2007 at 07:45

(COOH)2 + PCl5 -> (COCl)2 + POCl3 + H20

Neither HCl nor CO2 is evolved in that reaction. But that water will destroy part of the oxalyl chloride and that will form HCl, CO and CO2. Or it might react with the excess of PCl5 producing HCl and more POCl3. Or it might hydrolyze part of the POCl3. Also giving rise to HCl (and H3PO4 or partially chlorinates derivatives.)

(COCl)2 + H20 -> 2HCl + CO + CO2

PCl5 + H2O -> 2HCl + POCl3

POCl3 + 3 H20 -> H3PO4 + 3 HCl

which may account for the 50% yield if the oxalyl chloride reacts with water formed. But if so then there ought to be PCl5 remaining at end, dissolved in POCl3. Per, see what remains if you distill off the POCl3 fraction (bp 107). Maybe you won't have used up all that nice PCl5.

oxalic acid MW 90 (anhydrous)
PCl5 MW 208

So the procedure calls for a 2 to 1 excess of the PCl2. And still a 50% yiels.

But just by looking at it, there ought to be more than 300 g POCl3 and a little more than 125 g in theory of oxalyl chloride so maybe 60-65 g is the realistic guestimate.

To be fair, TCT also gives a 52% yield.

At least this methos is a really good POCl3 prep.

Need I mention that this mixture of two solids needs to be completely ptotected from moisture for the 2-3 days Staudinger (whom Sartori is reciting from) teaches is needed for this reaction?

[Edited on 4-10-2007 by Sauron]

Eclectic - 3-10-2007 at 08:41

"Neither HCl nor CO2 is evolved. But that water will destroy part of the oxalyl chloride and that will form HCl, CO and CO2."

WTF!? :o You mean it's not, but it is?

Anyway, have a way to get rid of the HCl which, like Schrodenger's cat, must be in a state of quantum indeterminacy. :P

I have, in fact, done this reaction myself on exactly this scale. :cool:

[Edited on 10-3-2007 by Eclectic]

Sauron - 3-10-2007 at 19:55

Clearly, what I said was that the reaction of PCl5 and oxalic acid (anhydrous) does not generate any gaseous product.

But I have edited it for clarification.

But the water formed in the reaction reacts either with the oxalyl chloride or with the POCl3 formed or with PCl5 (which is present in excess perhaps for this very purpose.) I can't predict which of those three will occur but whichever, it will produce HCl.

Since the final product is entirely liquid, presumably the PCl5 has all reacted or is in solution in the POCl3 (which certainly dissolves it.)

Ecletic I bow to your superior experience, I never made (COCl2) this way as I was put off by the cost, the low yield and the long rxn time. How much purified product did you get?

POCl3 boils at 107 C. The instruction is to fractionate and collect everything from 60 to 100 C. Then refractionate several times. Oxalyl chloride boils IIRC 62-64 C. So any decent column will cleanly seperate the two, given a >40 C delta in bp. A reasonable reflux ratio is implicit but the fractionation ned not be a heroic one. I'd doubt that more than two distillations are needed. A fast strip and a slow fractionation.

After the strip I'd take off the POCl3, of which there is beaucoup, and see if there's any PCl5 left in the pot.

Assuming that Per started with 100 g red P he ought to have obtained >600 g PCl5, enough to convert maybe 135 g oxalic acid to c.100 g oxalyl chloride, with more than 600 g POCl3 for lagniappe (as we used to say in Nawlins.)

I just ordered 2 Kg oxalyl chloride (about 600 ml per kg). Frigging stuff is $500 a Kg delivered.

I will at some point be trying the following preps of same:

1. TCT in acetone w/TEA
2. Phthaloyl chloride 2-3 mols neat
3. Benzorl chloride (3-4 moles) neat

All with anhydrous oxalic acid.

Another one that just popped up is diethyl oxalate with 2 mols of phthaloyl chloride and chlorosulfonic acid.

The first is in the lit. 52% yield, 3-4 hrs stirring at RT

The rest are in the lit for RCOOH to RCOCl and in a few cases for dicarboxylic acid. But no examples of oxalyl chloride.

However as they are all a lot more convenient than PCl5 (given that red P is a problem for me) they are well worth a try.

[Edited on 4-10-2007 by Sauron]

Eclectic - 4-10-2007 at 05:35

I think I got about 56 grams oxalyl chloride, with the balance almost pure POCl3. It was a long time ago, but I seem to remember capturing HCl with a scrubber.

I thought PCl5 + ROOH ->ROCl + POCl3 + HCl via abstraction of oxygen and intramolecular rearrangement.

Not sure of the mechanism, but I doubt free H2O ever exists in the reaction.

Aldrich has PCl5 by the kilo really cheap in their catalog.

[Edited on 10-4-2007 by Eclectic]

Sauron - 4-10-2007 at 06:47

It's unimportable here. Defense Ministry regs.

And I'm going by the equation in Sartori, see the attachment above. Your equation makes perfectly good sense for a monobasic acid though.

I suppose we could go to the original article in Ber., and see what Staudinger says.

I'm not sure it really matters.

Bottom line is, 400 g of PCl5 gets you <60 g oxalyl chloride, for a lot of work. You have to really WANT the stuff to go this way.

Eclectic - 4-10-2007 at 07:07

I suspect the poor yield is due to oxalyl monochloride decomposition, or reaction of oxalyl chloride with oxalic acid and decomposition to HCl, CO2, and CO.

It's a good prep for POCl3 with a useful byproduct. :(

Sauron - 4-10-2007 at 07:40

Agreed. I did pull up the very first prep of oxalyl chloride, from diethyl oxalate and PCl5. No equation given. This is Compt.Rend. 114 p122-123, attached here.

And I requested the Staudinger paper from Ber. which is sixteen years later. Just out of curiosity.


Of the other preps (known and speculative) that I want to try, the cheapest and easiest is the known TCT one.

Next cheapest is the benzoyl chloride method of H.C.Brown, which of course Brown never claimed for oxalyl chloride.

Phthaloyl chloride is known to work with dicarboxylic acids, often giving the cyclic anhydride when possible. Obviously not possible with oxalic. And no lit. on this.

The same reagent plus chlorosulfonic acid takes esters straight to acyl chlorides so maybe diethyl oxalate to oxalyl chloride, but those reagents are maybe a lil bit exotic for many members.

I'm hoping the TCT will work out. I can't afford too many $500 kilos of oxalyl chloride. TCT is cheap.

[Edited on 4-10-2007 by Sauron]

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Sauron - 4-10-2007 at 16:02

And here is the Ber. paper from 1908.

My German is none too perfect so I could have missed something but I saw nothing about water forming. So maybe Sartori is (ha ha) all wet.

Staudinger tried equimolar oxalic acid and PCl5, also ethyl oxalate and PCl5. Very low single digit % yields. For yields in the 45-60% range 2 mols PCl5 per mol oxalic acid (or diethyl oxalate) are required, so that is one per carboxylic acid function and therefore your equation is correct, Eclectic. HCl is produced not water.

Staudinger also describes a reaction between (COCl)2 and H2S proceeding through a S-heterocyclic intermediate that falls apart to COS and CO.

[Edited on 5-10-2007 by Sauron]

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Per - 5-10-2007 at 11:28

The synthesis was entirely successfully.

I started with 6g of red P, just to see how it works and because of this small amount I didn´t calculated the yield, all what I can say is that I´m now the owner of about 2g oxalyl chloride. (Took it in the freezer and then outside, the liquid was molten when still ice crystals remained on the surface of the flask)

I just putted the anhydrous oxalic acid into the Erlenmeyer flask with the PCl5, at this time I couldn´t see any reaction. Until this time I hadn´t red Saurons further posts so I didn´t wait 3 days and just heated the mix on the heater.
Firstly just vapours of HCl, as imagined, were appeared.
Then the rise liquid took the remaining PCl5 from the wall of the flask, the oxalic acid was distilled of.
By further heating the POCl3 goes over and nearly at the end of this something occurred I didn´t thought of.
White crystals appeared in the flask with the distilled POCl3, they came over with the POCl3 and where also in my cooler and the entire apparatus.

Firstly I thought it could be remaining oxalic acid, but then by rinsing the apparatus with water I noticed a very violently reaction and much HCl was involved, fortunately I did it outside.
Now I´ve a mix containing PCl5 and liquid POCl3.

So, Sauron you were totally right when you said that PCl5 could remain at the end.

And thanks for the very useful information’s, especially for the one in german :) is the entire book free online available?
By the way, I´m worse in reading france stuff but that shouldn´t matter.

Don´t expect that it was easy to get red phosphorous in Germany, they´re all thinking that you´re going to build a bomb with it, just like Al-Kaida or something like that, really paranoid time. Also it wasn´t really cheap 1kg/85€, I just got it because of my connections.

Sauron - 5-10-2007 at 15:52

Chem. Berichte prior to 1902 is available free from Gallica but you have to download page by page.

All of this crucial German journal is available from Wiley.

And Wiley's site is searchable and has DOI's for articles, and with that information you can request articles from References.

Bravo, Per!

Starting with 6 g red P 200 mmol you ought to have ended up with about 40 g PCl5 and that should have been enough to convert 9 g anhydrous oxalic acid into c.5 g oxalyl chloride.

If you had let the reaction run its course I do not think there would have been any unreacted PCl5, because 2 mol PCl5 to one mol anhydrous oxalic acid is the required amount stoichiometrically and not excess at all, as it turns out. I was wrong and Ec;ectic was right.

With a Kg red P you can make >6 Kg PCl5. If you have the patience to generate 5 Kg Cl2.

Probably more than you need.

Per - 9-8-2008 at 04:24

POCl3 + C2O4Na2

Yesterday I did a further experiment in preparing oxalyl chloride:

I placed 10g dry sodium oxalate in an Erlenmeyer flask and added 7,6g phosphoryl chloride.
Nothing happened, but after a few minutes I could see drops at the glas wall, and they became more and more, I know phosphoryl chloride, it would never evaporate so fast at room temperature.
Then I heated the stuff on a hot plate, this gave a very vigorous reaction for a few seconds, white fumes appeared and I could collect a volatile liquid.
It evaporated fare to fast to be POCl3.

All seemed that this must be oxalyl dichloride.

I hoped the reaction could take place like this but I was almost sure that it wouldn´t work, but I want to try it:

3 Na2C2O4 + 2 POCl3 --> 3 C2O2Cl2 + 2 Na3PO4

After this attempt I tested the "oxalyl chloride" in an solution of ethyl acetate, chlorophyll from tee and 30% H2O2, unfortunately absolutely no light was emitted.

With cold water it didn´t reacted as vigorous as oxalyl chloride would do, but it dissolved more quickly as POCl3 would do.

So could anybody tell me what´s formed in this reaction?

PS, the preparation with PCl5 worked fine, POCl3 is a very good solvent in this synthesis, for the chlorination of red phosphorus as well as for the chlorination of oxalic acid.

Sauron - 9-8-2008 at 08:32

I can't tell you what you made but obviously it was not oxalyl chloride.

There is no need to call it oxalyl dichloride, as the half chloride of oxalic acid decomposes immediately. So there is no oxalyl monochloride.

There are only two reagents I know of capable of chlorinating oxalic acid and as you ought to know these are PCl5 and TCT.

TCT is cheaper and easier to obtains.

They are of approximately the same efficiency. TCT on an equimolar basis with anhydrous oxalic acid in acetone and in presence of 2 mols TEA at RT for 3-4 hours well stirred gives a slurry of cyanuric acid from which a 52-54% yield of oxalyl chloride can be extracted with CCl4.

Yields from the reaction of PCl5 and anhydrous oxalic acid, which requires lots of PCl5, are also quite moderate (which is to say modest.) POCl3 is byproduct.

From that alone you ought to have reckoned that POCl3 would not work. If it could work, then the amount of PCl5 could be reduced and the POCl3 would finish the job.

Oxalyl chloride is intensely irritating to eyes and nose/throat, reacts explosively (very violently) with water and almost as violently with lower alcohols. When you have made it you will know at once, and this really ought to be in a good hood. Oxalyl chloride was a candidate war gas in WWI. Its reactivity and short shelf life precluded such use.

The dry distillation of TCT and sodium oxalate in a sealed tube produces very poor resulkts, although the same procedure with sodium benzoate works well for preparation of benzoyl chloride.

chemoleo - 9-8-2008 at 08:58

Over at lambdasyn I read that phosphgen can be prepared by passing oleum over tetrachloromethane, the COCl2 is simply collected in a cold trap. The other product being HSO3Cl, chlorosulphonic acid. The suggestion was made that hexachloroethane could be used to make...presumably the oxalylchloride?
Although I think that the carbon linkage may not survive this treatment.

Sauron - 9-8-2008 at 09:40

Phosgene is usually prepared (by this sort of method) from oleum and CCl4 or CHCl3 and chromic acid.

I also doubt that the C-C bond will survive hot oleum. Phosgene is more likely.

1,1,1,3-tetrachloroethane is I suppose the one you mean. The utility of obtaining chlorosulfonic acid by this route is somewhat offset by the hazards of phosgene. With oleum available, just passing dry HCl in will give you chlorosulfonic acid (and a lot of heat! that must be efficiently removed or you start to distill SO3 which you do not want to do.)

1,1,1,3-tetrachloroethane is a source of the trichloromethyl radical and so is hexachloroethane. Anything that breaks the C-C bond in one case will also do so in the other.

Klute - 9-8-2008 at 10:23

Per, considering the high molecular weight of PCl5 and the small mw of H2O, your oxalic acid might have contained enough water to destroy a certain amount of PCl5? How did you dry your oxalic acid?

The oven method causes a fair amoutn of subliamtion, especially at 150°C.. Azeotropic drying with pet ether for example works well, see my post on the subject. Fot your use, I would recommend drying the filtered anhydrous oxalic acid over P2O5 for example for a day or two, under vacuum if possible.

Sauron - 9-8-2008 at 10:44

For best results I'd buy anhydrous oxalic acid in package size consistent with experimental scale and always use a freshly opened bottle and use in in entirety.

The additional cost is nothing compared to the laboriousness and uncertainty of drying the cheaper hydrated acid thoroughly.

As klute says, the drying temperature is close to decomposition temp, oxalid acid sublimes all over the oven interior, and it is hard to gauge how complete the drying is. Given cost of reagent, the unfavorable MW, and the crappy yield under the best of conditions, this sucks.

woelen - 9-8-2008 at 11:28

I have 100 ml of oxalyl chloride (for just a few euros, it was an old bottle, 20 years old), but it still contains a colorless liquid which fumes in contact with air. Surprisingly, the reaction of oxalyl chloride with water is not that violent. I added a few drops to water. It reacts with water, giving bubbles of gas, but it definitely is not explosive, it is similar to the reaction speed of SOCl2, which also is not explosive.

It is remarkable though, that no oxalic acid remains behind. After all oxalyl chloride had dissolved, I heated the water, driving off all CO2. No precipitate was formed with a solution of Ca(OH)2 when this was added. If oxalic acid were present I would expect a white precipitate of CaC2O4.

Also, Sauron writes about short shelf life, but if properly stored, the shelf life can be very long. The bottle I have has a very special cap, which really well assures that no oxalyl chloride goes out of the bottle or that humid air can get inside.

Sauron - 9-8-2008 at 12:15

Twenty years old? I'm afraid that oxalyl chloride is rather notorious for having a short shelf life, also for building up pressure in as it decomposes. I have a couple of 100 g bottles that are 3-4 years old and I would not trust them unless I distill them and get the right bp. It is best to use oxalyl chloride on a JIT basis - just in time. Order it for when you need it and use it right away.

The high price and short shelf life of oxalyl chloride are the major reasons why it has not become more widely used as a reagent, despite its unique properties.

You can beat the high cost by making your own from TCT and anhydrous oxalic acid.

Then use it quick.

[Edited on 10-8-2008 by Sauron]

12AX7 - 9-8-2008 at 12:55

So what was the liquid, if not oxalyl chloride? Now I'm curious.


woelen - 9-8-2008 at 14:37

Yes, I also am curious now. The stuff is a colorless liquid which is absolutely clear without any solid particles, which is fuming in air. No excessive pressure was present in the bottle. When I received it, it still was sealed, so it never was opened in all those years.

According to this MSDS the liquid is stable under ordinary conditions:

I'm quite sure that if the stuff did decompose in a few years, and if pressure builds up in the containers, then that would be mentioned by the MSDS. I also checked a few other MSDS's and none of them mentions limited stability or shelf life.

Sauron - 9-8-2008 at 21:30

You put too much faith in MSDS. The instability of oxalyl chloride over time (limited shelf life) is mentioned in the literature. I mean the peer reviewed journals, not the monkey spew MSDS.

See the attached first page of JACS paper by H.C.("Boron")Brown, one of the most eminent organic chemists of the centiry, in which he recounts Roger Adams' work on oxalyl chloride, but notes that the high cost and instability of the reagent precluded its general use.

I have seen similar remarks elsewhere in the literature.

Given 20 years internal pressure can slowly dissipate even through a seal. Consider the difficulty of containing Br2 at ordinary temperatures even in nominally gas-tight containers.

If you are set up to handle oxalyl chloride at all, as appears to be the case, why not set up a 250 ml fractionation apparatus and slowly distill the contents of that two decade old bottle? The boiling point of oxalyl chloride is no secret. The procedure will not take long. At its end you will have a far better idea what is in there. Either it is pure oxalyl chloride or it is not.

You might also want to take its density and its refractive index.

If the stuff is wholly or partly decomposed then taking the trouble to QA it will prevent you from ruining an experiment at some point.

The physical constants are

BP 63-64 C
Density 1.455 @ 20 C
Refractive Index 1.4304

The first two ought to suffice if you have no refractometer.

Oxalyl chloride normally breaks down to CO, CO2 and HCl by exposure to humidity.

At high temperatures it decomposes to CO and phosgene. This decomposition is also brought about by UV irradiation or interaction with AlCl3.

I am hoping you have a good bottle, because that probably means my two bottles are also likely to be good.

[Edited on 10-8-2008 by Sauron]

[Edited on 10-8-2008 by Sauron]

Attachment: f_ja01273a014.pdf (116kB)
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Per - 10-8-2008 at 02:28

How did you dry your oxalic acid?

Yes, finding out if the oxalic acid is absolutely dry is really a problem, I just dry it overnight in an oven at about 100 to 140°C, the sublimation is moderate and the crystalls always seems to be dry the next morning.

woelen, may you can put a few drops of the 20 year old oxalyl chloride in a solution of ethyl acetate, a bit tee or any other dye which works and a few drops H2O2, if light is emitted, you can be nearly absolutely sure that it still contains oxalyl chloride.

I can confirm that it doesn´t explode in contact with water, may mine still contains POCl3, does it form a azeotrope with oxalyl chloride?
Also no oxalic acid is formed again with water as woelen said, only in the vapour phase, oxalic acid is formed again. I could observe that by storing oxalyl chloride in an sealed Erlenmeyer flask, after a couple of weeks a few crystals appeared at the wall of the flask.


Good page, is it from the Library of University of California or from the page you can only download page for page?

[Bearbeitet am 10-8-2008 von Per]

Sauron - 10-8-2008 at 02:58

Hydrolysis of oxalyl chloride proceeds through the half chloride which is extremely unstable and falls apart to CO, CO2 and (HCl.

(COCl)2 + 2 H20 ->[ ClC(O)-C(O)OH + HCl ]
-> CO + CO2 + 2 HCl

The overall reaction yields CO, CO2 and 2 HCl

You can analyze for oxalyl chloride by hydrolyzing and measuring the CO given off or, titrating the HCl or both. Details and lit.refs in Sartori.

[Edited on 10-8-2008 by Sauron]

woelen - 10-8-2008 at 07:10

Per, I tried your suggestion and I only can say WOW! I did not know of this reaction, but it is beautiful, I might even make a web-page about this. I knew of the synthesis with 2,4,6-trichlorophenol, but not of plain oxalyl chloride itself being fluorescent in its reactions.

I took appr. 1.5 ml of dichloromethane (I hardly have any ethylacetate, so I decided that DCM could do the job as well). I dissolved 0.1 ml of my oxalyl chloride in it. The liquids perfectly mix, no 2-phase system could be observed at all.

In another 1.0 ml of DCM I dissolved some green fluorescent dye. I'm not sure what dye it is, it has a very intense green color and is both water-soluble and non-polar solvent soluble.

When the liquids were mixed, the green color immediately changed to a light yellow color. So, the oxalyl chloride reacts with the dye. The dye only was present in TINY amounts, probably just a mg or so.

To this liquid, I added some 50% H2O2 in water. Everywhere, where the aqueous layer and the DCM layer are in contact, there is a very bright green light. Even in daylight, it can be seen easily. In a dark place it really looks amazing. The fluorescence does not last for a long time. In appr. 30 seconds all is over. A lot of bubbles are produced, due to decomposition of the oxalyl chloride with the water present and due to the decomposition of the H2O2.

I think that this experiment could be even better if I could extract the H2O2 in some organic solvent in which the oxalyl chloride and dye also can be dissolved, such that no two-layer system is obtained. I tried extracting it in DCM, but that does not work. H2O2 does not dissolve in DCM.

woelen - 10-8-2008 at 10:11

I also did the experiment wth ethylacetate, although I only have a little amount of that. This experiment was a total failure. I did the following:

- Take 1 ml of ethylacetate and add the green dye. The dye did not dissolve easily, only part of it dissolved, the rest remained sticking to the glass.
- Add 2 drops of oxalyl chloride (appr. 0.1 ml). This mixes with the ethylacetate and the dye becomes yellow again.
- In another test tube, add 0.5 ml of 50% H2O2 with 1 ml of ethyl acetate. Shake well. After this, the upper layer is sucked up with a small pipette.
- Add this ethylacetate with extracted H2O2 to the test tube with the oxalyl chloride and ethylacetate.

When the last step is taken, then a lot of gas is produced at once and a thick cloud of HCl appears. No fluorescence at all. When that happened I stepped back and waited in another place until the reaction stopped. I think that both the H2O2 in the added ethylacetate and the oxalyl chloride decompose at once. Remarkably, only little heat was produced. Immediately after the violent reaction, the test tube only was luke warm, I think something like 40 C.

[Edited on 10-8-08 by woelen]

Per - 11-8-2008 at 01:59

Pages about the peroxodichemilumineszenz:










Curious that the ethyl acetate don´t worked, may it´s more suitable for TCPO and DNPO, in mine Experiment it worked also with oxalyl chloride.
Dichloromethane is of course the better solvent in this reaction, but normally ethyl acetate is the cheaper one.

I´m always using bis(2,4-dinitrophenyl)oxalate, because I can´t get 1,3,5-trichlorophenol.

I had also the idea of water free H2O2 in an organic solvent, but I hadn´t the time to work it out.
May one could dissolve the H2O2 in a solvent which is not soluble in concentrated sulphuric acid, so the organic phase could be left over it
or the water could be removed by aceotropic distillation, may chloroform works, or decomposes H2O2 at nearly 70°C? Hopefully it don´t forms a aceotrope with H2O2.
Vacuum distilled H2O2 is maybe too hazardous and not everybody has a vacuum pump.

Klute - 11-8-2008 at 04:28

Maybe an equivalent like perborate or persulfate could work, as a finely dispersed solid? It could release minute amounts of H2O2, or peroxo radicals, in presence of traces of water, along with a dash of PTC?

Otherwise, you could try making a solution of peracetic acid in ethyl acetate, adding conc. H2O2 to AcOH dissolved in AcOEt... I guess AcOOH is more soluble in AcOEt than straight H2O2.

woelen - 18-8-2008 at 06:42

I stumbled upon this patent. It might be interesting for someone with access to trichloroacetyl chloride (CCl3COCl) or the skills to make this nice little beast. The other chemicals, required for this are easy to obtain or make (ethylene glycol and chlorine gas).

The funny part of this patent is that the trichloroacetyl chloride is regenerated, so a small quantity could be used for making a lot of oxalyl chloride.

The process, however, only is something for the most skilled and most equipped persons over here. I myself will not try this at my home :P.

[Edited on 18-8-08 by woelen]

Sauron - 18-8-2008 at 13:32

Trichloroacetyl chloride is not very stable, it likes to fall apart to CCl4 and Cl2 and CO, the elements of phosgene.

This is vis the trichloromethyl radical.

I think that patent has been kicked around herre before.

Sauron - 25-8-2008 at 22:01

The question of this thread topic, will SOCl2 chlorinate anhydrous oxalic acid to oxalyl chloride, was answered in JACS a LONG time ago.

The answer is NO. SOCl2 fails with oxalic acid.

On the other hand it succeeds with malonic acid where all other chlorinating reagents fail. See my thread on malonyl chloride, the paper describing the failre of thionyl chloride with this substrate ((COOH)2) is there.

L. McMaster, F. F. Ahmann
J. Am. Chem. Soc.; 1928; 50(1); 145-149.
DOI: 10.1021/ja01388a018

[Edited on 26-8-2008 by Sauron]

Per - 4-9-2008 at 01:46

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
3RCH2COOH + PX3 > 3RCH2COX + H3P03

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?

Sauron - 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.

ScienceSquirrel - 4-9-2008 at 02:54

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.

Sauron - 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]

ScienceSquirrel - 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.

Sauron - 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


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


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]

ScienceSquirrel - 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]

Sauron - 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.

S.C. Wack - 4-9-2008 at 05:39

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.

Sauron - 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.

ScienceSquirrel - 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]

Sauron - 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]

Attachment: US2816140.pdf (205kB)
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Sauron - 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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Sauron - 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]

Attachment: carothers.pdf (2.5MB)
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Sauron - 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]

Attachment: 3037-3040.pdf (209kB)
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Per - 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.

Sauron - 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.

Per - 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?

Sauron - 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.

Per - 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.

Sauron - 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.

halogenstruck - 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]