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Eclectic
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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. 
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Sauron
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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]
Attachment: CR114p122_123.pdf (93kB)
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Sauron
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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
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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.
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Sauron
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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.
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Per
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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.
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Sauron
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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.
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chemoleo
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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.
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Sauron
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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.
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Klute
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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.
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Sauron
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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.
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woelen
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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.
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Sauron
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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]
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12AX7
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So what was the liquid, if not oxalyl chloride? Now I'm curious.
Tim
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woelen
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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:
http://www.jtbaker.com/msds/englishhtml/o6180.htm
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.
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Sauron
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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
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| Quote: | | 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.
f_ja01273a014.pdf
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]
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Sauron
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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]
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woelen
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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.
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woelen
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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]
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Per
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Pages about the peroxodichemilumineszenz:
- http://www.chemie.uni-jena.de/institute/oc/weiss/peroxyoxala...
- http://kaltes-licht.fsla.at/chemolumineszenz/frame_rechts_ch...
- http://kaltes-licht.fsla.at/experimente/frame_rechts_expo_ch...
- http://www.shsu.edu/~chm_tgc/JPPdir/JPP1999/
- http://chemmovies.unl.edu/chemistry/dochem/DoChem113.html
- http://www.shsu.edu/~chm_tgc/chemilumdir/chemiluminescence2....
- http://www.lumigen.com/documents/chemexplained.shtml
- http://www.glowspace.com/classroom.htm
- http://www.old.uni-bayreuth.de/departments/ddchemie/umat/che...
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.
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Klute
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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.
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woelen
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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).
http://www.freepatentsonline.com/4301092.html
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  .
[Edited on 18-8-08 by woelen]
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Sauron
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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.
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Sauron
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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.
ACTION OF THIONYL CHLORIDE ON ORGANIC ACIDS1
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]
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