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Author: Subject: Synthesis of unsym.-heptachloropropane and hexachloropropene (updated)
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[*] posted on 2-10-2010 at 06:25
Synthesis of unsym.-heptachloropropane and hexachloropropene (updated)


In the book "Excuse me Sir, would you like to buy a kilo of isopropyl bromide?" which was recently uploaded to the Library by Polverone and which is very much worth reading, I read about how Mr. Gergel made heptachloropropane from perchloroethylene, chloroform and aluminum chloride.
I immediately thought "I'd like to do this!" and soon found the preparation on Orgsyn.
Moreover, the article about the discovery of the AlCl3-catalysed addition of chloroform to perchloroethylene by Prins and Boeseken turned out to be available for free.

I found that the exothermic nature of the reaction does not pose nearly as much of a problem as Mr. Gergel made it seem.
It probably only becomes potentially hazardous when working on a massive scale, like Mr. Gergel with his multiple runs in 12-litre-flasks.

Purification of the reagents

The chloroform was shaken twice with conc. sulfuric acid to remove the ethanol that was added as a stabilizer.
The first washing with H2SO4 produced a slight but noticeable exotherm as the alcohol was absorbed by the acid. No exotherm occured with the second acid washing.
It was then washed with water, Na2CO3 solution and water again, dried over CaCl2 and distilled over phosphorus pentoxide to render it perfectly dry and alcohol-free. It was stored in the freezer until use.

The technical perchloroethylene does not seem to require any purification, as a small test batch that I made with C2Cl4 straight from the bottle went fine and gave a high yield of good product.
Neverteless, I washed the C2Cl4 used for the depicted preparation with aqueous HCl to remove amine stabilizers, then with water, dried it over CaCl2 and distilled over P2O5.
I don't think that this made any difference.

The commercial AlCl3 was used straight from the bottle, a fine, somewhat yellow, free-flowing powder that gives off HCl fumes.

Preparation

I essentially followed the Orgsyn procedure, but used a little more C2Cl4 in order to better utilize the large CHCl3 excess.

180g Tetrachloroethylene and ca. 300g (200ml) chloroform were placed in a 1-litre round-bottom flask. 27g anhydrous aluminum chloride were added, and the flask was fitted with a reflux condenser which carried a drying tube filled with CaCl2.


The mixture was then heated with a bunsen burner. It assumed a red-brown color and continuously darkened as it it started to reflux.


After a few minutes of gentle refluxing, the boiling slowly grew more intense despite the heating not being increased.
The mixture soon reached a vigorous boil and continued boiling upon removal of the heat source:


A continuous stream of condensate was pouring down from the reflux condenser, but everything remained under control and the exothermic reaction gradually subsided over the next fifteen minutes.
The heating was now resumed and the mixture was refluxed overnight, for about 15 hours:


The next morning, its color had changed to dark green and it looked like this:


From below:


The mixture was cooled to room temperature, and the AlCl3 formed a precipitate on the bottom of the flask.
The supernatant was decanted from the sediment into a 500ml separatory funnel.

This is the sediment in the flask:


I added 100ml water to this black sediment. This turned out to be a mistake. It seems like the sediment consisted almost entirely of active AlCl3, and with the water it erupted into white smoke. The flask continued spewing acrid smoke for over a minute, and was extremely hot afterwards.
On the left side, you can see the black supernatant in the separatory funnel, covered with water:


Only a small amount of black lower phase remained in the flask where I hydrolyzed the AlCl3, and this was not further worked up.

The supernatant from the decantation of the AlCl3 (i.e. the main part of the reaction mixture) was thoroughly washed with three 100ml portions of water, which did not produce any noticeably exotherm.
After that, its color had changed from black to orange and it became almost clear:


It was dried over calcium chloride, with frequent shaking. This turned it clear:


A vacuum distillation apparatus with oil bath was set up.
You can see the frying fat in the oil bath/magnetic stirrer combo melting:


First the chloroform was distilled off at atmospheric pressure and up to 160°C oil bath temperature:


Then the mixture was cooled a bit, and very slowly and carefully, a vacuum was pulled.
This is the pump I used, a very old, obscure model from Vacuubrand that has 4 separate membrane pump heads on one motor shaft. Originally, those were connected in parallel, giving a useless 100 torr vacuum. After disassembly and cleaning of the pump heads, I connected three of them in series and added a gas ballast valve. Now it goes down to 18 torr- significantly better than my water aspirator, which I previously used for all my vacuum distillations. I checked the solvent resistance of the rubber diaphragms and they are resistant against chloroform.


Some residual chloroform and a small amount of perchloroethylene came over in the forerun.
Then the product started distilling at around 120°C vapor temperature (sorry, I can't give an exact boiling point, the vacuum varied somewhat, and the product even boiled over a slight bit at one time- this was my first time distilling with a diaphragm pump).

This is the distillate- essentially clear, only a few drops of the colored sump fell into it from the boiling over.


The distillate was a supercooled melt- it did not solidify despite the cooling water being colder than the 30°C melting point.
Upon adding a seed crystal of heptachloropropane from a previous batch, it rapidly started crystallizing, warming itself up to 30°C in the process:


It was remelted and poured into a wide-mouthed plastic dish with screw-on lid to solidify (the crystal cake is hard and difficult to remove from a glass dish).
The lid was screwed on tightly, as the solid has a significant vapor pressure, evident from the smell, and sublimates readily at room temperature, similar to iodine.
The plastic dish was hit with a hammer to break loose the crystal cake and coarsely crush it.
The crystal shards:


The final product: 280g from this batch, together with the previous batch, giving over 300g in total and filling a 250ml bottle to the edge:


Discussion

I obtained a yield of 280g (91%), which is almost exactly the literature yield from the Orgsyn procedure.
The product has a pronounced sweetish camphor-like smell which quickly fills the laboratory.

Dehydrochlorination to Hexachloropropene: Coming soon!
If you wonder why one would want to make hexachloropropene: look at patent US1320869.
Also, it can be used to make anhydrous metal chlorides from the oxides (e.g. Brauer: UCl4 from UO2) and from hydrated chlorides.
Look at this post from benzylchloride1 about the preparation of anhydrous chromium(III)chloride. This is the only previous mention of hexachloropropene on this forum.

[Edited on 12-10-2010 by garage chemist]




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[*] posted on 2-10-2010 at 10:32


Congratulations! Nice work and interesting (though obscure) target you choose.
If nothing else, it could certainly become an viable, though weird, choice of an reagent for preparing acyl chlorides from carboxylic acids.
Here are some additional ideas for your new exotic reagent applications:

According to EP0902019, unsym-heptachloropropane can be reacted with guanidine to give some interesting chlorinated 2-aminopyrimidines, though in poor yields. Possibly the same reaction principle could be applied with some other binucleophiles to give other chlorinated heterocycles (with hydrazines, thioureas, etc.).
SU423790 describes conversion of unsym-heptachloropropane to Cl2CHCCl2CO2H by "by hydrolyzing Cl2-CHCCl2CCl3 in oleum contg. I2 at 50 °C".
The next one is not really useful as it is essentially describing the retro reaction of what you did, but it is an interesting account exactly due to this curiosity. In short, the abstract of Zhurnal Organicheskoi Khimii, 7 (1971) 439-443 says: "The primary reaction during the thermal decompn. of CCl3CHClCCl3 at 180 over AlCl3 was dehydrochlorination to give CCl2:CClCCl3; the latter reacted further to give CCl4, CCl2:CClCCl:CCl2, C2Cl4, C2Cl6, COCl2 (frp, oxidn. of ClCCCl), and a polymer with empirical formula C5Cl4. Under similar conditions, CCl3CCl2CHCl2 yielded 1:1 CHCl3-C2Cl4. Carbonium-ion mechanisms for these reactions are presented."




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[*] posted on 2-10-2010 at 16:15


Very nice writeup as usual. Like Mr. Gergel, are you planning to ultimately prepare trifluoroacetic acid and derivatives?



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[*] posted on 2-10-2010 at 16:58


Thank you both.
No, I am not planning on making trifluoroacetic acid as I don't have SbF3, or any antimony compounds for that matter.
I want to explore the possible uses of hexachloropropene as a chlorinating agent for both inorganic and organic chemistry, especially for the preparation of acyl chlorides or anhydrides in the latter case.
I have already found a procedure for MoCl5 from MoO3 and hexa by simple refluxing followed by removal of the volatiles under vacuum. I'm sure that many more metal chlorides can be prepared that way.
I'm also interested in possible uses for the byproduct trichloroacrylic acid.




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[*] posted on 2-10-2010 at 17:02


Great work! If this brings acyl chlorides and anhydrides within reach, this will be a huge step forward for the amateur chemist.



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[*] posted on 2-10-2010 at 17:27


Something I've been thinking about : might it be possible to make acyl chlorides and anhydrides from trichloroethylene and sulfuric acid? TCE is hydrolyzed to chloroacetic acid by 75% sulfuric acid. Depending on what the mechanism is here, perhaps it's possible to get an acyl chloride or anhydride instead, by adjusting concentrations and reaction conditions?

EDIT: It's apparently possible and generalizable to other 1,1-chloroalkenes as well! Patent #3742047

http://www.freepatentsonline.com/3742047.html

[Edited on 3-10-2010 by madscientist]




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[*] posted on 2-10-2010 at 17:58


Well done garage chemist! And thanks for the beautiful write-up. As you and benzylchloride1 have indicated this may well prove to be a relatively facile route to the much desired acyl chlorides.



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[*] posted on 2-10-2010 at 22:44


@madscientist:
The activity of hexachloropropene as a chlorinating agent comes from its three allylic chlorines, which are easily exchanged.
In making acyl chlorides, trichloroacryloyl chloride or trichloroacrylic acid is produced as byproduct- the double bond is not attacked. Only a bit of zinc chloride is required as catalyst, not 100% H2SO4.
Trichloroethylene doesn't have allylic chlorines, so you seem to need a different catalyst for it to react with carboxylic acids.
The high-boiling trichloroacrylic acid (and chloride) would probably be easier to separate from the product acyl chlorides than the chloroacetyl chloride and acid obtained as byproduct when using trichloroethylene.
Is trichlorethylene even still available today? It's banned here, even perchloroethylene isn't available OTC.




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[*] posted on 3-10-2010 at 11:46


TCE is still available - I saw a gallon of it sitting on the shelf at the hardware store recently.

I see that route as being more useful just for acetyl chloride: TCE + acetic acid --(sulfuric acid)--> chloroacetyl chloride + acetyl chloride. Boiling acetyl chloride is 50.9C, chloroacetyl chloride is 107C. But other acyl chlorides it certainly won't separate nearly as cleanly as with your proposed method - didn't mean to discount the usefulness of what you're doing here! :)




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[*] posted on 3-10-2010 at 18:04


Are you sure about that number?...
http://www.google.com/patents/about?id=QCBqAAAAEBAJ&dq=1...




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[*] posted on 4-10-2010 at 09:28


Be careful with the synthesis of hexachloropropene. Only add the heptachloropropane in small portions as the reaction is extremely exothermic once it gets going, I used a mechanical stirrer to stir mixture. I do not have my procedure available since I am away at school, but I think that I used acetone and sodium hydroxide as the other reactants. The yield was good and the infrared spectrum looked like the commercial Aldrich material after distillation.



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[*] posted on 11-10-2010 at 19:05


Here is my synthesis of hexachloropropene.
I didn't use Mr. Gergels method with NaOH, but instead used the procedure from this patent, GB841358A, which uses an aqueous suspension of calcium hydroxide as the base.
There is no dangerous exotherm with this method, and the usage of the weaker base Ca(OH)2, with a pH of the suspension of 12, is supposed to reduce side reactions like hydrolysis that may occur when using the stronger base NaOH.
The use of Ca(OH)2 suspension for the dehydrochlorination of polychlorinated aliphatic hydrocarbons is a standard industrial procedure for the production of trichloroethylene and perchloroethylene from 1,1,2,2-tetrachloroethane and pentachloroethane, with practically quantitative conversion and yield.

Preparation

I an 500ml round-bottom flask, 65g technical calcium hydroxide (hydrated lime) were slurried into 250ml water by vigorous shaking, a stirbar was added, and the flask was clamped over a magnetic stirrer.


220g of heptachloropropane were added, and the flask fitted with a reflux condenser.


The mixture was heated with a bunsen burner, first carefully, until the heptachloropropane had all melted into a liquid lower phase.
Then the stirring was started, and the mixture brought to a reflux.
The stirrer was set to the highest possible speed that did not uncouple the stirbar. Since this is a two-phase reaction with the phases being almost insoluble in each other, the reaction can only occur at the phase interface. Powerful stirring which turns the two phases into a fine emulsion is therefore very important.


The mixture was refluxed with vigorous stirring for 6 hours (perhaps this was a mistake, as the patent specifies a temperature of 90°C, not 100°C).
After stopping the heating and stirring, a drop of phenolphthalein solution was added to see whether the aqueous phase was still alkaline.
This was done due to the questionable quality of the Ca(OH)2, which produced quite strong fizzing when a bit was added to HCl, showing considerable uptake of atmospheric CO2 during storage.
The pink color showed that some unreacted Ca(OH)2 was still present at the end of the reaction, as intended:


Now the mixture was acidified by careful addition of HCl to dissolve the residual Ca(OH)2.
As everything dissolved, the mixture cleanly split into two phases:


The lower phase was washed twice with water, then twice with Na2CO3 solution (there seemed to be some acidic impurity in the crude product that produced some CO2 in the first carbonate wash) and again with water.
It was dried with calcium chloride.
The product has such a high density that the calcium chloride floated on it:


There was some brown flocculent crud, probably from dirt in the lime, therefore the liquid was filtered.

It was then distilled in membrane pump vacuum. It went over practically completely at about 95°C. Only after nearly all of the product had distilled did the temperature start to rise, and the distillation was stopped at this point.


The product, ca. 160g of a clear, mobile heavy oil with pleasant sweetish smell:


The yield was ca. 83%, which is good, but not as high as the patent claims for this process (93%).
One possible cause could be the temperature being higher than the specified 90°C- perhaps this has caused some hydrolysis of the product to trichloroacrylic acid and its secondary products. This could be the source for the observed acidic impurity in the crude product after acidification of the aqueous layer.

If I do this synthesis again, I will try to keep the temperature at 90°C, and also use more water to slurry the Ca(OH)2- the mixture became somewhat harder to stir during the reaction, and thorough mixing could not be ensured throughout the whole reflux time.

Here is the patent that describes the preparation of acyl chlorides from the acid and hexachloropropene:
GB1320869.
If this actually works, it would be a great alternative to the usage of phosphorus chlorides- not only is hexachloropropene much easier to make, but it is also an amateur-friendly reagent because it is easy to store and handle (doesn't react with moisture, doesn't give off corrosive or choking fumes, and it even smells relatively pleasant).





[Edited on 12-10-2010 by garage chemist]




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[*] posted on 11-10-2010 at 19:56


Well done on a masterful synthesis of this exotic chemical.

I appreciate your fluent use of the bunsen burner. You show that it indeed is not outmoded by electric mantles.




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[*] posted on 17-1-2011 at 08:01


Does anyone know if a second addition of CHCl3 could be carried out on the hexachloropropene? Nonachlorobutane would yield perchloro-2-butene after dehydrohalogenation. Presumably the end reaction product would either be dichlorofumaric acid or dichloromaleic anhydride (mp. 120C) allowing for the prep of even heavier acyl chlorides than hexachloropropene.



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[*] posted on 18-1-2011 at 12:50


I was wondering that as well. If hexachloropropene can add a molecule of CHCl3, perhaps perchlorobutene can do so too,
and one could build a series of perchlorinated short-chain alkenes.

What I would also like to know is how one can add chlorine to hexachloropropene and what the properties of octachloropropane are.
Mr. Gergel claims to have done this in his book- on page 117: "I laboriously chlorinated hexachloropropene, then fluorinated the octachloropropane so produced."
I can't find any information on octachloropropane, no physical properties, nothing. Yet it must be in the literature.
Can anyone look up this substance, properties and how to make it from C3Cl6?

When perchlorobutene can be made then it could also be possible to make perchlorobutane from it, and the series of perchlorinated alkanes.




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[*] posted on 18-1-2011 at 16:06


Quote: Originally posted by garage chemist  
I can't find any information on octachloropropane, no physical properties, nothing. Yet it must be in the literature.


Take a look in the Formelregister of Beilstein.
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[*] posted on 26-1-2011 at 21:15


Anyone tried out the use of Hypochlorite on Citric acid to make 1,1,1,3,3-pentachloropropan-2-one? It is of interest as a source of 1,1,1,3,3,3-hexachloropropan-2-one, which would provide a useful protecting group and trichloroacetic acid.



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[*] posted on 10-2-2011 at 15:46


The refractive index of the hexachloropropene that I produced by the above method was measured at 1,5495.
Signa-Aldrich gives a value of 1,549, so my product seems to be of good purity.
The sample of hexachloropropene from the test batch that was not base-washed had an index of 1,5492, so there is almost no difference between the two (they smell the same as well).

Nothing done with the C3Cl6 yet, I couldn't find the time.

aliced25: Do you have a source for that reaction of hypochlorite with citric acid? I've never heard of something like that.

Pentachloroacetone is produced by chlorination of a mixture of monochloroacetone and dichloroacetone with Cl2 gas, first at RT, then later at 100°C until no more chlorine reacts. The substance is isolated by fractional distillation, bp 192°C. I have a procedure.
The last hydrogen atom does not seem to get substituted by Cl2.

EDIT: I looked up Octachloropropane.
mp: 160°C, bp: 268-269°C at 734 mmHg. Beilstein only lists the preparation from 1,2,3-trichloropropane and excess ICl3
at 200°C.
It decomposes to C2Cl4 and CCl4 when its vapor is passed through a tube at 300°C. This might be of interest for those wishing to prepare some CCl4, but it would be a somewhat lengthy and work-intensive route.

[Edited on 11-2-2011 by garage chemist]




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[*] posted on 11-2-2011 at 15:14


Quote: Originally posted by garage chemist  
EDIT: I looked up Octachloropropane.
mp: 160°C, bp: 268-269°C at 734 mmHg. Beilstein only lists the preparation from 1,2,3-trichloropropane and excess ICl3
at 200°C.
It decomposes to C2Cl4 and CCl4 when its vapor is passed through a tube at 300°C. This might be of interest for those wishing to prepare some CCl4, but it would be a somewhat lengthy and work-intensive route.


The one Ergänzungsband also referenced in the index must have more information; those basically need a trip to the university library. Polychlorinated acetones (e.g. hexachloroacetone) resulting from chlorination of citric acid and citrates references are also seen in Beilstein.
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[*] posted on 19-12-2012 at 16:11


I would like to report my successful repetition of the unsymmetrical heptachloropropane procedure. I scaled down the orgsyn procedure to handle 10g of AlCl3, which is all I had (See note below). Solvent purification was carried out much like Garage Chemist did. I used 15% aq. HCl to wash the perchloroethylene, followed by saturated NaHCO3, and finally water. The perchloroethylene was dried over CaCl2 and decanted from the dessicant for use. The chloroform was washed with one small portion of concentrated sulfuric acid (~10ml for 120g of chloroform containing 1% MeOH), then saturated NaHCO3, and then dried over CaCl2 and decanted from the dessicant for use.

A 500ml RBF with a stir bar was charged with 10.00g of anhydrous aluminum chloride, 62.00g of perchloroethylene, and 112.05g of chloroform. A Freidrich condenser with a CaCl2 guard tube was mounted on the flask, and a propane torch flame was swept along the bottom of the flask for perhaps 10 seconds to initiate the reaction. It became hot, much more so than the torch alone would have made it, but on this scale, it did not boil. Without additional external heating, it cooled down, and the flask was heated externally with a hotplate and stirring. The mixture began as an orangey-brown and darkened to a deep brown during 16 hours of very gentle reflux.

The cooled reaction mixture was poured over ~250ml of small chunks of ice in a beaker. The hydrolysis of the AlCl3 was not nearly enough to melt all the ice. Some waxy solid was observed to separate out in the chloroform phase and float on top of the aqueous phase. The mixture was warmed in a water bath to melt the ice. Small amounts of chloroform were added to dissolve as much waxy material as possible and it was loaded into a separatory funnel. The organic phase was dark brown and the aqueous phase, cloudy and tan. The organic phase was drained off after shaking vigorously and allowing to settle, and the aqueous phase discarded. A second wash with water provided similar phases, and a layer of "gunk" that was initially assumed to be emulsion on the aqueous/organic interface. Thus, it was transferred with the organic phase. A third wash provided an almost clear aqueous phase, but a dark brown organic phase. An attempt was made to let the "gunk" settle like an emulsion after draining off the bulk of the organic phase. A small layer of organic phase separated and was drained off several times, leaving the real culprit, a pasty mass of tar-stained aluminum hydroxide entrapping small amounts of organic phase. The organic phase was dried over a liberal amount of anhydrous CaCl2 for several hours, decanted, and slowly fractionated through a Vigreux column to recover chloroform. Since the amount added to dissolve product had not been quantified, I can not comment on recovery.

The residue froze on standing overnight and was set up for vacuum distillation with an air condenser and my teflon diaphragm pump (~30torr). On applying vacuum, some bubbles formed (residual chloroform) and some condensate appeared in the stillhead, condenser, and vacuum adapter between 25 and 30C that is very likely traces of unreacted perchloroethylene. The apparatus was disassembled and this material removed mechanically and by heating with a blowdryer.

On resuming vacuum distillation, material began condensing in the stillhead as early as 80C, but did not begin to pass over until the stillhead reached 125C, which was rapidly rising in temperature at the time. The vast bulk of the product passed over between about 133C and 137C, the latter being reached only with material condensing far up into the thermometer adapter (likely over-immersing the thermometer and giving a falsely high reading). The stillpot contained a very small amount of black residue. The product was a water-clear heavy, somewhat viscous (supercooled) liquid. On adding a seed crystal scraped from the condenser, the entirety of the product crystallized to a white solid in seconds, becoming quite warm in the process. The product weighed 95.76g, an 89.8% yield from perchloroethylene.


Note) This material had been prepared from Al metal in DCM, with catalytic I2, into which dried chlorine gas was fed slowly. The resulting product was freed of DCM (now dark brown from polycondensation products), and sublimed at atmospheric pressure with a CaCl2 guard tube, leaving a charred residue. The bulk material was faintly yellow and some of the last material to sublime was orangey. This is presumably from ICl3.




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[*] posted on 12-5-2013 at 21:59


Quote: Originally posted by garage chemist  


aliced25: Do you have a source for that reaction of hypochlorite with citric acid? I've never heard of something like that.



Sorry I only just noticed you asked, I cited the article (and uploaded it) on this thread (http://www.sciencemadness.org/talk/viewthread.php?tid=8378&a...). I'd personally be very interested to see if it can be made to work, as it would presumably break down to either dichloromethane & trichloroacetic acid or more probably chloroform and dichloroacetic acid when reacted with a-amino acids.

(It would be a very interesting thing to use SbF5 on though, as a potential precursor to TFA)

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Diagnosing Consecutive Reactions of Hypochlorite: pH and Oxidative Decarboxylation/Halogenation

Benjamin J. Gilliotte, Cynthia L. Sanders, L. Kevin Wall & Robert G. Landolt

J. Org. Chem
Vol.51(16) 1986, pp.3233-3234
DOI: 10.1021/jo00366a040

Abstract

In addition to the classic “haloform” reaction of methyl ketones,’ trihalomethanes result readily from hypohalous acid/hypohalite reactions of certain carboxylic acids and amino acids, as well as from activated B-diketones, meta-dihydroxylated aromatics, and quinones. With some of these substrates, halogenation is but one component of a complex series of competitive and/or consecutive reactions.

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Note they cite several "old" journal articles and they also test it. Now, they don't say they succeeded, but they don't say they disproved it either. Someone with analytical equipment trying this would be a godsend. Trying the fluorination with SbFx salts would also be interesting to read about.

Attachment: Gilliote.etal.Diagnosing.Consecutive.Reactions.of.Hypochlorite.pH.and.Oxidative.Decarboxylation.Halogenation.pdf (284kB)
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There is also Dr Lassar Cohn, Laboratory Manual of Organic Chemistry, MacMillan & Co, London, 1895 (https://ia700306.us.archive.org/35/items/alaboratorymanu00la...) p.182

Quote:
To secure the exposure of a large surface to the action of chlorine a device like that used by Cloez (Bull. Ch. 39, 636) may be employed. He dissolved citric acid in one and a half times its weight of water, and allowed this solution to drop onto pieces of pumice in a vertical cylinder, while a stream of chlorine passed upwards to meet it. The product of the action was pentachloroacetone


[Edited on 13-5-2013 by aliced25]




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UnintentionalChaos
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[*] posted on 19-1-2014 at 13:07


It only took me 13 months, but I finally got around to dehydrohalogenating my unsym-heptachloropropane. My source of calcium hydroxide was "Mrs. Wages Pickling Lime" obtained from a Walmart several years ago. The packaging is foil lined and addition of acid showed no appreciable effervescence.

To a flask containing 93.76g of heptachloropropane (0.329mol) was added 24.38g of Ca(OH)2 (100% excess) slurried in 75ml of distilled water. A large stirbar was added. The flask was warmed until the hepta melted forming a dense lower phase. Stirring was turned nearly to maximum and the two phases mixed into an emulsion. The mixture was initially raised to reflux, then heating was adjusted so it remained at 93+/-3C for 6 hours with stirring. Upon cooling, the mixture thickened considerably. 35ml of 31.45% HCl was added and with shaking, all remaining calcium hydroxide dissolved. The reaction mixture split cleanly into a pale yellow-green upper aqueous phase and a colorless, but slightly hazy lower phase. The lower phase was isolated with a seperatory funnel, and shaken with small portions of saturated aq. sodium bicarbonate until no more CO2 evolution was noted.

The lower phase was then washed with 40ml of saturated NaCl brine, but this did not clear the organic phase as hoped. The still-hazy organic phase smelled considerably like heptachloropropane and was left to stand over anhydrous CaCl2 overnight, after which it had cleared. The liquid was decanted from the CaCl2 into a 250ml RBF. A stirbar was added and the product distilled at ~30torr in a short path stillhead. Product passed over smoothly from 100-103C, rising relatively quickly to 105C when only a small amount of liquid remained in the stillpot. At this point, the vacuum was cut off.

The product smells faintly of heptachloropropane and is a dense, colorless liquid weighing 65.10g (79.5% yield) I suspect that larger batches would improve yield by reducing mechanical losses and may make fractionating the final product worthwhile.




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[*] posted on 21-1-2014 at 23:38


Inorganic Synthesis reports many uses for hexachloropropene such as the synthesis of tungsten hexachloride, which is a really nice purple-black colored crystalline compound. It is very hydroscopic and needs to be prepared under nitrogen using Schlenk glassware. The synthesis involves refluxing WO3 with an excess of the hexachloropropene for several hours, cooling, filtering with a Schlenk filter and washing with CH2Cl2. I would highly recommend the synthesis for anyone who likes working with coordination complexes, especially those involving high oxidation state tungsten.

For a detailed procedure, see Inorganic Syntheses Volume 9, page 133. VCl3, NbCl5, MoCl5, can be prepared analogously. 10 mL of hexachloropropene are used for each gram of metal oxide, not exactly cheap with respect to the hexachloropropene.




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