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Author: Subject: properties and reactions of tetrachloroethene and trichloroethene
woelen

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properties and reactions of tetrachloroethene and trichloroethene

I now have two new chemicals added to my lab. I could obtain them cheaply: 1 liter of CHCl=CCl2 and 1 liter of CCl2=CCl2, trichloroethene and tetrachloroethene (some people call them trichloroethylene and tetrachloroethylene). Both are colorless liquids with a sweetish smell. The tri is much more volatile than the tetra. Both chemicals have a warning on the bottle that they might cause cancer, so I am careful not to inhale too much of the vapor.

I have done quite a few experiments and I found that these chemicals are much less reactive than I expected. I also have searched sciencemadness extensively and although quite some information is scattered over the forums, there is not a nice thread which covers these chemicals in detail.

I found that adding a halogen to these chemicals across the double bond is very slow or does not work at all. I tried with chlorine and with bromine. Is this due to steric hindrance, because of the large chlorine atoms around the small C-atoms, partially covering the double bond between the C-atoms?

I have done some research on oxidation and some sources are talking about formation of dichloroacetic acid from these chemicals when oxidized. I found that oxidizing them from aqueous solution, combined with hefty shaking to keep good contact between the two liquid phases hardly works. It is very very slow, even when the liquid is heated.

Next, I tried a more vigorous oxidizer. I dissolved some CrO3 in warm (not hot!) conc. H2SO4. I added this to some CCl2=CCl2. Nothing happens! Some of the CrO3 goes to the organic phase, but no visible reaction occurs. When I did the same with CHCl=CCl2, then a reaction starts. The liquid starts foaming a little and the red color of the haxavalent chromium changes to green. A colorless gas is produced. When a lot of warm water was added, white fumes (HCl?) could be observed. What could this colorless gas be. There was a fairly strong spicey smell (not really unpleasant) and when I smelled that, I immediately went away, not knowing what it was and not wanting to take the risk of poisoning myself. Cleaning up the stuff I did later, when there was no smell anymore. This gas cannot be dichloroacetic acid. Given the structure of CHCl=CCl2 and assuming breakdown of the C-C bond, I only can imagine the gas to be CO2 or (bad bad!!) COCl2. That's why I immediately left the lab when I smelled something peculiar.

Are there any people over here, who know of interesting reactions with CHCl=CCl2 and CCl2=CCl2? Also getting more insight in what I did would be welcome.

[Edited on 17-7-12 by woelen]

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Lambda-Eyde
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I know you're not into organic chemistry, but tetrachloroethylene can be reacted with chloroform in the presence of AlCl<sub>3</sub> to give unsym. heptachloropropane, which in turn can be dechlorinated to hexachloropropene (which is a much more interesting chemical than it sounds).

Trichloroethylene is a great solvent for grease, but I don't know of any reactions it takes part in...

[Edited on 17-7-2012 by Lambda-Eyde]

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woelen

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Yes, I read garage chemist's thread about heptachloropropane. I consider trying that on a smaller scale, let's say a few grams of reagents. This seems doable with standard glassware and equipment. I have all reagents needed for that.

[Edited on 17-7-12 by woelen]

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Mailinmypocket
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Also, sulfur dissolves nicely in trichloroethylene
kavu
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With a quick search through reaxys:

Trichloroethene is mostly used as an electrophile in substitutions, common example is williamson ether synthesis. Also some radical type chemistry is used to form dichlorovinyl groups. Interesting compound, 1,2-dichloroacetylene, can also be prepared by heating trichloroethane with sodium hydroxide in DMSO/Et2O (Journal fuer Praktische Chemie (Leipzig), 1989 , vol. 331, # 1 p. 145 - 148). Full hydrolysis of trichloroethane with NaOH yields glycolic acid (Schmerling,L. Journal of Organic Chemistry, 1975 , vol. 40, p. 2430 - 2434).

Similar substitution type chemistry with S, N and O has been investigated for tetrachloroethene as well. For both compounds some pericyclic reactions leading to highly chlorinated cyclobutane rings has been published (for example Nonoyama, Shinji; Yonezawa, Noriyuki; Saigo, Kazuhiko; Hirano, Tsuneo; Hasegawa, Masaki
Chemistry Letters, 1987 , p. 487 - 490) . Interesting thing is nitration of these highly halogenated ethenes (Biltz Chemische Berichte, 1902 , vol. 35, p. 1550). Some coupling reactions with Pd and Ni have also been investigated, this seems to be a viable route to alkyne coupled aromatics.

A really interesting compound, sodium tetrathiooxalate, can be prepared as described in Kuropatov, Viacheslav; Klementieva, Svetlana; Fukin, Georgy; Mitin, Aleksander; Ketkov, Sergey; Cherkasov, Vladimir; Abakumov, Gleb; Budnikova, Yulia Tetrahedron, 2010 , vol. 66, # 38 p. 7605 - 7611.

Reason for the reduced reactivity of the pi-bond results from electronegativity of chlorine. Nucleophilicity of the double bond is reduced due to chlorine, resulting in a HOMO bulged at chlorine ends ( http://ed.augie.edu/~sghussai/homo.jpg ) rather than high electron density at the double bond.
 Quote: When I did the same with CHCl=CCl2, then a reaction starts. The liquid starts foaming a little and the red color of the haxavalent chromium changes to green

Heating trichloroethene with sulfuric acid (and a bit of water) leads to formation of chloroacetic acid. Under oxidative conditions the alcohol formed after initial protonation and addition of water might get oxidized or coupled to another trichloroethene. A reasonable mechanism could be drawn at least for the formation of dichloroacetylchloride. Similar reaction (at least through the hypothetical mechanism) is not possible with tetrachloro compound as there's only one hydrogen. This would also match up with the smell as the acid chloride has a more pungent odor than chloroacetic acid and is readily hydrolyzed by moisture. (Though both do stink a LOT)

[Edited on 17-7-2012 by kavu]
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You could try the angry genie, dichloroacetylene:

I also did some work with chlorinating tetrachloroethylene and ended up with some very large shard-like solids that crashed out (hexachloroethane I presumed) although I was going for carbon tetrachloride, I never got around to fractionating. I used a UV source to try to push things along and to make the radicals to cleave the C-C bond.

Finally I purposefully did some small scale experiments with oxidants (chromic acid solutions) to attempt to make small quantities of phosgene (older preparations for phosgene from chloroform using chromic acid detail that other completely halogenated hydrocarbons can be utilized). The solutions built pressure quickly but I had no means to test for phosgene, at the time so it was likely an experiment that I should have never carried out.

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woelen

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I tried the angry genie, dichloroacetylene. I now remember that I tried before. I indeed had a little amount of CHClCCl2 some years ago, but all of it has evaporated over the last two years. The stuff is amazingly eager at leaving its container. The new CHClCCl2 I now have is stored in a good laboratory glass bottle and I hope that this lasts longer than a few years.

Now I had a really impressive experience

I took 2 ml of CHCl=CCl2 and put this in a test tube.
* I added so much solid NaOH that appr. 75% of all liquid was filled up with the granules of NaOH. After this, nothing happened. The NaOH goes to the bottom and does not react at all.
* I added a small amount of water and started shaking strongly. The water mixes with the NaOH and dissolves most of it. Finally, I obtained a two-layer system, the top layer being the CHCl=CCl2 and the bottom layer being the saturated (and quite hot) solution of NaOH, with a small quantity of solid NaOH at the bottom. Still no reaction occurs.
* I carefully added appr. 1 ml of DMSO. This DMSO remains floating on the CHCl=CCl2, but there is no sharp interface between the layers. After gentle shaking, all of the DMSO dissolves in the CHCl=CCl2 layer.
* Next, I started shaking somewhat more vigorously. The layer of CHCl=CCl2/DMSO still remains on top, but the interface with the aqueous NaOH layer becomes much more ragged and it looks as if the aqueous NaOH layer mixes somewhat with the organic layer. After a few minutes I heard gentle crackles, one crackle each second or so. The color of the organic layer slowly changed from colorless to brown. The brown material is a flocculent solid, dispersed in the liquid. The liquids at this point were only luke-warm, most of the heat of dissolving the NaOH in water was gone already.
I kept on swirling gently, and slowly, a white dense fume was produced in the test tube, which remained above the now brown/olive green organic layer. Suddently there was an orange flame in the test tube and much more white fumes were produced. The soft crackles still could be heard, each second or so.
* I decided to heat the liquids in the test tube somewhat. I only heated slightly, maybe to 60 C or so, just a little too hot to keep it touched with bare hands for a long time. The heating leads to somewhat more frequent crackling and the organic layer becomes totally opaque brown with an olive green hue. Suddently at the open end of the test tube a little dark orange/red flame appeared and tremendous amounts of soot were produced. The flame slowly moved inside the test tube, producing even more soot. The flame disappeared when it reached the surface of the liquid. This process repeated many times. Every half a minute or so, a new small flame appeared at the open end of the test tube, which went inside and then disappeared. Each time a fantastic black cloud of soot was produced.
* Finally, I decided to quit the experiment and filled the test tube with water such that the liquid is just under the rim. Many small bubbles were produced, and as soon as these bubbles reached the surface of the water, they burst out in red/orange micro-flames, each micro-flame producing a puff of black soot.

This is a great demo, but the amount of soot is incredible. It best can be done outside. The smell of the soot is bad. It has no typical fire-smell, but it has a bad chemical odor. I only smelled it slightly, but it is bad, really bad.

My next experiment will be to do this experiment with a test tube, which is stoppered and collecting the gas bubbles in another test tube under water (the first half I will not collect, because of mixed air). Then I'll do some experiments with the pure gas. It is very flammable, but an interesting experiment might be to ignite the gas by just holding a warm (not hot) glass rod or metal stick in the gas.
Unfortunately, I hardly find any useful info about the gas. I know it is colorless and toxic and pyrophoric at temperatures just above room temperature, but I can't find much more info on it. Is it safe to collect the gas in a glass tube? Or is there a chance of explosion, sending splinters of glass around the place?

I think that the DMSO does not react in this experiment, but that it facilitates more intimate mixing of the hydroxide and the trichloroethylene. Without the DMSO there is a sharp boundary between the aqueous layer and the organic layer and there may be insufficient contact.

[Edited on 20-7-12 by woelen]

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jon
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you know this might have some utility now they found chloroacetates and phenylacetates to cure cancer
but its not being developed because it can't be put under patent and the $they make of you off the last 6 months of your life is like more than you spend on average during your life. http://truedemocracyparty.net/2012/03/cancer-cures-dca-thc-n... Give me librium or give me meth! Patrick Henry.... edgecase Hazard to Self Posts: 65 Registered: 4-11-2011 Location: Ontario, Canada Member Is Offline Mood: Viscous I suspect it is being dumped from the market, due to the mandate to reduce CFCs. Many local sources of "surplus" goods have automotive brake parts clearer for heavy discounts. The best deal I've found is 600g for$3.99 CAD, pure tetrachloroethylene and CO2 as propellant. Other brands were 535 or 500g with other ingredients (some had trichloroethylene), and more expensive.

My first thought was that this would be a useful solvent for chlorinations, where the product can be distilled off and the high-boiling solvent (BP 121.1 deg C) can be used as a "chaser".

[Edited on 2012-7-22 by edgecase]
garage chemist
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http://www.versuchschemie.de/htopic,6344,hexachlorethan.html
http://www.versuchschemie.de/htopic,12817,hexachlorethan.htm...
http://forum.lambdasyn.org/index.php/topic,1645.0.html
I have successfully done this myself once, though I do not know the optimal conditions for this reaction. Gentle heating and refluxing have been tried, with and without UV irradiation, and there is always quite some chlorine escaping, though product is being formed at a sufficient rate in all cases.
Addition of AlCl3 as catalyst, as described in the last link, is not recommended since it may cause side reactions (e.g. addition of C2Cl6 to C2Cl4 giving decachlorobutane- this is only a hypothesis by me) and darkening of the reaction mixture.
The chlorination has to be stopped before 60% of the C2Cl4 have reacted because the reaction mixture would otherwise solidify from precipitated product even at its boiling point.
The crude C2Cl6 can be obtained pure by repeated recrystallization from 95% Ethanol.
It is a solid which sublimates without melting when heated at atmospheric pressure. Its melting point must be determined in a sealed capillary.

Chlorination of trichloroethylene would give pentachloroethane, a liquid with relatively high boiling point. I was once interested in the possible application of this as a solvent for the chlorination of red phosphorus to PCl3 since the high b.p. would make the distillative separation from PCl3 easy. CHCl3 and CCl4 cannot be used because of the close boiling points, and also because red P only forms PCl5 when chlorinated in these solvents due to insufficient temperature.

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edgecase
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 Quote: Originally posted by garage chemist Tetrachloroethylene readily adds chlorine to form hexachloroethane:

So are you suggesting that it would *not* be a suitable solvent for chlorinations, due to it's own reactivity with Cl(g) ? What about for bromination?
garage chemist
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I don't think it could be used for chlorinations, the addition is quite rapid even at room temperature. No idea about brominations, though I remember BromicAcid has once reacted bromine with tetrachloroethylene at room temperature and obtained a waxy solid.

Pentachloroethane derived from trichloroethylene, on the other hand, would be an interesting and attractive solvent for chlorinations. Radical substitution of the last hydrogen atom should be extremely slow even at reflux temperature from what I've read.

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 Sciencemadness Discussion Board » Fundamentals » Organic Chemistry » properties and reactions of tetrachloroethene and trichloroethene Select A Forum Fundamentals   » Chemistry in General   » Organic Chemistry   » Reagents and Apparatus Acquisition   » Beginnings   » Responsible Practices   » Miscellaneous   » The Wiki Special topics   » Technochemistry   » Energetic Materials   » Biochemistry   » Radiochemistry   » Computational Models and Techniques   » Prepublication Non-chemistry   » Forum Matters   » Legal and Societal Issues