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Author: Subject: synthesis of methyl chloride with phosphorus trichloride
dactyl
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[*] posted on 21-12-2016 at 10:10
synthesis of methyl chloride with phosphorus trichloride


Is it reasonably possible to synthesize pure methyl chloride through this reaction:
PCL3 + CH3OH = CH3CL + H3PO3
Metacelsus
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[*] posted on 21-12-2016 at 12:55


First, that equation isn't balanced.

Yes, phosphorous trichloride converts alcohols to alkyl chlorides, so this would work. However, I think you'll also get significant amounts of trimethyl phosphite produced.




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AJKOER
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[*] posted on 21-12-2016 at 18:22


A small scale impure preparation would be to light a methane torch and insert the burning CH4 flame into a generated stream of chlorine gas. Target reaction:

CH4 + Cl2 → CH3Cl + HCl

so the gas stream should contain equal volumes of methane and chlorine.

Unfortunate side reactions with an incomplete methane combustion in chlorine or excess generated chlorine:

CH3Cl + Cl2 → CH2Cl2 + HCl
CH2Cl2 + Cl2 → CHCl3 + HCl
CHCl3 + Cl2 → CCl4 + HCl

[Edit] Per comments relating to possible formation of some COCl2, which is commercially produced by passing purified carbon monoxide and chlorine over a catalyst (porous activated carbon), I would recommend scrubbing the exit gases with water removing any COCl2 and HCl as:

COCl2 + H2O → CO2 + 2 HCl

Decant organic products and further separate products by distillation.

Support, as I have not attempted this experiment: Per Wikipedia on CH2Cl2 (link: https://en.m.wikipedia.org/wiki/Dichloromethane), on COCl2 (link: https://en.m.wikipedia.org/wiki/Phosgene ) and CH3Cl (link: https://en.m.wikipedia.org/wiki/Chloromethane ), to quote from the latter:

"A smaller amount of chloromethane is produced by heating a mixture of methane and chlorine to over 400 °C (752 °F). However, this method also results in more highly chlorinated compounds such as dichloromethane, chloroform, and carbon tetrachloride and is usually only used when these other products are also desired."

[Edited on 22-12-2016 by AJKOER]

[Edited on 22-12-2016 by AJKOER]
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clearly_not_atara
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[*] posted on 21-12-2016 at 18:44


I'm sort of concerned that in the chlorination of a methane flame in air it is possible that some CO will be produced and react with available Cl2, resulting in a straw-like odor and breathing difficulties.

I think the HCl/ZnCl2/MeOH method is probably fine. If that doesn't cut it then a chlorinating agent like PCl3 or SOCl2 or the system SO2Cl2/HOBt
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[*] posted on 21-12-2016 at 19:59


Quote: Originally posted by AJKOER  
A small scale impure preparation would be to light a methane torch and insert the burning CH4 flame into a generated stream of chlorine gas. Target reaction:

CH4 + Cl2 → CH3Cl + HCl

so the gas stream should contain equal volumes of methane and chlorine.

If you have a flame, aren't you also feeding some oxygen into the mix? There are more possible side reactions than you have listed. But maybe a setup where you switch O2 for Cl2 would work? That strikes me as feasible.




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[*] posted on 21-12-2016 at 21:01


I'm guessing you could actually produce hundreds of liters of chloromethane and hydrochloric acid this way with a tube furnace?



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[*] posted on 22-12-2016 at 04:13


Quote: Originally posted by Metacelsus  
First, that equation isn't balanced.

Yes, phosphorous trichloride converts alcohols to alkyl chlorides, so this would work. However, I think you'll also get significant amounts of trimethyl phosphite produced.


Trimethyl phosphite, methyl dichlorophosphite, and dimethyl chlorophosphite would all form as intermediates. However, the methyl groups on these will undergo Sn2 substitution by Cl- in the reaction mixture, forming a strong P=O bond and ultimately forming H3PO3. The methyl chloride (gas, bp -24.2 C) will be easily separated from this mixture.

On a separate note, I think we're underestimating the challenges of high temperature chlorination. It truly is mad science. I recommend using a brick furnace, but it may well corrode before your reaction is complete. And for sure: CH3Cl is going to be more reactive than CH4 to radical chlorination. Selectivity is going to be an issue, regardless of your stoichiometry. I expect a hot, fuming, poisonous, ozone-depleting mess from this.

EDIT:

Instead of a 400 C furnace, why not use a tungsten lamp? Fill a vessel with methane and chlorine and irradiate it. UV light will weaken the Cl-Cl bond for homolytic cleavage, allowing Cl radicals to tear up C-H bonds. This sort of photochemistry is not economical on scale, but I see no reason why it can't work on lab scale.

[Edited on 12-22-16 by DDTea]




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[*] posted on 22-12-2016 at 04:20


I have tried the chlorination of alkanes (methane, propane, butane) from a flame. It does not work. The alkane starts burning with an incredibly dirty sooty dark orange flame as soon as chlorine comes in contact with the burning gas. The gas/smoke mix escaping from this is very pungent and besides HCl contains a lot of soot, and some stuff which has a strong and acrid smell, but which definitely is not cleanly chlorinated alkane, but partially burned material.

It is an interesting experiment in itself, especially the flame color is weird, it is really a dark flame, dark orange, probably because of the intense sooting of the flame.

[Edited on 22-12-16 by woelen]




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[*] posted on 22-12-2016 at 04:32


I think that if you inhaled the reaction mixture from methane and chlorine, the HCl would kill you before any COCl2 got a chance.
Is there anyone to whom that distinction matters?

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[*] posted on 22-12-2016 at 13:35


Comments from US Patent 1,891,859, "Production of carbon and hydrogen chloride", appear to support Woelen's observations. To quote:

"It is known that gaseous hydrocarbons such as methane, ethylene or acetylene may react with chlorine to give, among other products, carbon and hydrogen chloride, whether a mixture of the reactant gases be exploded or whether one of the gases be caused to burn in the other. This invention consists in the production of carbon and hydrogen chloride from hydrocarbon-containing gases or vapours, such as coal distillation gases, natural gas, cracking still gas, vaporised hydrocarbons and the like by causing the gases to combine in a flame burning in an atmosphere of air.
According to one form of the invention chlorine is burned in an envelope of hydrocarbon-containing gas surrounded by air, so that the outer surface of said envelope burns in air."

Link: https://www.google.com/patents/US1891859

[Edited on 22-12-2016 by AJKOER]
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[*] posted on 22-12-2016 at 18:48


On DDTea's post

Wouldn't a pressure-equalizing funnel be needed to collect the methyl chloride gas perhaps with the gas being collected from to the top of the pressure-equalizing funnel?
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[*] posted on 22-12-2016 at 18:51


correction: from the top
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[*] posted on 22-12-2016 at 21:47


Ullmann's Encyclopedia of Industrial Chemistry says that the process is conducted at 350-450 C. A Bunsen burner flame might be too hot. Also, any heating is supposed to take place before the chlorine and methane are mixed.


Attachment: ChlorinatedHydrocarbons.pdf (2.3MB)
This file has been downloaded 2820 times



[Edited on 23-12-2016 by JJay]




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[*] posted on 23-12-2016 at 03:04


I think the process of making CH3Cl from betaine hydrochloride looks interesting.
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[*] posted on 23-12-2016 at 10:31


Per Wikipedia on hypochlorous acid (link: https://en.m.wikipedia.org/wiki/Hypochlorous_acid ), to quote:

"HClO reacts with alkanes to form chloroalkanes and water, illustrated here by its reaction with methane:

CH4 + HClO → CH3Cl + H2O"

Per another source: Theoretical and Experimental Chemistry, September 2012, Volume 48, Issue 4, pp 227–232, "DFT analysis of the mechanism for the gas-phase chlorination of methane in the HOCL–H2O system", by S. L. Litvinenko, to quote:

"The DFT method was used to study gas-phase chlorination reactions CH4 + HOCl (1) and CH4 + HOCl + H2O (2). These reactions entail singlet–triplet (s→t) preactivation of HOCl and terminate in a reverse t→s transition. At 298 K, the barrier ∆G≠t for reaction (2) is higher than for reaction (1) by 1.5 kcal/mol, while the singlet transition state TSs-2 lies higher than TSt-2 by MG= 21.4 kcal/mol."

Link: http://link.springer.com/article/10.1007/s11237-012-9265-7

Apparently, a gas phase reaction, but I can also see how this reaction could proceed with hypochlorous acid and a small excess of methane in UV light via a radical pathway as well. For example, some expected reactions:

HOCl + hv → .OH + .Cl
.OH + CH4 → H2O + .CH3
.Cl + CH4 → HCl + .CH3
HCl + HOCl = Cl2 + H2O
.Cl + .Cl → Cl2
Cl2 + .CH3 → CH3Cl + .Cl
.Cl + .CH3 → CH3Cl
......
Source: See http://www.chemguide.co.uk/mechanisms/freerad/ch4andcl2tt.ht... and also Table l in "Large losses of total ozone in Antarctica reveal seasonal ClOx/NOx interaction", by JC Farman, link: http://eesc.columbia.edu/courses/v1003/readings/Farman.etal....

One could also combine approaches by adding a uv photolysis to the gas phase reaction.
-----------------------------------------------------------------------------------------

Note, one can prepare HOCl by the action of CO2 on a mixture of aqueous NaOCl (chlorine bleach) and CaCl2. Remove the CaCO3 precipitate leaving dilute HOCl and NaCl. One way to remove the NaCl from the HOCl and concentrate the hypochlorous acid is to distill half of the volume of the HOCl/NaCl and discard the remaining half. Per Watts Dictionary of Chemistry (please see comments and links in a prior SM thread at https://www.sciencemadness.org/whisper/viewthread.php?tid=17... ), the more volatile than water HOCl/Cl2O is driven largely off first, so the concentration of the hypochlorous acid is nearly doubled and the NaCl is removed. Note, concentrated HOCl is increasingly unstable and should be cooled, free from strong light and used promptly.

[Edited on 24-12-2016 by AJKOER]
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[*] posted on 24-12-2016 at 11:31


Interestingly, I can across a reference, to quote:

"CH4 + Cl2→ HCl + CH3Cl

This reaction does not ordinarily occur in aqueous solutions and therefore is of little significance in environmental engineering and science."

Link: Page 220 at https://www.google.com/url?sa=t&source=web&rct=j&...

But, as aqueous chlorine implies HCl and HOCl, my gas phase reference to reactivity of hypochlorous acid on methane is perhaps more correctly addressed as a solely gas phase event.

I also recall reading that the optimal pH for the HOCl species is around pH 6.
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