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Author: Subject: The Electrolytic Chlorination of Ethanol
Metacelsus
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[*] posted on 5-5-2014 at 13:01
The Electrolytic Chlorination of Ethanol


87.6 ml (1.50 mol) of absolute ethanol were added to 400 ml of 12 M hydrochloric acid. The mixture was electrolyzed in a cell with a stir bar at 5 V and 1.2 A for 165 hours. The current density was 24 mA/cm2. Platinum electrodes were used for both anode and cathode.

I thought this reaction would produce chloral, which I planned to isolate by fractional distillation. However, I wound up getting chloroform instead (b.p. 59-65 oC). The yield of chloroform was 68 ml (0.84 mol), density 1.49 +/- 0.02 g/cm3. Interestingly, the reaction mixture only had one phase. A possible explanation (although I doubt it) could be that the reaction did form chloral, but that it decomposed upon heating.




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[*] posted on 5-5-2014 at 14:53


Two German patents on the chlorination of ethanol to produce chloral are in the attached archive :P There's threads on producing chloral on SM already ;) Do you have any references supporting your hypothesis :cool:

[Edited on 5-5-2014 by leu]

Attachment: ChloralPatents.zip (448kB)
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[*] posted on 5-5-2014 at 18:18


With platinum electrodes, it's hard to say what exactly you were making. Why not use graphite? That may be better suited for both cathode and anode in HCl solution. If you use a platinum cathode, it might be necessary to use a divided cell, where the alcohol is being added to the anode compartment. This is because platinum happens to be very good at reducing things under cathodic conditions. In other words, what is happening at the anode could be undone at the cathode, with chloroform building up from a side reaction. But I'm not an organic chemist, so your mileage may vary on that one.
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[*] posted on 5-5-2014 at 23:16


I'm amazed Chedite that 50cm^2 of platinum electrodes with 5V across it in a 12M HCl solution only drew 1.2A. Any ideas why the low current?



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[*] posted on 6-5-2014 at 02:45


Results are interesing, but not very surprising.
Iodoform and bromoform can be obtained in similar way - electrochemical iodination/bromination of ethanol.
Apparently chlorination also may give chloroform.
Explanation of this reaction is not easy, especially that homogenous mixture is obtained. Possibly, formed chloral is oxidized to chloroform during heating.
I think that results of this experiment are strongly conditions-dependent (concentrations, temperature, current density, amount of electricity... etc)




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[*] posted on 6-5-2014 at 04:47


@DeltaH
I didn't directly measure the voltage across the cell. My power source (a wimpy cellphone charger) may have been outputting less than the rated 5V.

I usually use a rectifier on a variac (this can easily give up to 20A) but I had burned out the rectifier and was waiting for a new one to arrive.

Maybe I will try a graphite cathode to prevent reduction (if I ever do this again).




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deltaH
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[*] posted on 6-5-2014 at 05:22


Ah okay... that explains it. The devil's in the details :D

Yeah I also think graphite should work much better! Will make a mess of carbon powder though, but not an issue really.

[Edited on 6-5-2014 by deltaH]




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[*] posted on 6-5-2014 at 05:42


Quote: Originally posted by Cheddite Cheese  
I thought this reaction would produce chloral, which I planned to isolate by fractional distillation. However, I wound up getting chloroform instead (b.p. 59-65 oC). The yield of chloroform was 68 ml (0.84 mol), density 1.49 +/- 0.02 g/cm3. Interestingly, the reaction mixture only had one phase. A possible explanation (although I doubt it) could be that the reaction did form chloral, but that it decomposed upon heating.

Chloral can decompose to chloroform, but this occurs in basic media. I can't see any mechanism for such a thing occurring in acidic media. You did distill the reaction mixture directly?




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[*] posted on 6-5-2014 at 10:58


Yes, I did.



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[*] posted on 13-6-2014 at 09:12


Maybe the chloral is oxydised to trichloroacetic acid and this afterwards lead to decarboxylation...
CCl3-CH=O + O(2-) --> CCl3-CO2H +2 e(-)
CCl3-CO2H --> CCl3-CO2(-) + H(+)
CCl3-CO2(-) --> CCl3-CO2° + 1 e(-)
CCl3-CO2° --> CCl3° + CO2
CCl3° + 1e(-) --> CCl3(-)
CCl3(-) + H(+) --> CHCl3

It is strange that in the mix you don't find the dimeric termination product:
2 CCl3° --> Cl3C-CCl3 (hexachloroethane)




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[*] posted on 13-6-2014 at 10:08


Well, I wasn't looking for it. Anyway, decarboxylation of trichloroacetic acid is plausible, according to this paper:



Attachment: V70N02_097.pdf (601kB)
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