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Topic: Synthesis of trichloromethanesulfinyl chloride  (Read 17449 times)

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Offline Babcock_Hall

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Re: hydrogen sulfide as a reactant
« Reply #15 on: February 06, 2016, 06:32:12 PM »
That's an interesting suggestion.  If one could create the [CCl3]- anion and react it with thionyl chloride, one would have the desired compound.  I don't know how practical this is offhand.

Offline BRSM

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Re: hydrogen sulfide as a reactant
« Reply #16 on: February 07, 2016, 10:34:51 AM »
There's no need to do put yourself through H2S. There's actually a very reasonable synthesis of trichloromethanesulfinic acid from trichlorobromomethane and sodium dithionite, two very common and safe reagents, that was reported in the literature a couple of decades ago (Inorg. Chem. 1992, 31, 492-494). From that, the chlorination with thionyl chloride that you propose seems reasonable and well precedented (there's a procedure in the same paper).

Doesn't seem like more than a day's work for both steps. Have fun!




Offline Enthalpy

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Re: hydrogen sulfide as a reactant
« Reply #17 on: February 07, 2016, 01:13:34 PM »
With thionyl chloride, I had imagined to react chloroform, both in gas phase - but this is uneducated guess.

Offline BRSM

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Re: hydrogen sulfide as a reactant
« Reply #18 on: February 07, 2016, 02:33:34 PM »
The trichloromethyl anion is kind of a tricky beast as it readily collapses to the dichlorocarbene. Double addition is also problematic - that's one reason people don't make acid chlorides from carbanions and phosgene, for example. Also, gas phase reactions are much easier to talk about than do on scale with standard lab glassware.

If I had to make this compound I'd definitely play it safe and go for the above literature prep. You could do it in a day, and it's very likely to work. No point re-inventing the wheel.

Offline Babcock_Hall

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Re: hydrogen sulfide as a reactant
« Reply #19 on: February 09, 2016, 09:03:17 PM »
Thank you very much; this seems quite do-able.  The paper in Inorganic Chemistry uses pretty much the same chlorination procedure as the paper in Liebig's Ann. Chem.  The only difference that I can see is in the ratio of thionyl chloride to sulfinic acid.  It is 2.1 to one in the 1992 Inorganic Chem. paper and about 3 to 1 in the 1973 Liebig's paper.

Offline Babcock_Hall

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I wonder what's become of SALLE
« Reply #20 on: March 02, 2016, 09:26:56 AM »
I just redistilled some thionyl chloride, and I will attempt this synthesis shortly.  When I read through the first step, I was initially puzzled by the fact that the acetonitrile layer is separate from the aqueous layer.  he high concentration of sodium bicarbonate is probably responsible for this effect.  See for example Leggett et al., Anal. Chem. 1990, 62, 1355-1356. and http://www.chromatographyonline.com/salting-out-liquid-liquid-extraction-salle

It is odd, however, that the desired compound at this point sodium trichloromethanesulfinate preferentially goes into the acetonitrile layer.  I thought that SALLE was mainly used for polar but neutral molecules, but I accept that it must do so in this case.  The other thing I am concerned about is the drying step.  We can apply a vacuum at room temperature for a day or so, but we don't have the equipment needed for applying a vacuum at high temperature.

Offline zarhym

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Re: hydrogen sulfide as a reactant
« Reply #21 on: March 04, 2016, 04:01:35 AM »

How dangerous is 150kg of 98% H2O2 refluxed at 150°C considered to be?

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For my information: would SOCl2 be a possible reactant for the desired product?
I mean you won't repeat their procedure at that scale. But it won't be a lot of work to test the hazard of reflux H2O2 in grams.

Offline Babcock_Hall

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Re: hydrogen sulfide as a reactant
« Reply #22 on: March 04, 2016, 08:35:20 AM »
I tried to redistill thionyl chloride this week with mixed results.  I used a fractionating column, and the main cuts were less yellow than the first cuts.  However, it was not until I took a couple of small cuts near the end of the distillation that I was able to obtain very pale yellow liquid that boiled around 76.5 or so.  I may try combining the two large cuts and redistilling, after stirring with triphenylphosphite today, if I have time.

I may try the synthesis in Inorganic Chemistry as early as next week.  I will not be trying the procedure with H2O2 for a couple of reasons.  One is that I think that the synthesis in Inorganic Chemistry is easier and safer.  I am still not sure what my best option is to dry the product the first step, however.

Offline phth

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Re: hydrogen sulfide as a reactant
« Reply #23 on: March 04, 2016, 08:35:26 PM »
Some of the impurities are SCl2 S2Cl2, etc.  Distill with PPh3

Offline Babcock_Hall

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Re: hydrogen sulfide as a reactant
« Reply #24 on: March 05, 2016, 09:35:53 AM »
Do you think that P(Ph)3 is better than P(OP)3?  I was planning to use the latter, based on a reference from Friedman and Wetter, Journal of the Chemical Society A 1967, 36-37.
« Last Edit: March 05, 2016, 09:48:25 AM by Babcock_Hall »

Offline phth

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Re: hydrogen sulfide as a reactant
« Reply #25 on: March 05, 2016, 04:28:23 PM »
PPh3 is a solid.  Phosphites are better pi acids, I would just look at the tolman cone angle chart.  You can treat it's ability for the νCO to go up is the ability to act as a π acid.  Be wary of working with liquid phosphites because they are very toxic, may be flammable in air, and they may form azeotropes with your liquid.  PPh3 is really easy to handle, it doesnt boil easily, and it is air stable.

Offline phth

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Re: hydrogen sulfide as a reactant
« Reply #26 on: March 05, 2016, 04:46:58 PM »
by air stable I mean it won't somtaneously combust.

Offline BRSM

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Re: hydrogen sulfide as a reactant
« Reply #27 on: March 06, 2016, 03:09:27 PM »
I don't see any reason not to use P(OPh)3 if that's what the literature prescribes; I used several hundred mL without trouble this week (to make triallyl phosphite; way cleaner than going from PCl3, FYI).

I would just correct a couple of points from above:

1. It's not pyrophoric. At all. You can measure it in a measuring cylinder. Incidentally, I am not aware of any phosphites that are, although I'm not an expert.

2. It's not going to boil---boiling point is 360 °C which is almost the same as that of PPh3, even though the latter is a solid.

3. I doubt that it would azeotrope with thionyl chloride, which has a boing point almost 300 °C lower. Can't think of any azeotropes that span that difference in boiling points. I didn't struggle to distill triallyl phosphite and allyl alcohol away from it this week.

4. Can't think what relevance cone angle and pi acidity have here. I mean it's not like the thionyl chloride is back-bonding the phosphite; we just care about pure nucleophilicity to sop up "Cl+"-type impurities.

If you've got a literature reference for doing this, I would trust that more than us randoms on the internet. No point taking a risk if you have the reagent and the procedure. I'll also note that this is a method described in The Purification of Laboratory Chemicals, which is normally pretty solid.

Offline BRSM

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Re: I wonder what's become of SALLE
« Reply #28 on: March 06, 2016, 03:21:57 PM »
I just redistilled some thionyl chloride, and I will attempt this synthesis shortly.  When I read through the first step, I was initially puzzled by the fact that the acetonitrile layer is separate from the aqueous layer.  The high concentration of sodium bicarbonate is probably responsible for this effect.  See for example Leggett et al., Anal. Chem. 1990, 62, 1355-1356. and http://www.chromatographyonline.com/salting-out-liquid-liquid-extraction-salle

It is odd, however, that the desired compound at this point sodium trichloromethanesulfinate preferentially goes into the acetonitrile layer.  I thought that SALLE was mainly used for polar but neutral molecules, but I accept that it must do so in this case.  The other thing I am concerned about is the drying step.  We can apply a vacuum at room temperature for a day or so, but we don't have the equipment needed for applying a vacuum at high temperature.

This seems odd to me too. I have had acetonitrile be biphasic with an aqueous phase (sat. aq. Na2S2O3), but I have never heard of people salting salts out of an aqueous phase. I have limited experience of sulfinates, but I remember phenyl sulfinate being reasonably organic soluble, so you might be okay. Suck it and see, I guess.

Regarding drying, you don't really need anything special to dry at 80 °C---a pump, oil bath, round bottom flask and hot plate are all you require.  However, if you really can't manage that, I would think that 24 hours under a good vacuum at room temperature would suffice. I guess the most important think is to get rid of the organic solvents as I would not expect trace water to affect the protonation. There's 2% water in the acid anyway.

Offline phth

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Re: hydrogen sulfide as a reactant
« Reply #29 on: March 06, 2016, 05:24:25 PM »
Yes, phosphites probably are not that flammable, but still strong neurotoxins.  It could form a positive azeotrope, which can change the distillation, so I would agree with BRSM going with the literature examples.
3. Pi acidity changes how it interacts with the impurities.  It's the ability to interact with electron density of the impurities which has everything to do with the distillation.  I was not referring to cone angle, just that it is a common reference comparing phosphorus L type ligands electronic properties.  It has everything to do with phosphorus' ability to be effective.

I don't see any reason not to use P(OPh)3 if that's what the literature prescribes; I used several hundred mL without trouble this week (to make triallyl phosphite; way cleaner than going from PCl3, FYI).

I would just correct a couple of points from above:

1. It's not pyrophoric. At all. You can measure it in a measuring cylinder. Incidentally, I am not aware of any phosphites that are, although I'm not an expert.

2. It's not going to boil---boiling point is 360 °C which is almost the same as that of PPh3, even though the latter is a solid.

3. I doubt that it would azeotrope with thionyl chloride, which has a boing point almost 300 °C lower. Can't think of any azeotropes that span that difference in boiling points. I didn't struggle to distill triallyl phosphite and allyl alcohol away from it this week.

4. Can't think what relevance cone angle and pi acidity have here. I mean it's not like the thionyl chloride is back-bonding the phosphite; we just care about pure nucleophilicity to sop up "Cl+"-type impurities.

If you've got a literature reference for doing this, I would trust that more than us randoms on the internet. No point taking a risk if you have the reagent and the procedure. I'll also note that this is a method described in The Purification of Laboratory Chemicals, which is normally pretty solid.

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