Chemical Forums
Chemistry Forums for Students => Organic Chemistry Forum => Organic Chemistry Forum for Graduate Students and Professionals => Topic started by: Babcock_Hall on January 23, 2016, 10:36:10 AM
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There is a starting material that I would like to have that is no longer commercially available (trichloromethanesulfinyl chloride). The only published route I have found involves the use of hydrogen sulfide gas, which is quite toxic. I have standard glassware available to me. How difficult would it be to use hydrogen sulfide in a reaction? The earliest I would perform this reaction is a couple of months from now, and I have a little NaSH in the lab at the moment, which I could potentially use to generate H2S if that would be safer.
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I think generating it in situ would be the safest method because you could control the amount of H2S generated. Could you generate the H2S in situ from the NaSH or Na2S and acid?
Or you could generate in buchner flask by having the sulfide placed in and slowly adding from addition funnel and lead the gas through the tubing into the reaction. Im not sure how to get rid of the excess gas after the reaction is over, maybe add something like NaOH to convert it back to NaSH?
Not sure if I answered your question
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You can probably buy lecture bottles of H2S. If not use a Kipps generator. https://en.wikipedia.org/wiki/Kipp%27s_apparatus (https://en.wikipedia.org/wiki/Kipp%27s_apparatus)
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What're you trying to do with it?
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I would like to reduce trichloromethanesulfonyl chloride to trichoromethanesufinic acid with H2S then chlorinate it to the sulfinyl chloride. Of course if there were another route to the sulfinyl chloride, that would be great.
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maybe oxidation of sulfenyl chloride? or rather thiol to sulfenic acid?
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Trichloromethanesulfenyl chloride is commercially available, but I would need a method of oxidizing it but not over oxidizing it.
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I looked into a clever book and found something you might want to look at:
http://www.orgsyn.org/Content/pdfs/procedures/CV5P0709.pdf
http://www.sciencedirect.com/science/article/pii/S0040403900842938
It could work on the substrate you want to make what you desire.
Im not sure if the needed disulfide is available or can be easy made (maybe from trifluorobromomethane and Na2S2?)
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http://www.3bsc.com/index/pro_info.php?id=98035
http://www.chemos.de/productinfo.php?ART=140755
http://www.hanhonggroup.com/en/?sType=4&keyword=25004-95-9&button2=&mod=products&act=list
Just buy it.
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I have contacted a few of the firms that come up in searches. One company wanted over $10,000. Some companies did not even respond.
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There is an easy way to synthesis this chemical. According to the literature I found (attached below), trichloromethanesulfinyl sulfenyl chloride(95.4kg), H2O2 (150kg, 98%) and KF(89kg) are refluxed at 150 degree C to synthesis trichloromethanesulfinyl chloride and used in situ. The sulfinyl chloride is pumped into another container as a gas. The container you need should tolerant HF (No glass).
Here is the reference. It's written in Chinese.
CN102633722 page 10&11
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Funny, hydrogen sulfide is not toxic in small quantities is actually beneficial to cells. The experiemtn was to take a mouse/rat and remove its kidney. Next they sewed it back in. The animal exposed to the H2S livedwith 20% death versus control 70% death. H2S is apart of the natural digestion, garlic, horseradish (isothiocanates). It's prolly worse because its stinky than toxic at small levels. Cysteine, homocysteine are natural releasing agents of H2S.
This guy does some good work on detection and ways to measure it: http://chemistry.uoregon.edu/profile/pluth/
In addition, photo labile dithianes are possible sources from within my department:
http://pubs.acs.org/doi/abs/10.1021/ol401118k
I can ask questions. All they study is sulfur chemistry. Us too..
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Funny, hydrogen sulfide is not toxic in small quantities is actually beneficial to cells
It is a dose that makes a poison, nothing new here. Still, it is quite dangerous.
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Funny, hydrogen sulfide is not toxic in small quantities is actually beneficial to cells. The experiemtn was to take a mouse/rat and remove its kidney. Next they sewed it back in. The animal exposed to the H2S livedwith 20% death versus control 70% death. H2S is apart of the natural digestion, garlic, horseradish (isothiocanates). It's prolly worse because its stinky than toxic at small levels. Cysteine, homocysteine are natural releasing agents of H2S.
This guy does some good work on detection and ways to measure it: http://chemistry.uoregon.edu/profile/pluth/
In addition, photo labile dithianes are possible sources from within my department:
http://pubs.acs.org/doi/abs/10.1021/ol401118k
I can ask questions. All they study is sulfur chemistry. Us too..
It's beneficial in small quantities. However, in real synthesis it's always VERY LARGE quantities. We are not talking about ppm level in biology. In chemistry, most likely the chemical are used in mmol to kmol level.
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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?
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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.
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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!
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With thionyl chloride, I had imagined to react chloroform, both in gas phase - but this is uneducated guess.
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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.
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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.
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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.
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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.
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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.
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Some of the impurities are SCl2 S2Cl2, etc. Distill with PPh3
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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.
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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.
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by air stable I mean it won't somtaneously combust.
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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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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.
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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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Thanks for all of the replies. Sulfuryl chloride boils at 68-70 °C; therefore, reducing it to thionyl chloride is an attractive approach. In reading through the paper in JCS Chem. Soc., it sounds as if triphenyl phosphite reacts with thionyl chloride and other sulfur chlorides, as well as sulfuryl chloride, just not as quickly. The authors used 0.5 hours of reaction time. It might be that long reaction times would create more problems.
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Thanks for all of the replies. Sulfuryl chloride boils at 68-70 °C; therefore, reducing it to thionyl chloride is an attractive approach. In reading through the paper in JCS Chem. Soc., it sounds as if triphenyl phosphite reacts with thionyl chloride and other sulfur chlorides, as well as sulfuryl chloride, just not as quickly. The authors used 0.5 hours of reaction time. It might be that long reaction times would create more problems.
It is easier for phosphorus to take an oxygen from sufuryl chloride than thionyl chloride because the average bond order of the oxygens is different. It's a better bet not to go over. Thionyl chloride is the thermodynamic product on the way to phosphonate.
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I redistilled some of the moderately yellow material after stirring with P(OPh)3. The second fraction was pale yellow, and the third fraction was almost clear.
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In the past couple of days I have tried to reproduce the first step of the synthesis from Inorganic Chemistry, the synthesis of sodium trichloromethylsulfinate. After the synthesis, I removed solvent, added methanol and filtered away some precipitate, and removed the solvent with rotary evaporation. I dried for about 5.5 hours at 80 °C under vacuum and for a couple of hours at room temperature under vacuum. The mass was roughly 33% higher than theoretical. The authors of the paper claimed a 60% yield at this point. Therefore, I put the material back on high vacuum overnight at room temperature, and the mass is virtually unchanged.
The next step of the synthesis is to acidify with sulfuric acid and then to distill the product, which is the free sulfinic acid. Should I proceed to the next step, should I make an additional attempt to dry the material, or should I try the methanol step over again?
ETA
If the impurity is a salt, it might be NaBr, NaHCO3, or the sulfur-containing byproduct created by the attack of dithionite on bromotrichloromethane.
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Here is the reaction scheme from Zhang et al., Inorganic Chemistry 1991, 32, 492-494. I am not sure how to proceed, given my greater than theoretical yield.
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Yuan Fa Zhang, Robert L. Kirchmeier, Jean'ne M. Shreeve
Inorg. Chem., 1992, 31 (3), pp 492–494
DOI: 10.1021/ic00029a028
are you talking about this paper?
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The procedure says 'the residue was treated with 50 mL of absolute methanol'.
I would guess removing water from the system is the key in this synthesis, otherwise it's very likely you end up have some water in the crystal structure.
If this is the reason, you can try drying your crude solid from acetonitrile/water layer and also dry your methanol before you mix them together.
In organic synthesis, organic layers are often dried with MgSO4 or Na2SO4. Since you are synthesis a salt, you may have to find another drying agent.
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Yuan Fa Zhang, Robert L. Kirchmeier, Jean'ne M. Shreeve
Inorg. Chem., 1992, 31 (3), pp 492–494
DOI: 10.1021/ic00029a028
are you talking about this paper?
Yes. With respect to the methanol, I used a high grade of commercial solvent. I have decided to dry some myself if I repeat either the methanol step or the whole synthesis. However, I think that another source of water is the acetonitrile layers that I combined. There must have been at least some water in the acetonitrile. However, the protocol that I am following does not call for drying the combined acetonitrile layers.
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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.
Forgive my ignorance, but I still can't see the relevance of Pi acidity here. Pi acidity refers to the ability of a compound to accept electron density into empty Pi orbitals. This interaction is impossible with SOCl2, S2Cl2, SCl2 etc --- they simply don't have orbitals of the correct symmetry for this kind of interaction. Seriously, draw me an orbital diagram; if I'm wrong I'd like to know why. Have a look at a metal-phosphine two-way bond (https://en.wikipedia.org/wiki/Metal_phosphine_complex) and explain to me how chlorine is doing that.
All that's going on here is attack of the phosphite lone pair on an electrophile of the "Cl+" type: S2Cl2, SCl2, whatever. One way movement of electrons. Simple, first year orgo. No electron density is moving from the electrophile to phosphorus. Thus pi acidity---the ability of the phosphine to accept electron density---does not seem relevant to me. The phosphite is just sucking up "Cl+" donating impurities in SOCl2 as a sigma nucleophile. That's why the alternative of olefin-rich terpines like linseed oil is offered in The Purification of Laboratory Compounds. Really, any soft nucleophile will do. Pi acceptor ability isn't needed here.
Also, nobody's serious proposing the preparation of SOCl2 from sulfuryl chloride (SO2Cl2), right? SOCl2 is dirt cheap, and actually making it from any other substance is not going to be worth your time. Firstly, I wouldn't count on phosphites to reduce SO2Cl2 in the way you mean, as SO2Cl2 is chlorine rather than an oxygen donor. Look at its reactivity --- it's a chlorinating reagent! Can anyone find me a reaction where a nucleophile attacks SO2Cl2 on oxygen? And even if you could reduce SO2Cl2 to SOCl2, in my opinion you'd be crazy to even bother.
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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.
Forgive my ignorance, but I still can't see the relevance of Pi acidity here. Pi acidity refers to the ability of a compound to accept electron density into empty Pi orbitals. This interaction is impossible with SOCl2, S2Cl2, SCl2 etc --- they simply don't have orbitals of the correct symmetry for this kind of interaction. Seriously, draw me an orbital diagram; if I'm wrong I'd like to know why. Have a look at a metal-phosphine two-way bond (https://en.wikipedia.org/wiki/Metal_phosphine_complex) and explain to me how chlorine is doing that.
All that's going on here is attack of the phosphite lone pair on an electrophile of the "Cl+" type: S2Cl2, SCl2, whatever. One way movement of electrons. Simple, first year orgo. No electron density is moving from the electrophile to phosphorus. Thus pi acidity---the ability of the phosphine to accept electron density---does not seem relevant to me. The phosphite is just sucking up "Cl+" donating impurities in SOCl2 as a sigma nucleophile. That's why the alternative of olefin-rich terpines like linseed oil is offered in The Purification of Laboratory Compounds. Really, any soft nucleophile will do. Pi acceptor ability isn't needed here.
Also, nobody's serious proposing the preparation of SOCl2 from sulfuryl chloride (SO2Cl2), right? SOCl2 is dirt cheap, and actually making it from any other substance is not going to be worth your time. Firstly, I wouldn't count on phosphites to reduce SO2Cl2 in the way you mean, as SO2Cl2 is chlorine rather than an oxygen donor. Look at its reactivity --- it's a chlorinating reagent! Can anyone find me a reaction where a nucleophile attacks SO2Cl2 on oxygen? And even if you could reduce SO2Cl2 to SOCl2, in my opinion you'd be crazy to even bother.
Yes, you sir are right, but I'm not so good with words sometimes; no one but Chem engineers would turn SO2Cl2 to thionyl chloride. Phosphorus has empty pi orbitals like this(https://upload.wikimedia.org/wikipedia/commons/6/66/Connelly-Orpen-PR3-pi-acceptor-orbitals.png).
If chloride has pi electrons, which are not solely sp3, then it will constructively interfere with any empty orbital. Even lone pairs will constructively interfere with other lone pairs aka the alpha effect eg HOOH. Once the phosphite forms a complex with Cl+, it can form a trig bipyramid complex with a the rest of the complex. It will better accept electrons here if it is more pi acidic, P(OPh)3 also is more favorable for steric reasons. Therefore, it will drive the reaction in the forward direction if there is less electron density...phosphorus has accessible d orbitals so it acts like a metal
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Also, nobody's serious proposing the preparation of SOCl2 from sulfuryl chloride (SO2Cl2), right? SOCl2 is dirt cheap, and actually making it from any other substance is not going to be worth your time. Firstly, I wouldn't count on phosphites to reduce SO2Cl2 in the way you mean, as SO2Cl2 is chlorine rather than an oxygen donor. Look at its reactivity --- it's a chlorinating reagent! Can anyone find me a reaction where a nucleophile attacks SO2Cl2 on oxygen? And even if you could reduce SO2Cl2 to SOCl2, in my opinion you'd be crazy to even bother.
I think that sulfuryl chloride is a common impurity in thionyl chloride, and one that is difficult to remove by distillation because of their closeness in boiling points. The purpose of reducing it to thionyl chloride is to produce thionyl chloride of greater purity.
With respect to the excess mass in the synthesis of sodium trichloromethanesulfinate (the first step in the synthesis), I have heard back from one of the authors of the paper, who was friendly but who was not able to offer any advice.
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Yesterday I added sulfuric acid and attempted to bring the solid into solution, but it remained a partial suspension. I attempted a distillation at reduced pressure. An unpleasant gas was given off, and I collected very little product. All I obtained was a little bit of an orange liquid and a yellow semi-solid.
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3Read a patent a while ago but I can't find the correct method. SO2+Cl2. The way the reaction is controlled is through a cylindrical chamber of hot catalyst; the extent of the reaction varies by Δx, so it's really easy to make sufuryl chloride or thionyl chloride. Sufuryl chloride is the shorter distance of the two
I think it's work a scifinder search. Pass it through activated carbon with metals to stabilize the stationary phase in a distillation column. They tested the spiked carbon by oxidative stress of the catalyst using thermal gravimetry temperature step. After a certain point the carbon, spiked with e.g. a 1000 ppm Fe..., would suppress the formation of the byproducts such as CCl24, CHCl3, DCM.
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I think my first attempt to make sodium trichloromethanesulfinate can be declared to be a bust, but I won't have time to try again for a little while. That may be a good thing; perhaps like Sherlock Holmes, I need to sit on a pillow and smoke a bunch of bowls of tobacco until it comes to me. I don't remember which story that was...
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Hey Babcock_Hall, are you still interested in this ClSOCCl3 intermediate? It may be a candidate for the intermolecular photocoupling
http://www.chemicalforums.com/index.php?topic=81721.0
if heat doesn't suffice. Sketch appended.
Starting with SOCl2 and CHCl3:
- Both have a usable vapour pressure, 16kPa and 26kPa at +25°C, still 5kPa and 8kPa at 0°C
https://en.wikipedia.org/wiki/Thionyl_chloride
https://en.wikipedia.org/wiki/Chloroform
and ClSOCCl3 is expected to be less volatile hence exit the reaction zone, avoiding a further step. - At 254nm from a low-pressure mercury lamp, SOCl2 molecules have 6.3*10-18cm2 cross-section, but CHCl3 only 2.0*10-23cm2 and HCl <<3.4*10-23cm2, CCl4 (represents C2Cl6) 1*10-21cm2, N2 nothing
http://satellite.mpic.de/spectral_atlas
At +25°C and half the SOCl2 vapour pressure, the exponential attenuation distance is 0.4mm, and at 0°C 2.5mm. - Under the small assumption that an absorbed photon breaks a Cl-SOCl, the °Cl radical abstracts an °H from CHCl3 to create HCl and °CCl3. Since H-CCl3 is 8kJ weaker than at t-butyl it must proceed quickly even at 0°C.
http://staff.ustc.edu.cn/~luo971/2010-91-CRC-BDEs-Tables.pdf - Does °CCl3 break the abundent Cl-SOCl? I have no data for the bond dissociation energy. The unwanted reaction seems to be energetically neutral, the activation energy is unknown. That's a big uncertainty. Cold must improve.
- °SOCl and °CCl3 hopefully recombine as C2Cl6, SxOyClz and, with 50% yield at best, ClSO-CCl3. Fortunately, the reactants are much cheaper than the sought intermediate.
- 28W light from 87W electricity make 5mol/day photons, and at 50% best possible yield, 519g/day ClSOCCl3.
http://www.light-sources.com/germicidal-uvc/products/low-pressure-mercury-lamps/high-output-quartz
How common is such a photochemical reactor in labs? Besides toxicity, the setup looks rather simple.
Marc Schaefer, aka Enthalpy
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I don't know how common that sort of equipment is, but I will ask around. My first attempt at the synthesis did not lead to anything good, but I hope to make this a summer project.