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Author: Subject: HELP! need thionyl chloride(SOCl2) synthesis
S.C. Wack
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[*] posted on 10-1-2005 at 21:07


I don't see how that JCS article got in a footnote somewhere. Looking at it, it's very hard to believe. It was only a single sentence at the end of a long article, they speculated that a small amt was formed to explain their FeCl2 -> FeCl3 via HCl and SO2 results. Not that they found any, much less isolated some.
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[*] posted on 12-1-2005 at 13:06


Where does the Oxygen go.
mick


Perhaps if dry HCl is led into dichloromethane or chloroform then excess dry sulfur dioxide led into this mixture , the DCM distilled off, leaving the thionyl chloride

edit mick



[Edited on 12-1-2005 by mick]

[Edited on 12-1-2005 by mick]

[Edited on 12-1-2005 by mick]
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[*] posted on 12-1-2005 at 14:06


with everyone's indulgence may i rephrase the question---------------

how does one synthesis thionyl chloride without the use of any phosphorus compounds?
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[*] posted on 21-2-2005 at 16:58


S.C. - Heh, that figures. Its amazing how often things turn out that way, that one little foot note with a low probablility of being correct is blown out of porportion by the hopeful wishing of us amateur (well, me at least ;) ) chemists. I just hope that my idea for a methodology isn't copy pasted by someone as a factually correct way to produce useable quantities of thionyl chloride.

Anyways, I ran into a rather odd problem (this was some weeks ago, I've been quite preoccupied with other matters) when I attempted to synth SCl2. I'm fairly sure my setup was satisfactory: Cl2 was made by the action of muriatic acid on bleaching powder, and was then led into the bottom of a large testube filled with CaCl2 via glass tube, and from there through a rubber hose into a second glass tube which reached into the bottom of a 300ml RBF (it entered through the thermometer adapter of a distilling head) that had 64g of sulfur in a molten state contained therein. The apparatus was set up for distillation through a 300mm glass west condenser with ice water running through it.

Now, I don't know what the hell I was thinking (or not thinking), but I didn't put a gas trap on the other end, off the vacuum adapter. Needless to say, it didn't take long to fill the lab with a haze of lung and eye burning diatomic chlorine. (As a side note: even chromed steel rusts in the presence of chlorine and moisture. :o )

After it was all said and done, I had <1ml of bright orange S2Cl2 in my distillation flask, one hell of a mess, and some mild pain on deep inhalation. The oddest thing though, and the real point of my tale, is that in the flask which contained my molten sulfur, there was left this dark gray-brown material, which frothed rather than melted when heated. This material proved to be inert to most acids, and the only method of cleaning whcih worked was to reflux the mess with 50% NaOH solution for an hour and then scrape the insides with a glass rod while still hot. Like I said, one hell of a mess. Got my RBF back though.

Anybody have any idea what this bizarre polymer-like stuff is? (I still have it in a beaker.) Oh yea, the Sulfur I am using is Fisher reagent grade precipitated flour. Surely this is not causing the problem?

If someone with success at this synthesis could step foward and give an account, I would be ever so grateful.

Sincerely,
Nature's Natrium




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[*] posted on 21-2-2005 at 18:45


Quote:
Originally posted by Natures Natrium
The carbon I have on my hands is preactivated (aquarium carbon) and claims to contain less than 2% water. Thus, is the soak in CaCl2 solution really necessary?

Yes, if you look at the article, yields are greatly improved by very long soaking in CaCl2 followed by through drying and quick usage.

One thing I've been wondering though, most of my references state that SCl2 decomposes right at or around it's boilng point of somewhere around 70C, when it's used in the process described in the article Solo posted though the temperatures employed in the catalyst tube are somewhere around 150C I recall. Possibly the decomposition of SCl2 in the gas phase is what allows for this conversion?

As for your S2Cl2 synth Natrium, <1 ml, how on earth did you even manage the mililiter? I would have expected better results, your gas bubbling tube didn't dip below the sulfur did it? In case you haven't already checked you should look at Garage Chemist's post on disulfur dichloride. Just because it's neat. Also the 'Inorganic Synthesis' series recomends taking a round bottom flask and melting sulfur in it, then moving the flask around to coat the whole inside in a layre of sulfur, then apply gentle external heating while running a stream of chlorine into the flask or something like that.

Anyway, back to the method Solo posted, it recomends running the SCl2 / O2 mixture over the activated carbon in contact with it for a somewhat long period of time, but the problem would be controling the SCl2 to O2 ratio. Nitrognen would not be a hazardous contaminant so it would be great to just suck air through a dessicant (ala axehandle design) and bubble it into the SCl2 relying on the vapor pressure to make a good ratio of SCl2 to air and just passing this saturated mixture into the catalyst tube. Anyone know a good way to figure out what temperature the liquid SCl2 should be at to have the proper vapor pressure to combine with a slow bubbling of oxygen to get this to work to give around 70% yield like the article states?




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[*] posted on 21-2-2005 at 19:46


Quote:
I would have expected better results, your gas bubbling tube didn't dip below the sulfur did it?


Funny you should mention that, as it did indeed dip below the surface of the molten sulfur. I will take a look at those posts you mentioned and refine my methodology (to something that actually works). Still doesn't explain the wierd outcome regarding the molten sulfur turning into some bizarre suprisingly chemically inert material though.

As far as your question goes, I have little to no data on SCl2 other than bp. There is a vast amount of data on the net regarding S2Cl2, but so far my searches have been fruitless. If I manage to find something, naturally I will make a point of sharing it.

Sincerely,
Nature's Natrium

EDIT: After reading Garage Chemists post, it looks like my main mistake was simply not letting the reaction run long enough (since I had to stop because of chlorine faster than he did). I will however take your suggestions into consideration, and attempt to blow the chlorine right down onto the surface of the sulfur rather than bubble it through it. Actually, maybe I will ran this bad boy twice in 32g increments, just to see what qualitative and quantitative differences can be observed in the differing methods.

[Edited on 22-2-2005 by Natures Natrium]




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[*] posted on 21-2-2005 at 20:02
SCl2


According to Brauer, it is a dark red fluid with a pungent chlorine-like odour, which easily decomposes according to S2Cl2 + Cl2 <--> 2SCl2 (note, it's an equilibrium).
It is sensitive to humidity (water) and reacts with it to H2S2O3, H2S(n)O6 und H2SO4.
It is soluble in hexane without decomposition, hence I should think that any liquid aliphatic hydrocarbons should do (such as lamp oil :) ).
The boiling point is 59.6 deg C, and its density at 20 deg C is 1.62 g/cm^3




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[*] posted on 24-2-2005 at 21:24


The method Solo posted gives thionyl chloride in roughly 70% yield, but the bulk of the remainder is sulfuryl chloride. Looking at the patents KABOOOM posted for the purification of thionyl chloride they use aluminum chloride mixed with sulfur as a catallyst for the conversioin of the sulfuryl chloride impurity to SO2 and Cl2. Problem is, in the patent sulfuryl chloride is present in less then 1%, whereas the procedure that Solo brought up can produce it in >30% depending on the conditions. At least any sulfur chlorides could be somewhat easily sparated.

Anyone know if the aluminum chloride trick would work for such large concentrations of SO2Cl2?




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[*] posted on 25-2-2005 at 04:30


@Natures Natrium:
Don't blow the chlorine on the surface of the sulfur! Insert the bubbling tube as deep as possible into the molten sulfur.
Also, heat the sulfur really strong, near its boiling point.
S2Cl2 (the main product of this synthesis) begins to distill after some minutes, not immediately. It stays in the molten sulfur until the S2Cl2 concentration is too high and the sulfur can't hold it anymore. Molten sulfur seems to attract S2Cl2, just like H2SO4 attracts water.
The S2Cl2 can be converted into SCl2 by bubbling chlorine into it until it is dark red.

I also had this grey mass in the boiling flask after cooling. I think it is a mixture of mostly sulfur, S2Cl2 and higher chlorosulfanes (SxCl2), which may be responsible for the color.
I think it is best to use up all the sulfur in the synthesis to avoid the cleaning problems. Just make sure you have enough chlorine- producing compounds in your gas generator.
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[*] posted on 8-3-2005 at 11:16


Can you say SO2? (from heating the sulfur)

thionyl chloride is a nasty reagent. I'd rather make dimethyl sulfate than this.

but good luck hehe.
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[*] posted on 8-3-2005 at 11:47


thionyl chloride is a nasty reagent. I'd rather make dimethyl sulfate than this.

Thionyl chloride is nasty stuff, it will make cough, it will burn you, it will react violently with water but you can smell it and you should know where you stand. Can be neutralised by adding to cold water.

Dimethyl sulphate has no smell but a bit of a vapour pressure. To paraphase a film, it will methylate your bones, it will methylate you blood, it will methylate your DNA.

mick

Thionyl chloride is hydrolysed fast, dimethyl sulphate is hydrolysed slowly and can last in the body for a while.

mick



[Edited on 8-3-2005 by mick]
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[*] posted on 16-3-2005 at 09:54
so far, so good


Well, I has some success in getting this reaction to work. I took your advice garage chemist, and although all of it was helpful I think the most pertinant piece is heating the sulfur up as hot as possible. It also helps immensely to wrap the reaction flask and distillation head in a couple layers of aluminum foil. (I set my hotplate on 4, and it has a temp range of 150-700C, so that should be about 330C plate temp.) The reaction does have a bit of delay wherein the flask contents turn into a viscous black goo, then the orangish S2Cl2/SCl2 begins to distill over. It is very important to keep the chlorine flow continuous, as the suck-back will draw sulfur into the gas inlet tube where it will solidify and clog the flow. This time I also remembered to put a gas scrubber on the end of the apparatus, and it turned out that 1.5L of H2O with half a pound of NaOH was suffecient for the entire reaction (which took about 4-5h from beginning heating and chlorine flow). Also an absolute must is a suck back bottle between the recieving flask and the 2L beaker with the scrubbing wash in it. I had to empty the suckback bottle twice (poured back into beaker) when I failed to keep the chlorine flow consistant. Also, I was intitially worried that the CaCl2 would not be effecient enough a dessicant, but my worries proved to be unfounded.

EDIT: Screwed up the math, lets try again:
My total yield is 98g (60mL) of orangish red liquid from 64g of sulfur. (d1.63, yield: <73% S2Cl2 OR <46% SCl2 ) Not great, but a lot better than last time. :D

Right now I have my RBF and distillation head on a sonic bath, trying to free up the ground glass joint. It turned out the silicon grease I was using reacted, and formed a black crust inbewteen the joints. God please don't let me break these things trying to get them apart...my glass blowing skills suck and I haven't built an annealing oven yet. :(

Also, does anyone know of a good solvent for sulfur besides CS2? I am having a hell of a time getting the sulfur off all my glassware. I had intially figured that sulfuric acid would work well, but such is not the case. Although the merck states that sulfur has limited solubility in acetone, it doesn't appear to be working either. I would appreciate any advice on this matter.

[Edited on 16-3-2005 by Natures Natrium]




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[*] posted on 16-3-2005 at 12:46


There are limited good solvents for sulfur. Carbon tet and carbon disulfide are the only good ones I know of. Boiling with HNO3 will destroy it as will some other oxidizing combinations. Try heating with a blow torch, push come to shove. I had the same problem when I went all out distilling CS2 from charcoal and sulfur, the sulfur carried over onto everything and I found that a bottle brush and water was a good way to loosen it up and wash it away.
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[*] posted on 16-3-2005 at 13:51


People have said that S<sub>2</sub>Cl<sub>2</sub> was a good solvent for sulfur. Look in the appropriate thread.
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[*] posted on 16-3-2005 at 14:03


Quote:
Originally posted by neutrino
People have said that S<sub>2</sub>Cl<sub>2</sub> was a good solvent for sulfur. Look in the appropriate thread.


Uhm, that's supposed to be a joke, yes? ;)

Through a combination of things, mostly elbow grease and potassium permanganate I am making progress. I also got the glass joint unstuck by rapidly heating the outside jacket in a blue bunsen flame while pulling the two pieces apart.:)

Then I ended up breaking the RBF when I was trying to dig the goo out with a glass rod. :( :mad: Crap, I only have 2 RBFs left out of my set, a 100 and a 500. Although my own foolish mistakes take the bulk of the blame, I don't intend to ever buy kimax again. Both the pyrex and bomex glassware I have has lasted while all my kimax shit shattered at the wierdest opportunities. (ie a drop to the floor wasnt a problem, but a slight tap and they break into a million pieces.)




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[*] posted on 16-3-2005 at 14:32


Heating with NaOH, etc. will remove the S.

Was the Kimax used? You never know what happened to it before. The glass remembers previous abuses. Because I have a lot of Kimax and never had problems with it.
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[*] posted on 16-3-2005 at 16:26


Quote:
Originally posted by Garage Chemist
S2Cl2 is interesting because it can dissolve up to 67% sulfur. It also dissolves white phosphorus and is therefore an ideal substitute for CS2. It's also non- flammable.


I was referring to that comment made in the disulfur dichloride thread.

Your bomex outlasted your kimax? That is very odd, as in my experience kimax is good quality glass while bomex is worthless garbage.
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[*] posted on 16-3-2005 at 16:52


Quote:
Originally posted by S.C. Wack
Heating with NaOH, etc. will remove the S.

Was the Kimax used? You never know what happened to it before. The glass remembers previous abuses. Because I have a lot of Kimax and never had problems with it.


Well, actually no, I bought it brand new. I suppose I have put it through its paces, but it seems like everyone else talks about how thier glassware has lasted forever, and I am just fed up with the stuff breaking. I have broken 2 1L erlyenmeyers, 8 thermometers (I have no shortage of elemental mercury :o), a 500mL filtering flask; 50, 250, and 300mL ground glass RBFs; a 300mm west condenser, more glass rods and tubes that I care to count, and a couple of glass stem funnels. That being said, I cannot lie to myself with such hard empirical data presented so blatantly before me. I am hard on glassware, although god knows I try not to be. I dont know if it is clumsiness or bad luck, but it is heart rending to see this much good labware go to waste.

At this point I am going to have to save up and just buy a whole new 24/40 kit, although I have no intention of spending $550 like I did on the last one. Now that is truly painful. At that price I expected these pieces to last forever. :(

EDIT: neutrino, I see what you are getting at, I just cant see using the S2Cl2 I made to clean up the mess from the reaction which made it! It just seems redundant somehow. As for the bomex, well, all the bomex pieces I have are nice and thick. All the kimax and pyrex tends to be thinner glass. Although I have seen others professing such beliefs, I cannot imagine why someone would want thin glass instead of thick.

[Edited on 17-3-2005 by Natures Natrium]




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[*] posted on 17-3-2005 at 14:32


I believe that there is a good reason why crucibles and flasks are thin, they withstand temperature swings better. I attribute my good luck with glassware to cautious heating and cooling. If distilling, I only turn up the heat on the bath low until the contents of the flask have warmed up. When heating with a flame, the flame is increased and decreased in steps. This is what I was taught, don't know how important it really is, its just habit.

There is another route to SOCl2 that doesn't get mentioned much. S and Cl2O @ -12C. I think that would have got a (rated as: dangerous) at the Hive. Wouldn't be my first there, but here I just wanted something on-topic to say. You can see what Wurtz had to say back in the day at:
http://gallica.bnf.fr/document?O=N003019
after entering 460 (the first page of the article) in the Aller Page box.
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[*] posted on 17-3-2005 at 16:08


I believe that I was successful in converting the S2Cl2/SCl2 solution into mostly SCl2. It took a shitload of chlorine to get the job done, but here are the results. Uhm, crap, I thought these forums hosted images? I guess will I will just add as an attatchment. I also got a shot of my improvised catalyst tube which I expect I will be using within a couple of days. Too bad I didnt think to get a pic of the orange-red liquid that was the result of the chlorination of molten sulfur, it was purty. :D

SC Wack, I can see your point about thermal expansion and glass thickness. However, like you, I try to heat and cool all my glassware as gently as possible. The only thing I ever directly broke from thermal shock was a thermometer. Considering how much stuff I have broken from physical impact, I am thinking that thicker glass would be a better choice for me.

Anyways, pics:

Attachment: 2005_03_17.rar (64kB)
This file has been downloaded 534 times





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[*] posted on 22-3-2005 at 08:44


I'm sorry that this comes so late, but I found a SCl2 preparation in Brauer that doesn't require a distillation setup for the S + Cl2 reaction.

The important thing is that ground sulfur rods (i.e. sulfur which has been molten and solidified) react with chlorine at room temperature.
Note that flowers of sulfur and precipitated sulfur cannot be used directly. You have to melt them, let the melt solidify and break it up into small pieces.

The procedure goes like this:
A 1 Liter two-neck round-bottom flask with a reflux condenser on the middle neck and a chlorine inlet tube on the other neck is charged with 200g coarsely ground sulfur rods. A rapid stream of dry chlorine is introduced until the contents of the flask have liquefied completely by forming S2Cl2.
Then 0,1g iron powder is added as a catalyst and the chlorine addition is continued for half an hour while cooling the flask to 20°C with a water bath.
The formed dark red solution (contains mostly SCl2, along with S2Cl2 and dissolved chlorine) is left to stand for an hour.
Then 2ml of PCl3 are added as a stabiliser (prevents the decomposition of SCl2 during distillation) and the liquid is distilled over a short fractionating column, collecting the fraction that comes over at 55- 62°C. It is fractionated again and the fraction boiling at 60°C is collected in a flask which contains a few drops of PCl3. The resulting SCl2 is very pure. Yield about 70%.
S2Cl2 can also be used as a starting material, of course.

SCl2 can be stored a few days in presence of PCl3, it decomposes slowly to S2Cl2 with dissolved Cl2. It can be re-purified by distilling with a few drops of PCl3.

I think that the PCl3 can be substituted by adding a gram of red phosphorus (from matchbox strikers) while still bubbling chlorine into the mix.

I'm going to make some S2Cl2 this way, by using 50g sulfur and stopping the chlorine addition when the liquid starts to become red (I don't want SCl2, I want S2Cl2 for making acetic anhydride).
I won't dry my chlorine because I don't have a suitable apparatus, but this shouldn't cause problems because I'll distill the S2Cl2 afterwards.

[Edited on 22-3-2005 by garage chemist]

[Edited on 22-3-2005 by garage chemist]

[Edited on 22-3-2005 by garage chemist]
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[*] posted on 23-3-2005 at 05:58


Quote:
Originally posted by garage chemist
I'm sorry that this comes so late, but I found a SCl2 preparation in Brauer that doesn't require a distillation setup for the S + Cl2 reaction.

The important thing is that ground sulfur rods (i.e. sulfur which has been molten and solidified) react with chlorine at room temperature.
Note that flowers of sulfur and precipitated sulfur cannot be used directly. You have to melt them, let the melt solidify and break it up into small pieces.


Ah, I wondered why my precipitated flour wouldn't react with chlorine unlil it melted, even thought the Merck states that sulfur combines with haolgens in the cold. I am eagerly awaiting to hear how this goes for you.

As for me, I came up with a little brain storm this morning for increasing the amount of sulfur chlorides I have. Since bubbling in chlorine into S2Cl2 produces SCl2 in a 1:2 molar ratio, and refluxing SCl2 with sulfur reprouduces S2Cl2 in a 1:1 ration, every repitition of this cycle would double the molar amount of sulfur chlorides.

S2Cl2 + Cl2 -> 2SCl2
2SCl2 + 2S -> 2S2CL2

This would be a good way of building up a stock once S2Cl2 has already been procured.

The catalysis is currently on hold, as I foolishly broke the tip off of my bootleg catalysis tube when trying to modify it, and now I have to figure out a way to rebootleg it.

I will say that I did prepare the catalyst as I stated earlier, by leaving in CaCl2 solution for three days, refluxing for an hour, filtering, and washing with methanol. The contents were then heated in a flask until no more vapor evolved, and then (while still being heated quite hot) was put under hard vacuum for an hour.

Still, I found there to be significant quantities of water left in the carbon. I don't think the importance of thoroughly drying the catalyst can be underscored enough.

I wasn't able to produce in thionyl chloride in my first run, but really this isn't suprising. I don't have real precise control over the temp of my ni-cr wire, I changed the catalyst preperation, and there was some residual water left in the catalyst. I used up about 20mL of my SCl2, which is why I have been thinking of easier ways to replinish my supply.

Lastly, it didn't look to me like hardly any SCl2 came over with the dried air until the waterbath that the SCl2 containing flask was in hit about 50C. At this point I finally started to see small droplets condensing in the distillation head. Naturally, I still have no idea what the optimization point is.

Once I make more chlorides of sulfur, and rig my system up to work again, I intend to give it another shot.




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[*] posted on 23-3-2005 at 07:47


The last three hours, I've been trying to make S2Cl2 without a distillation setup.

I've put my sulfur chunks in the described apparatus and began to add the chlorine.
Nothing happened. No chlorine was absorbed and the chunks looked the same like before, just the green gas was over them.

I heated the sulfur and this started the reaction. After some chlorine was absorbed, the liquid became less viscous. I heated it until the produced S2Cl2 refluxed and continued to add chlorine. The liquid had a very dark red color, like hot molten sulfur and stayed like this throughout the entire procedure. I always kept the mixture at its boiling point with a bunsen burner.
I added the calculated amount of chlorine over the course of 2h, but near the end, the stupid chlorine generator frothed over and some of the TCCA/HCl sludge was pushed through the tubing. I had no time to react, and in the moment the sludge touched the red boiling liquid, there was a violent reaction, the liquid frothed a lot and it sounded like water being poured into hot oil (the liquid was at ~200°C AND reacted with the water). A white cloud of SO2 erupted from the condenser and filled my entire fume hood for a few seconds.
I disconnected the chlorine apparatus and turned off the burner.

I weighed my liquid, and this is the strange thing: from 50g sulfur, I only got 64g of red oily product. It remained liquid after cooling, so it must be S2Cl2. I'm going to distill it tomorrow, no more experiments with nasty sulfur chlorides for today.
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[*] posted on 24-3-2005 at 10:02


Some sulfur has crystallized out of the liquid in nice rhombic crystals. Therefore I now have a saturated solution of sulfur in S2Cl2.

The method from Brauer sucks. Better bubble chlorine through molten sulfur in a distillation apparatus.

BTW, hot NaOH solution is perfect for cleaning glassware from sulfur. I put some boiling 20% NaOH in my round- bottom flask, stoppered its two necks and shook it. After a few minutes, it was perfectly clean.

KClO3 can oxidise PCl3 to POCl3 (Does anyone have a procedure for this? I'd greatly aprecciate some info on reaction conditions!).
Maybe this also works with SCl2?
It is also possible that it would be oxidised to SO2Cl2, but SO2Cl2 reacts with SCl2 to SOCl2, so using stochiometric amounts of SCl2 and KClO3 could work to produce SOCl2.


[Edited on 24-3-2005 by garage chemist]
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[*] posted on 11-4-2005 at 01:58


Why couldn't you react DMSO with 2HCl (both are cheap and easy to get), and collect the two methanes given with each thionyl chloride to run your car seeing as how gas is getting too expensive lately.
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