Sciencemadness Discussion Board - HELP! need thionyl chloride(SOCl2) synthesis

HELP! need thionyl chloride(SOCl2) synthesis

navarone - 31-1-2004 at 14:16

zzup bros?
look guys....does any body of u know how to make thionyl chloride SOCl2?
i would really need it man!
and possibly with easy getting chemicals.
u know dont come out with crasy substances like cyan,palladium,arsenic or others stricted chems.
u know a recipe that even a kid would be able to do.

BromicAcid - 31-1-2004 at 14:33

If you looked hard enough online Im sure you could come up with a preparation but here is the basic one. Not OTC chems but oh well, maybe you could build off it. Next time though try and word your post so it seems like you're actually interested in it and not just like you expect immediate gratification.

Thionyl Chloride

PCl5 + SO2 -----> POCl3 + SOCl2

A well-dried, 250-ml., two-neck, round-borrom flask, equiped with a reflux condenser is connected to a CaCl2 tube, is loaded with 100 g. of PCl5. Sulfur Dioxide is intorduced though a gas inlet rube extending to the bottom of the flask. The SO2 is carefully dried by allowing it to bubble though two wash bottled containing H2SO4. The reaction which can be acceletated by shaking the flask, is completee after several hours when all the PCl5 dissolves. The products
are seperated by repeated careful fraction using an efficent collum. BP of SOCl2 77C and BP of POCl3 108C This preparation is generally not free of phosphorus compounds. The yeild of SOCl2 is 30g (50% of theroretical)

From Inorganic Preparations book I

Mumbles - 31-1-2004 at 18:36

There's another method, on this site no less posted my Solo. The title is Thionyl Chloride.

You pass SO3 into SCl2 to make SO2 and SOCl2. I've never tried it, but it looks like it would work. Perhaps you should consult the other thread. Those can be made from Sulfur, Chlorine, and some NaHSO4. All quite easy to obtain/make.

It's obvious you didn't search. This could of easily gone in that thread. Or you could of completely avoided making the thread by finding both answers provided thus far there.

guaguanco - 1-2-2004 at 23:52

Quote:

Originally posted by navarone
HELP! need thionyl chloride(SOCl2) synthesis

No you don't. Trust me, you really don't.

meselfs - 7-2-2004 at 16:15

IIRC, this is how you make it:

Pass a mix of sulpfur dioxide & chlorine through activated carbon or camphor catalyst.

Note that this stuff is slowly hydrolyzed to H2S04.

guaguanco - 7-2-2004 at 22:50

Quote:

Originally posted by meselfs
Note that this stuff is slowly hydrolyzed to H2S04.

Actually, it hydrolyses to SO2 and HCl

meselfs - 8-2-2004 at 18:29

Right.. I was thinking SO2Cl2 . Sorry.

interesting patents

KABOOOM(pyrojustforfun) - 9-2-2004 at 20:05

US 4,337,235 Purification of thionyl chloride
US 4,388,290 Purifying thionyl chloride with AlCl3 catalyst
US 4,759,826 Process for purifying thionyl chloride
US 5,498,400 Preparation of thionyl chloride and phosphorus oxyhloride from sulfur dioxide

acx01b - 24-10-2004 at 10:37

Hi mfs,

SO2 + Cl2 --> SO2Cl2 with camphor catalyst at 20°C! (this reaction looks so great)
do you have already tried it ?
i think i am gonna try it

SO2Cl2 + H2O --> H2SO4/HCl oleum

[Edited on 24-10-2004 by acx01b]

[Edited on 24-10-2004 by acx01b]

Ref: thionyl chloride synthesis

solo - 25-10-2004 at 08:52

try reading post [537708] over at the hive.....

Excerp
Summary

The oxidation of SCl2 by atmospheric oxygen has been investigated using activated charcoal as the catalyst. The reaction has been discussed in terms of two reactions. Firstly, the irreversible oxidation of SCl2 to SO2Cl2 :
SCl2 + O2 __> SO2Cl2

and secondly, the reversible reaction between the SCl2 and the SO2Cl2 to form SOCl2 :

SCl2 + SO2Cl2 <__> SOCl2

This oxidation has been investigated at the temperatures 193°C, 138°C and 320°C, and the equilibrium constant of the reversible equilibrium reaction has been determined for each temperature. From the temperature dependence of this equilibrium constant it is found that the reaction is exothermic to the extent of 18 kcal.

(Yield of SOCl2 at 193°C : ~73%)

..............

Dilaudid - 9-11-2004 at 10:51

Well I know a way... unfortunatly in Canada under the CanadaChemical Weapons Convention Act thionyl compounds are illegal.

I would assit if I could...but bound by Canadian Laws and George Bushs' reelection I wouldn't want to assist any terrorist organization

Edit by Chemoleo: That's good to know, but what's the point of your post?

[Edited on 9-11-2004 by chemoleo]

neutrino - 9-11-2004 at 14:25

It's doubtful whether anyone will come after you for telling a bunch of harmless hobbyists how to synthesize one simple chemical. You could always give us a reference to it or just tell one person who may or may not take it upon himself to break the anti-knowledge laws for the good of mankind. If they’ve made exchange of knowledge an international offence, then we all have much bigger things to worry about.

Reverend Necroticus Rex - 10-11-2004 at 10:09

I am quite interested in that synthesis using a camphor catalyst, could you perhaps post a patent number or something to that?

Chris The Great - 18-12-2004 at 18:18

Taken from Handbook of Preperative Inorganic Chemistry, Volume 1, Second Edition, pages 382 to 383.

Thionyl Chloride
SOCl2

I.
SCl2 + SO3 = SOCl2 + SO2
102.98 80.07 118.98 64.07

Flask A of the apparatus in Fig. 166 is charged with pure 65% oleum, flask B with 100g of SCl2 (see p. 370). Flask A is slowly heated in an H2SO4 bath, while B is cooled with an ice water bath. The stoichiometric quantity of SO3 is thus gradually distilled onto the SCl2. The reaction proceeds with SO2 evolution (use a hood!). Partial solidification of the flask contents feqeuntly occurs at the beginning. However, the contents should again be completely liquid when the addition is completed. If necessary, the flask is heated at the end for a short time on a water bath. Flask A is then disconnected, the ground joint at C is stoppered, and the mixture is slowly distilled in a column (use a hood!). The mixture must be protected from contact with atmospheric moisture. The middle fraction is further porified by repeated careful fractionation, with S added to the distillation charge in order to convert all sulfur chlorides present to S2Cl2. The contents are distilled through an efficient column until a completely colorless product, coming over at 76-77C, is obtained. The yeild is about 80% of theoretical.

Fig.166

II.
PCl5 + SO2 = POCl2 + SOCl2
203.37 65.07 153.35 118.98

A well-dried, 250mL, two neck, round bottom flask, equipped with a reflux condenser connected to a CaCl2 tube, is loaded with 100g of PCl5. Sulfur dioxide is introduced through a gas inlet tube extending to the bottom of the flask. The SO2 is carefully predried by allowing it to bubble through two wash bottles containing concentrated H2SO4. The reaction, which can be accelerated by shaking the flask, is complete after several hours, when all the PCL5 dissolves.
The products are seperated by repeated careful fractionation, using an efficient column. Bp. Of SOCl2, 77 C; of POCl2, 108C.
This preparation is generally not completely free of phosphorus compounds. The yeild of SOCl2 is 30g (50% of theoretical).

PROPERTIES:
Colorless, highly refractive liquid with an unpleasant, SO2-like odor. Mp. -104.5 C, bp. 77 C; d. 1.64. Significant dissociation to S2Cl2, SO2 and Cl2 occurs just above the boiling point. Hydrolyzes inwater to SO2 and Hcl; soluble in benzene and chloroform.

REFERANCES
I.A, Michaelis, Liebigs Ann. Chem. 274, 173 (1893).
II.H. Grubitsch, Anorganisch-praparative Chemie, Vienna, 1950, p. 294.

Much of the fancy glassware can probably be left out, and just distilled etc. with what you have on hand, as every synthesis in the book is on how to get super-uber-absolutely pure chemicals, which we don't need for most things.

chloric1 - 19-12-2004 at 12:30

Quote:

[quoteMuch of the fancy glassware can probably be left out, and just distilled etc. with what you have on hand, as every synthesis in the book is on how to get super-uber-absolutely pure chemicals, which we don't need for most things.

Yes but that applies if we only need to go one or two steps. You see if one is attempting a several step proceedure(advanced members) then the product at each step will be of a little less purity as the previous. Then again, if you are taking hydrated metallic chlorides and wanting to convert them to anhydrous chlorides then the phosphorus compounds impurites are helpfull if anything. Purity is application specific.

[Edited on 12/19/2004 by chloric1]

Natures Natrium - 3-1-2005 at 14:58

Quote:

Originally posted by solo
try reading post [537708] over at the hive.....

Excerp
Summary

The oxidation of SCl2 by atmospheric oxygen has been investigated using activated charcoal as the catalyst. The reaction has been discussed in terms of two reactions. Firstly, the irreversible oxidation of SCl2 to SO2Cl2 :
SCl2 + O2 __> SO2Cl2

and secondly, the reversible reaction between the SCl2 and the SO2Cl2 to form SOCl2 :

SCl2 + SO2Cl2 <__> SOCl2

This oxidation has been investigated at the temperatures 193°C, 138°C and 320°C, and the equilibrium constant of the reversible equilibrium reaction has been determined for each temperature. From the temperature dependence of this equilibrium constant it is found that the reaction is exothermic to the extent of 18 kcal.

(Yield of SOCl2 at 193°C : ~73%)

..............

Seeing as how the hive is currently unavailable, might you have all the details mentioned in that post? Or is that it? We have the temperature, but what about an actually example detailing reactant (and catalyst) amounts and a general description of the set up. Boiling S2Cl2 in a stream of air through a glass tube with activated carbon heated to 193c? If you have the details (inc. original ref.), please post them, many thanks to you.

Ref: thionyl chloride synthesis

solo - 3-1-2005 at 16:41

I haven't the posted thread , but I do have an article I can share......solo

Attachment: synthesis of thionyl chloride (trans_faraday_soc_43.).djvu (170kB)
This file has been downloaded 1933 times

Natures Natrium - 3-1-2005 at 20:06

Ah, thank you very much good sir. That information is most helpful. I dont know if I will get around to attempting it (going to try SO3 from NaHSO4 first), never the less it is good to know. :cool:

Natures Natrium - 6-1-2005 at 10:06

Something about this document has been bothering me. They show the use of SCl2, which is represented as chemically distinct from S2Cl2 (differing bp, etc), which in itself makes sense to my limited knowledge. However, I cannot find one single reference to SCl2 anywhere, not in the Merck, my textbooks, or any of a numerous amount of vendors with online catalogues. The implication in thier purity description from the lot obtained from a chem supply house is that SCl2 was made from S2Cl2 via either reaction with HCl or catalytic hydrogenation (oh, or maybe excess Cl2?). Either way, it does make this synth not quite as conveinent and useful as it might be. I guess the real question now is how easy and viable is a synthesis of SCl2 from S2Cl2?

PostScript : Put SO3 attempts on hold, this just seems too interesting.

Polverone - 6-1-2005 at 14:56

SCl2 certainly does exist. In Sidgwick's Chemical Elements and their Compounds it is said to be prepared by adding Cl2 to S2Cl2 and distilling.

Thomas Winwood - 7-1-2005 at 08:24

SCl2 is mentioned on WebElements (not the best resource for compounds, but good for some).

Natures Natrium - 8-1-2005 at 21:42

Thanks for the info guys. It made sense to me, but I feel better having other solid references to it.

Whilst I was perusing all the various posts on this site concerning thionyl chloride, I came upon this (http://www.sciencemadness.org/talk/viewthread.php?tid=790) thread. The procedure yet undisclosed (anybody ever find this ref?) states that thionyl chloride may be made by the action of SO2 on a saturated solution of HCl. Well, having some time on my hands (I ran out of MAPP gas for SO3 experiments :mad:

), I came up with this idea for a methodology in lieu of the details from the reference.

Saturate room temp anhydrous DCM (say, 500mL) with HCl. Set up for a distillation with a gas inlet reaching to the bottom of the DCM solution. Bring the DCM solution to a gentle boil, and begin a slow, steady flow of SO2 gas. Continue the gas flow until nearly all the DCM has distilled over (or the level of the DCM drops below the end of the gassing tube), cut the gas flow, but finish distilling off the DCM. If all goes well and the beanstalk grows, there should be thionyl chloride left in the flask. Change recieving flasks and distill the thionyl chloride for decent purity.

After looking at all the solvents available to me, I decided DCM would probably be the best based on several factors. One, its bp (40C) is well below that of thionyl chloride (76C), two, it doesnt react with thionyl chloride, and three, the azeotrope it forms with water (1%H2O) boils at 38.8C. The other solvent possibilities considered are CS2, isopropyl chloride, chloroform, and diethyl ether. Not that I have access to CS2, just that it has properties that would probably work in this reaction.

Anyways, just thought I would share some brainstorming. Criticism welcome.

Sincerely,
Nature's Natrium

EDIT: Oh wait, I just realized something. A 1% azeotrope means that disitlling 500mL of DCM would only remove a mere 7g of H2O, less than .5 moles. There has to be a better solvent...

EDIT2:500mL of propyl chloride or isopropyl ether would remove 15.4g. Interesting that because of density and azeotrope differences, they just happen to come out to removing the same amount of H2O per volume of solvent. Nevertheless, thats not even a mole of water. I am beginning to see the value of a dean-stark trap. On the other hand, in a perfect magic universe, .85 mol of water removal ought to equal 101g of thionyl chloride. It is becoming clear that in this variation of the idea, the amounts of HCl and SO2 would have to be tightly controlled to ensure a decent yield.

FINAL EDIT:Oh oh oh, chloroform would be even better for this. 500mL should be able to remove up to 36.75g of H2O, or roughly 2 moles. Also, its azeotrope has a lower bp than with isopropyl ether. If the chloroform is thoroughly saturated with HCl, and then about 2-2.5mol of SO2 bubbled in, the theoretical yield would be 237.9g (2 mol) of thionyl chloride. I think it is probably safe to say that is a lovely, vivacious day dream somewhat removed from reality though.

[Edited on 9-1-2005 by Natures Natrium]

[Edited on 9-1-2005 by Natures Natrium]

[Edited on 9-1-2005 by Natures Natrium]

Natures Natrium - 9-1-2005 at 08:41

Ok, got a question about activated C, since I have never really used such as a catalyst for anything. 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? Would doing so anyways help or hinder the catalysts activity? I intend to heat it under vacuum to drive off water either way.

Thanks,
Nature's Natrium

Need to find reference for Thionyl Chloride

FloridaAlchemist - 9-1-2005 at 17:19

Excess sulfur dioxide in conc Hydrochloric acid yields thionyl chloride.

SO2 + 2HCl = SOCl2 + H2O

Ref. JCS, 117, pg 1103, 1920

Just an Idea

FloridaAlchemist - 9-1-2005 at 17:28

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.

S.C. Wack - 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.

mick - 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]

jimwig - 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?

Natures Natrium - 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.

)

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

BromicAcid - 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?

Natures Natrium - 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]

SCl2

chemoleo - 21-2-2005 at 20:02

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

BromicAcid - 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?

garage chemist - 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.

enima - 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.

mick - 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]

so far, so good

Natures Natrium - 16-3-2005 at 09:54

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.

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]

BromicAcid - 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.

neutrino - 16-3-2005 at 13:51

People have said that S2Cl2 was a good solvent for sulfur. Look in the appropriate thread.

Natures Natrium - 16-3-2005 at 14:03

Quote:

Originally posted by neutrino
People have said that S2Cl2 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.

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.)

S.C. Wack - 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.

neutrino - 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.

Natures Natrium - 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

), 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]

S.C. Wack - 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.

Natures Natrium - 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.

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 877 times

garage chemist - 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]

Natures Natrium - 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.

garage chemist - 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.

garage chemist - 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]

IrC - 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.

BromicAcid - 11-4-2005 at 11:29

The reactivity of DMSO is fairly thoughly explained in the text "Technical Bulletin Reaction Solvent Dimethyl Sulfoxide" from Gaylord Chemical Corporation. According to it DMSO reacts with anhydrous HCl to give an ionic salt:

CH3SOCH3 + HCl ---> [CH3S(OH)CH3]+Cl-

Even if you could find a nice way to convert to thionyl chloride, it would readily react with the DMSO present:

CH3SOCH3 + SOCl2 ---> CH3SCH2Cl + SO2 + HCl

So really I see no easy way to it, you would have to break the S-C bonds and they are fairly stable, they will survive direct chlorination to give CH3SOCCl3. Unless the DMSO is under basic conditions halogens catalyze the decomposition to SO2 and CH3SO2CH3 and other crap. So I don't really see a feasible way to thionyl chloride through DMSO.

IrC - 11-4-2005 at 17:56

Is there a way using a catalyst to break the C-S bonds?

Oxychlorides of sulfur

Natures Natrium - 16-7-2005 at 21:50

Crap, I just wrote like 2 pages of stuff on this topic and lost it by clicking on the "review long topic" link. :mad:

Ok, let's try this again.

I have read somewhere that refluxing an equimolar amount of SO2Cl2 and SCl2 would result in conversion to SOCl2. It sounds nice, but I have my doubts, since in the trans-faraday article posted earlier in this thread they found the equilibrium of that reaction to lean toward thionyl at 193C (in the gas phase over activated carbon). I happen to have both sulfur dichloride and sulfuryl chloride on hand currently, so this would be mighty convienent. Anyone have any thoughts on this? Refluxing in the presence of activated carbon I'm not sure would work. The reason it would not work is related to my attempts at SO2Cl2.

Following the procedure outlined in Inorganic Lab. Preps., I have twiced now had success synthesizing sulfuryl chloride. The first time I took my 300mm West condenser and packed it full of my activated carbon*. Other than that it went exactly as described in the book, and I made about 30g SO2Cl2. Cleaned the excess Cl2 out (which I use in even greater excess since SO2 is a reagent that is precious to me

) by adding 3g Hg metal and shaking for a bit.

The second time I made it, I ran into a bit of a problem I did not notice in the first run but which might have been affecting my overall yield. I found that passing the gasses too quickly through the carbon, the end of my condenser (inside the 24/40 glass joint) began to get hot. This was interesting to note since the parts of the carbon which were inside the cooled part of the condenser were clearly "wet" with SO2Cl2, whereas the end carbon became dry and stayed that way. When this happened the flow of SO2Cl2 stopped entirely. I therefor hypothesize (guess) that hot activated carbon is better at splitting SO2Cl2 than combining it. (This is why I doubt the SO2Cl2 + SCl2 -> SOCl2 in the presence of carbon.) A look at US patent 5,498,400 shows that the yield of SO2Cl2 over carbon catalyst is much higher if the reaction is run at lower temps. This patent gives multiple detailed examples. I suppose it also probably says that the gas rates were too high, since the SO2 (limiting reagent) was being fully reacted before reaching the end of the tube. I then tried something a little different. I took the 500mL RBF that had been collecting the SO2Cl2, through some of the still damp activated carbon in (about an equal volume), and then proceeded to pump in chlorine gas until saturated, then SO2. I did this back and forth a number of times, stirrring/shaking by hand quite a bit, and managed to roughly triple the volume of SO2Cl2 I had collected. I ended up with a total of 59.6g SO2Cl2 after decantation from the activated carbon. So, once some is made, it is clear that a camphor like approach to this synthesis is possible. Whether or not it is to be preferred is dependent on disposition and equipment available, I suppose. It seemed to me to be a lot more effecient however, since very little Cl2 or SO2 managed to escape the SO2Cl2 liquid.

I also managed to make a bit of thionyl chloride with my bootleg catalyst tube, although the process was horribly ineffecient. In this expirement I passed Cl2 and SO2 over my activated carbon catalyst*, except this time it was heated to fairly high temperatures (guesstimation: >300C). After an hour or so I ended up with a measly 5mL of a reddish liquid. This stuff seems much nastier than SO2Cl2, SCl2, or S2Cl2. It's hard to describe the difference but essentially SOCl2 just instantly burned the crap out of the inside of my nostrils on the smallest accidental whif. The other compounds just tend to reek of SO2, Cl2, and HCl. (As a side note, I may take a long haitus from chlorides of sulfur soon, or any synth that requires chlorine...I got a whif of Cl2 earlier today and it smelled sweet and didn't physically bother me much!

) US Patent 5,759,508 also shows it is possible to obtain thionyl chloride this way, although curiously they don't state yield.

I made only one attempt at thionyl chloride based on the trans-faraday doc, and it was a waste since the carbon still contained way too much water even though it was heated on a hot plate under vacuum. Speaking of which...

* My Activated Carbon. I ended up getting this stuff truely anhydrous and activated by deciding to dump about 20mL of my SCl2 directly into the glass container which held the carbon. It boiled a bit and let out quite the stink before finally subsiding. I just let the carbon soak it all up. I have not had a chance to use it for sulfur dichloride to thionyl chloride yet, as was my original intention, although it has proven to be an excellent catalyst for SO2 + Cl2. Of course, it actually doesn't seem like it takes that much to get these two together, or break them apart.

Ok, main question: Will refluxing an equimolar mixture of SCl2 and SO2Cl2 lead to a decent yield of SOCl2?

Comments, opinions, and ideas are most welcome.

Sincerely,
Nature's Natrium