4-Stroke
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New (Simpler) Method for Making Sulfur Chlorides
Most of you know that sulfur chlorides are usually prepared by the standard procedure where chlorine gas is generated, dried, passed into molten
sulfur, and the sulfur chlorides formed get distilled as soon as they form, all at the same time. This method has a few disadvantages, such as the
need for:
-Much sophisticated glassware (a 2/3 neck RBF, stoppers, a chlorine generator, gas addition tube, either a heating mantle (expensive) or an
oil bath (messy), a distillation adapter, a condenser, and a receiving adapter & flask)
-Constant monitoring of the reaction
-The horrid cleaning process afterwards
-Longer reaction time
The method I propose removes all of these disadvantages (but note that it still requires you to have some sulfur chlorides to start with).
The only equipment needed is:
-Erlenmeyer flask (more accessible, cheaper, and easier to heat than a RBF)
-Condenser (any condenser will work, possibly even an air cooled one)
-Glass tubes (even some plastic ones will work)
-Rubber or silicone bungs
-Filter
Here is my proposed procedure:
Step 1: Setup a chlorine generator (the simplest way I could come up with):
A glass filtering flask is charged with 155 grams of TCCA (or 143 grams of Ca(OCl)2). A separatory funnel with 200ml of muriatic acid (or
270ml if Ca(OCl)2 is used) is placed on top.
Step 2: A gas dryer (to dry the chlorine) is set up. A test tube is stoppered with a bung and a glass (or a suitable plastic) tube is
inserted to the bottom. A second tube is connected that just barely enters the test tube. The test tube is filled with anhydrous (or monohydrate)
calcium chloride.
Step 3: 70ml (1 mole, 135 grams) of S2Cl2 is charged into an Erlenmeyer flask. A glass (or a suitable plastic)
tube is placed into the S2Cl2.
Step 4: The entire apparatus is assembled.
Step 5: The muriatic acid is allowed to drip in and the evolution of chlorine is started. The chlorine starts bubbling through the
S2Cl2.
Step 6: After all of the muriatic acid has been used up, the S2Cl2 should have changed color from yellow to red
as it was converted into 2 moles (205 grams, 127ml) of SCl2.
Step 7: The chlorine generator, dryer, and input tube is disconnected and removed. The Erlenmeyer flask is charged with an excess
(>65 grams) of sulfur, fitted with a condenser, and set up for reflux.
Step 8: The solution is refluxed until it changes color to back to yellow and most of the sulfur is gone.
Step 9: The condenser is removed and the solution is filtered to remove the unreacted sulfur.
Step 10: Done! You should now have more S2Cl2 than you started with. Theoretically, 1 mole (135 grams,
70ml) of S2Cl2 --> 2 moles (205 grams, 127ml) of SCl2 --> 2 moles (270 grams, 140ml)
S2Cl2.
I understand that the main problem with this method is getting any sulfur chlorides to start with, but it could be made by a method like this (https://www.youtube.com/watch?v=CgNxJHW513g) on a test-tube scale (or full scale if you have the equipment).
What do you think? Will it work?
[Edited on 11-11-2024 by 4-Stroke]
Ukrainium
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Tsjerk
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At least, I never worked with sulfur chlorides, you will have problems with suck back, and I guess your drying apparatus won't be sufficiently large.
Are you sure your method will be as efficient as you propose here?
Usually, especially with these will known reactions, there is a reason why they are done like the are done
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Bedlasky
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Sulfur reacts with S2Cl2 to form polysulfur dichlorides SxCl2. So any excess sulfur will react with your S2Cl2. If you want pure S2Cl2, you must
distill it to remove unreacted sulfur (SxCl2 break up during distillation back to S2Cl2 and sulfur).
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bnull
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Well, you still need stoppers, a chlorine generator, gas addition tube, some means to heat the stuff in order to purify it (see Brauer's Handbook of
Preparative Inorganic Chemistry, p. 370-372), a distillation adapter, a condenser, and a receiving adapter & flask. And sulfur chloride to begin
with.
Cleaning sulfur compounds from plastics, if you still intend to use them, is a mess.
In picture 4, you need a trap (I forgot how it is called) after the drying tube to prevent suck back. For the drying tube, a flask with concentrated
sulfuric acid. To make it even safer, two traps after the drying tube.
Filter the stuff? No, thanks.
It may work if you make some improvements. But even so I'll stick to the standard procedure. I'm a natural-born pessimist; I tend to
see what can go (very) wrong. There is an alternative process described somewhere in old publications, such as those journals of practical chemistry,
maybe even in the Journal of Chemical Education (probably before the 1980s).
Edit: Typo.
[Edited on 12-11-2024 by bnull]
Quod scripsi, scripsi.
B. N. Ull
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Boffis
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I would have thought that you will also need a vent with a chlorine scrubber on the final flask, at least to start with or you will never get the
chlorine into the final flask.
I have to say that I don't see any real advantage over the original process apart from slightly simpler and cheaper equipment. Maybe you shouldn't be
trying this reaction if you dont have the right equipment. Also chlorine attacks rubber stoppers terribly and you still have the chlorine delivery
pipes, drying jar etc.
I thought you might have come up with a one pot process, say adding TCCA to a chlorinated solvent containing a suspension of sulphur and bubbling HCl
gas through it 
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Monoamine
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It's a nice idea, perhaps if you want more S2Cl2. Also, really cool images! It's clear you put a lot of thought and effort into these illustrations.
Cudos!
But with the exact same setup you can also just generate S2Cl2 directly by blowing the chlorine into elemental sulphur. And that way, you don't need
to already have S2Cl2 on hand.
Actually, since you are refluxing the mixture in the last step anyway, I don't see how this will be less work or use less equipment than just
generating S2Cl2. Both approaches would be about the same amount of work and they use the same equipment as you describe in your process.
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4-Stroke
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Quote: Originally posted by Tsjerk  | At least, I never worked with sulfur chlorides, you will have problems with suck back, and I guess your drying apparatus won't be sufficiently large.
Are you sure your method will be as efficient as you propose here?
Usually, especially with these will known reactions, there is a reason why they are done like the are done |
Nope, I'm absolutely not sure if it is effective (or even viable at all), but that's why I'm asking for advice. And you're right, there probably is a
reason, but different methods are often a balance of time, yield, and the required equipment, but because different people need different things, some
methods might be "better" for some than for others.
| Quote: Originally posted by Bedlasky | | Sulfur reacts with S2Cl2 to form polysulfur dichlorides SxCl2. So any excess sulfur will react with your S2Cl2. If you want pure S2Cl2, you must
distill it to remove unreacted sulfur (SxCl2 break up during distillation back to S2Cl2 and sulfur). |
Oh, that's interesting. I haven't heard of that before. But even some, the procedure could be repeated many times (each time chlorinating the
polysulfur dichlorides all the way to SCl2) and then distilled only once at the end to yield pure S2Cl2.
| Quote: Originally posted by bnull |
Well, you still need stoppers, a chlorine generator, gas addition tube, some means to heat the stuff in order to purify it (see Brauer's Handbook of
Preparative Inorganic Chemistry, p. 370-372), a distillation adapter, a condenser, and a receiving adapter & flask. And sulfur chloride to begin
with.
Cleaning sulfur compounds from plastics, if you still intend to use them, is a mess.
In picture 4, you need a trap (I forgot how it is called) after the drying tube to prevent suck back. For the drying tube, a flask with concentrated
sulfuric acid. To make it even safer, two traps after the drying tube.
Filter the stuff? No, thanks.
It may work if you make some improvements. But even so I'll stick to the standard procedure. I'm a natural-born pessimist; I tend to
see what can go (very) wrong. There is an alternative process described somewhere in old publications, such as those journals of practical chemistry,
maybe even in the Journal of Chemical Education (probably before the 1980s).
Edit: Typo.
[Edited on 12-11-2024 by bnull] |
Fine, you still do need some equipment, but NOT a distillation adapter, receiving adapter, or a receiving flask. An Erlenmeyer flask is also a lot
easier to heat than a RBF.
Could xylene be used to clean plastic and glass tubes (since sulfur is soluble in it)?
What's so bad about filtering it apart from the smell?
The main point of the procedure is that it's a lot easier (and I suspect faster) than the conventional method. It is also easier to perform on a
bigger scale (since less equipment is required). It could also be performed in a lead-lined reaction vessel (anything from a small pot to a 55 gallon
drum covered with lead sheet inside) a lot more easily. Also less glassware to clean at the end 
But thanks for the references, I'll take a look at those.
| Quote: Originally posted by Boffis | | I would have thought that you will also need a vent with a chlorine scrubber on the final flask, at least to start with or you will never get the
chlorine into the final flask. |
I'm a little confused by what you are saying here. Do you mind explaining?
| Quote: Originally posted by Boffis | | I have to say that I don't see any real advantage over the original process apart from slightly simpler and cheaper equipment. Maybe you shouldn't be
trying this reaction if you dont have the right equipment. |
I actually do have most of the equipment required the standard procedure, but the goal of this one is that it is faster (I think), requires cleaning
less glassware, is a lot easier to perform on a larger scale in a lead lined reactor, and is just generally less of a hassle (working with molten
sulfur, chlorinating it, an distilling simultaneously vs just bubbling chlorine, adding sulfur, refluxing, and repeat).
| Quote: Originally posted by Boffis | | I thought you might have come up with a one pot process, say adding TCCA to a chlorinated solvent containing a suspension of sulphur and bubbling HCl
gas through it |
That's actually quite an idea I'll think a bit more about that.
| Quote: Originally posted by Monoamine | | It's a nice idea, perhaps if you want more S2Cl2. Also, really cool images! It's clear you put a lot of thought and effort into these illustrations.
Cudos! |
Thank you! Finally someone noticed
| Quote: Originally posted by Monoamine | | But with the exact same setup you can also just generate S2Cl2 directly by blowing the chlorine into elemental sulphur. And that way, you don't need
to already have S2Cl2 on hand. |
Yes, but this setup is still faster, easier to clean, and is much simpler to perform in a lead lined reaction chamber.
| Quote: Originally posted by Monoamine | | Actually, since you are refluxing the mixture in the last step anyway, I don't see how this will be less work or use less equipment than just
generating S2Cl2. Both approaches would be about the same amount of work and they use the same equipment as you describe in your process.
|
Yes but this process kind of like separates the entire reaction into different sections, so each one separately is easier to perform than all at once.
Ukrainium
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bnull
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| Quote: | | Fine, you still do need some equipment, but NOT a distillation adapter, receiving adapter, or a receiving flask. An Erlenmeyer flask is also a lot
easier to heat than a RBF. |
Nope. You need ALL the original equipment plus the erlenmeyer, the funnel, and the two traps (see below). The chloride is purified by distillation,
hence the need for the three pieces of glassware that you intend to ditch. You won't use the crude stuff that you have just filtered, will you?
| Quote: | | Could xylene be used to clean plastic and glass tubes (since sulfur is soluble in it)? |
Glass, yes, obviously; plastic, no. Plastics are usually very fond of solvents. Using xylene to remove sulfur from plastic tubing would make some
sulfur diffuse into it.
| Quote: | | The main point of the procedure is that it's a lot easier (and I suspect faster) than the conventional method. It is also easier to perform on a
bigger scale (since less equipment is required). |
Try it small scale first. Better yet, perform both processes, the standard and the one you devised, and compare them. This way you can see which one
is better, if there is something you missed.
| Quote: | | It could also be performed in a lead-lined reaction vessel (anything from a small pot to a 55 gallon drum covered with lead sheet inside) a lot more
easily. Also less glassware to clean at the end ;) |
Why would one need that much sulfur chloride?
| Quote: | | But thanks for the references, I'll take a look at those. |
Please do. Brauer's is a classic.
I understand you want to improve processes and make the life of amateur/home/hobby chemists easier. But hold on: how much sulfur chloride does an
amateur chemist need? Not gallons of the stuff, I believe (unless the chemist is making mustard gas).
My advice is, consult the literature, especially the old books on preparative chemistry (we have lots of them in the Library). There is always a
reason why the standard procedure became the standard procedure.
The pictures are pretty good, by the way. What program did you use?
Quod scripsi, scripsi.
B. N. Ull
P.S.: Did you know that we have a Library?
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4-Stroke
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| Quote: Originally posted by bnull | | Nope. You need ALL the original equipment plus the erlenmeyer, the funnel, and the two traps (see below). The chloride is purified by distillation,
hence the need for the three pieces of glassware that you intend to ditch. You won't use the crude stuff that you have just filtered, will you?
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Well, that depends on what you need it for (distillation might not be necessary for some reactions). Alternatively, if sulfur dichloride is desired,
the filtered mixture of polysulfur dichlorides and disulfur dichloride can be chlorinated all the way to SCl2, and since (I assume) there
are no "more chlorinated" sulfur chlorides, it should be pretty pure SCl2.
| Quote: Originally posted by bnull | | Glass, yes, obviously; plastic, no. Plastics are usually very fond of solvents. Using xylene to remove sulfur from plastic tubing would make some
sulfur diffuse into it. |
Ah yes, I haven't thought of that. Many plastics could even dissolve in solvents, so yeah, that was a pretty dumb suggestion from me.
| Quote: Originally posted by bnull | | Try it small scale first. Better yet, perform both processes, the standard and the one you devised, and compare them. This way you can see which one
is better, if there is something you missed. |
Yes, I will, but we are moving soon, so I can't really do much for the next couple month, that's why I just decided to post this here. But yes, when I
can, I definitely will try both methods, I've been wanting to do so for quite a while now.
| Quote: Originally posted by bnull | | I understand you want to improve processes and make the life of amateur/home/hobby chemists easier. But hold on: how much sulfur chloride does an
amateur chemist need? Not gallons of the stuff, I believe (unless the chemist is making mustard gas). |
As much as the chemist wants Well just recently I was asking about lead (and
other materials) lined reaction vessels, so I just had to include this here.
| Quote: Originally posted by bnull | | My advice is, consult the literature, especially the old books on preparative chemistry (we have lots of them in the Library). There is always a
reason why the standard procedure became the standard procedure. |
Yes, by now I'm pretty much running out of compelling arguments as to why my procedure is worth it (probably because it isn't ).
But I just wanted to ask, what do you think of this (https://www.youtube.com/watch?v=CgNxJHW513g) method? I found it some time ago, and it seems pretty good to me. Just mix TCCA with sulfur and
distill. He used the wrong (stoichiometrically) ratio though. What do you think?
Also, another question, why is sulfur chloride produced industrially by bubbling chlorine through barely molten (~120°C) sulfur, but in the standard
laboratory preparation method, the sulfur is heated much more (~250°C) and thus the sulfur chloride is distilled immediately as it is formed.
Considering that there is about a ~20°C range between sulfur's melting point and disulfur dichloride's boiling point, wouldn't it be easier to just
chlorinate the molten sulfur and distill it separately?
Thanks! It isn't a program, just the first website that came up. It's https://chemix.org/.
[Edited on 12-11-2024 by 4-Stroke]
Ukrainium
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bnull
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Apart from the sulfur fumes escaping the receiver a short distance from an open flame (O_o) and the possibility of the whole thing blowing up, it is
interesting. I wouldn't do it, but it is interesting.
| Quote: | | Also, another question, why is sulfur chloride produced industrially by bubbling chlorine through barely molten (~120°C) sulfur, but in the standard
laboratory preparation method, the sulfur is heated much more (~250°C) and thus the sulfur chloride is distilled immediately as it is formed.
Considering that there is about a ~20°C range between sulfur's melting point and disulfur dichloride's boiling point, wouldn't it be easier to just
chlorinate the molten sulfur and distill it separately? |
This is one thing I have no idea (at least at the moment). You may check these sections on sulfur monochloride and sulfur dichloride from PubChem.
| Quote: | | Thanks! It isn't a program, just the first website that came up. It's https://chemix.org/. |
Just added it to my bookmarks.
Quod scripsi, scripsi.
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4-Stroke
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| Quote: Originally posted by bnull |
Apart from the sulfur fumes escaping the receiver a short distance from an open flame (O_o) and the possibility of the whole thing blowing up, it is
interesting. I wouldn't do it, but it is interesting. |
Well, I've boiled gasoline in an open pot over an open flame before, so... But it does seem pretty stable (and the guy says he performed the reaction
"dozens" of times) so I think that it shouldn't blow up. But I'll be the one to check that in a couple month.
Ah I see, that's fine. I'll check those links.
Ukrainium
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Texium
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So… you’re an idiot?
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4-Stroke
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Well, there wasn't a lot of it, and I was kind of prepared for it to ignite (which it surprisingly didn't). But yes, I guess you could say so
But the great chemists of the 19th century distilled diethyl ether with an open flame (albeit not in an open container), so I don't think
that what I did was really that dangerous (considering that I was also doing it outside).
[Edited on 13-11-2024 by 4-Stroke]
Ukrainium
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clearly_not_atara
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| Quote: | | I thought you might have come up with a one pot process, say adding TCCA to a chlorinated solvent containing a suspension of sulphur and bubbling HCl
gas through it |
In theory, following the vanadium-chlorine cycle for thermochemical water splitting, you could follow the wacky chemistry of the vanadium chlorides.
- VCl3 can be dried by heating
- at >350 C, VCl3 (l) disproportionates into VCl2 (s) and VCl4 (g)
- at low temperatures, VCl4 (l, bp 148 C) slowly disproportionates into VCl3 (s) and Cl2 (g)
The equilibrium pressure of chlorine over VCl4 is pretty low, but it continually releases chlorine. The unusual disproportionation is apparently
driven by the lattice energy of crystalline VCl3. This was considered decades ago for hydrogen production from water, but probably isn't practical for
mass production because there are too many steps.
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6dthjd1
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I hope this was done outside.
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6dthjd1
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| Quote: Originally posted by 4-Stroke | Well, there wasn't a lot of it, and I was kind of prepared for it to ignite (which it surprisingly didn't). But yes, I guess you could say so
But the great chemists of the 19th century distilled diethyl ether with an open flame (albeit not in an open container), so I don't think
that what I did was really that dangerous (considering that I was also doing it outside).
Do you remember how much they distilled?
[Edited on 13-11-2024 by 4-Stroke] |
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6dthjd1
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| Quote: Originally posted by 4-Stroke | Well, there wasn't a lot of it, and I was kind of prepared for it to ignite (which it surprisingly didn't). But yes, I guess you could say so
But the great chemists of the 19th century distilled diethyl ether with an open flame (albeit not in an open container), so I don't think
that what I did was really that dangerous (considering that I was also doing it outside).
Do you remember how much they distilled?
[Edited on 13-11-2024 by 4-Stroke] |
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Sulaiman
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Bofis:
would have thought that you will also need a vent with a chlorine scrubber on the final flask, at least to start with or you will never get the
chlorine into the final flask.
4-stroke:
I'm a little confused by what you are saying here. Do you mind explaining?
.......,...............
The stopper in the Erlenmeyer used for bubbling has only one hole,
there is nowhere for excess gas to escape so internal pressure will increase
the pressurised chlorine gas will bubble up through the addition funnel the next time the stopcock is opened...
If the pressure is high a violent eruption of HCl and chlorine will occur.
Not good
CAUTION : Hobby Chemist, not Professional or even Amateur
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