Sciencemadness Discussion Board - Carbon disulfide synthesis (alternative routes)

I went through Reaxys a few months ago and compiled this:

Thanks a bunch for that! I shall look into this tomorrow.

wg48temp9 - 31-3-2019 at 13:28

I went through Reaxys a few months ago and compiled this:

I stripped out all C + S syntheses, and any ridiculous or unfriendly (to the hobby chemist, eg complex organic molecules) syntheses

There are some apparently simple methods there that use OTC reagents except for the CCl4 for example With S, iron ~(or aluminium or copper) in tetrachloromethane, at 120°C.

So apart from the steel pressure bomb to heat the CCl4 to 120C what if any are the other problems?

BromicAcid - 31-3-2019 at 15:57

There are whole mess of thiuram monosulfide and disulfides. Fungicides are the main use hence I mentioned OTC. I have seen them in dusting powders at various times for gardening but usually as now more than 10%, hence me mentioning small amounts. Most of the thiuram compounds I have run across are thermally unstable and at reasonable (100-250°C) temperatures undergo thermal decomposition to yield thioureas and carbon disulfide which distills out of the reaction mixture.

Assured Fish - 31-3-2019 at 20:20

There is a theoretical method starting from thiocyanate that could work.
It has been discussed in a previous thread, and another thread discussing the percyanic acid compound:
http://www.sciencemadness.org/talk/viewthread.php?tid=103650...
http://www.sciencemadness.org/talk/viewthread.php?tid=27676#...

It begins with the preparation of iso perthiocyanic acid, which involves treating a thiocyanate salt with an acid with a concentration over 40%.
3HCNS = HCN + H2C2N2S3

I have done this before and have a few grams of the supposed perthiocyanic acid sitting in a mason jar.

Alas this stuff gives off nitrous oxide whenever i open the jar, i believe it decomposes if the acid is too concentrated as i also observed large copious amounts of NO2 being given off during a few test runs to make it.
Its likely a small residual amount of acid is still present on this sample and when its exposed to the air, moisture gets in contact and causes the decomposition.
This stuff was made using ammonium thiocyanate and nitric acid, the exact yield of which i cannot say as its the collection of multiple small test tube scale experiments.
All I can say is that its rather difficult to filter, and can stain glassware if not cleaned off within a day or so.
Other than that though its very easy to make.

From the iso perthiocyanic acid it would be as simple as setting up a 3 neck round bottom flask in an oil bath with a still head and condenser in one adapter, and a thermometer in the other reaching down into the bottom of the flask and an inlet tube to gas the system with a steady flow of nitrogen or argon.
The oil bath can then be heated up to 180*C to 200*C and then the distillate collected, which can then be washed with water and redistilled hopefully collecting something at around 46*C.
I would run this myself but i dont want to do it in my current residence as preparing a highly volatile somewhat toxic, extremely flammable liquid that i cannot store is quite a dangerous exercise.

Keras - 1-4-2019 at 03:36

Quote: Originally posted by Assured Fish

There is a theoretical method starting from thiocyanate that could work.
It begins with the preparation of iso perthiocyanic acid, which involves treating a thiocyanate salt with an acid with a concentration over 40%.
3HCNS → HCN + H2C2N2S3

I have done this before and have a few grams of the supposed perthiocyanic acid sitting in a mason jar.

I was wondering about this article: F. Cataldo, Y. Keheyan: About a new class of inorganic polymers: the polythiocyanogens, in Polyhedron 21 (2002), especially: "It is shown that all the pseudohalogens belonging to the series of sulfur dicyanide S_y (CN)₂ or thiocyanogens polymerise very easily in the dry state, under moderate heating to produce a series of brick-red or orange solids known as polythiocyanogens with the general formula (S_y (CN)₂)_x with y an integer 1, 2, 3, 4."

Could it be what you obtained?

[Edited on 1-4-2019 by Keras]

Keras - 1-4-2019 at 04:20

I went through Reaxys a few months ago and compiled this:
I stripped out all C + S syntheses, and any ridiculous or unfriendly (to the hobby chemist, eg complex organic molecules) syntheses

Ok, I don’t have Reaxys and seem to have only limited access to Gmelin; couldn't find most of them, but the one using carbon tetrabromide seems interesting. However, until my batch of potassium bromide arrives, I only get access to pure iodine. I wonder if that could work with carbon tetraiodide…

[Edited on 1-4-2019 by Keras]

Assured Fish - 1-4-2019 at 20:17

Quote:

Could it be what you obtained?

After reading through that paper, i can safely say that i very much doubt that is at all what ive obtained, or that those polymers are what others have obtained before me.

For starters the reaction conditions used for the preparation of all those polymers were anhydrous, while they were washed with water in the work up in a few examples.
The conditions described don't even come close to what ive recreated.
Their paper discusses preparing compounds such as sulfur dicyanide which would require the thiocyanate to be reduced.
It just wouldn't make any sense.

The compound iso-perthiocyanic acid is a a differnet compound all together with a somewhat unknown structure.
https://sci-hub.tw/https://link.springer.com/article/10.1007...
Actually having just found this ref, i should probably try to test this given that ive got some sodium nitrite on hand.

There are a few papers on the subject but to get any results with google i found that using the name isoperthiocyanic acid yields better results than iso-perthiocyanic acid.

Quote:

I only get access to pure iodine. I wonder if that could work with carbon tetraiodide

This is not a good idea, i mean in theory you could attempt free radical halogenation of iodoform but ive never seen an example of this occurring which may be to do with carbon tetraiodide being rather unstable and the conditions required for free radical halogenation are both thermal and photochemical which the wiki states decomposes carbon tetraiodide back to iodoform and iodine.
Meaning you may end up with a very small amount of CI4 but the equalibrium lies heavily on the iodoform side.
https://en.wikipedia.org/wiki/Carbon_tetraiodide
Metasynthesis is instead the go to approach for preparing CI4, but even then its very expensive to make.
CCl4 on the other hand has been prepared on the forum:
http://www.sciencemadness.org/talk/viewthread.php?tid=14927

That thread should be moved to prepublications.

Keras - 2-4-2019 at 08:17

Coming back to CS₂, I found that it could also be produced by decomposition of thiourea at 230 °C.

Ref: Wang, Zerong Daniel; Hysmith, Meagan; Quintana, Perla Cristina: Theoretical study on the formation of carbon disulfide and ammonia from thermal decomposition products of thiourea (DOI:10.1142/S0219633614500229).

Thiourea is real cheap, so that could be a really convenient way to get cheap CS₂!

[Edited on 2-4-2019 by Keras]

clearly_not_atara - 2-4-2019 at 14:09

I went through Reaxys a few months ago and compiled this:

The most interesting one in here IMHO is the method using CuCN + S + heat, and the very similar Fe(SCN)3 + heat; if I can guess that the reaction of sulfur with a transition metal cyanide at a high enough temperature generates (perhaps transiently) the thiocyanate, they are the same: thermal decomposition of a transition metal thiocyanate. This gives a lot of possibilities for the thiocyanate and in particular CuCN can be prepared from OTC (and safe) ferr[oi]cyanides.

The similar methods using pyrolysis of thiourea are more likely to produce other gases as well (H3N might react with CS2 on cooling, and any formation of H2S is bad), while CS2, S2 and N2 should be the only significant volatile products from pyrolysing a metal thiocyanate (cyanogen and thiocyanogen should not be stable at these temperatures). (Transition metal thiocyanates should have lower decomposition temperatures than other thiocyanates.)

[Edited on 2-4-2019 by clearly_not_atara]

Keras - 15-5-2021 at 07:42

Coming back to CS₂, I found that it could also be produced by decomposition of thiourea at 230 °C.

Ref: Wang, Zerong Daniel; Hysmith, Meagan; Quintana, Perla Cristina: Theoretical study on the formation of carbon disulfide and ammonia from thermal decomposition products of thiourea (DOI:10.1142/S0219633614500229).

I’m sort of resurrecting or unearthing this thread because I would like to try this method. Since CS₂ auto-ignites at around 200 °C in air, the whole setup would be conducted under inert (argon) atmosphere. Planned setup is this: two-neck bottom flask, one neck fitted with an argon supply, second neck connected to a distillation apparatus, whose exit leads to two cascaded Drechsel bottles, the first being filled with mineral oil to avoid air getting in, the second with 10% bleach, since I read thermal decomposition of thiourea can also produce a small amount of HCN, even though not much is evolved at the planned temperature (circa 220 °C).

The RBF is flushed with Ar, then heated, with excess air bubbling out through the Drechsel bottles.

I quote: The main gaseous products of thiourea in pure Ar atmosphere between 182 and 330°C are found by TG-FTIR to be carbon disulfide (CS₂) and ammonia (NH₃) indicating decomposition of neighbouring thiourea molecules in the melt: 2 SC(NH₂)₂ (l) → CS₂(g) + 2 NH₃(g) + H₂NCN(l) Evolution of other gaseous species as cyanamide (H₂NCN), hydrogen cyanide (HCN), carbon dioxide (CO₂), and carbonyl sulfide (COS) are observed only at temperatures higher than 500°C.

That gives quite a margin of security, but one never knows. Everything is of course going to be done outside, and a multi-metre gas tube will be used to exhaust the gases out of the second Drechsel bottle far away from the setup.

Temperature should be stabilised around 220 °C, which gives the greatest output of CS₂.

Of course, dismantling the apparatus will be somewhat tricky, but I plan to do it by reopening the Ar source, disconnecting the stillhead, then placing a stopper on the flask and stopping the Ar flow, until the flask is cold enough to be opened without risking residual CS₂ self ignition.

Any comments on this blueprint?

EDIT: Since the other gas evolved is ammonia, using bleach might not be a good idea (chloramines formation)?

[Edited on 15-5-2021 by Keras]

Jenks - 15-5-2021 at 08:12

Since CS₂ auto-ignites at around 200 °C in air, the whole setup would be conducted under inert (argon) atmosphere.

According to wikipedia the auto-ignition temperature is 102C. A very old demo is to ignite carbon disulfide with steam or a boiling-hot surface. The other old demo was to pour a mixture of carbon disulfide and carbon tetrachloride into one's hand and set fire to it. Apparently the low heat of combustion of carbon disulfide, along with the cooling from evaporation of the tetrachloride, prevented being burned.

Keras - 15-5-2021 at 08:19

Quote: Originally posted by Jenks

According to wikipedia the auto-ignition temperature is 102C.

Oops. That’s lower than I suspected. The paper says red flames begin to flash at the surface of the liquid (when heated in air) when the temp reaches 200 °C, but that of course is the threshold for CS₂ formation.

Quote: Originally posted by Jenks

The other old demo was to pour a mixture of carbon disulfide and carbon tetrachloride into one's hand and set fire to it. Apparently the low heat of combustion of carbon disulfide, along with the cooling from evaporation of the tetrachloride, prevented being burned.

Didn't know about that one. My main problem was the rather extended explosive range of CS₂. But if operated under argon, everything should go smoothly.

Also, I was thinking about inserting another Drechsel bottle between the paraffin and the bleach one, this time filled with HCl in order to trap the ammonia and avoid it reaching the bleach.

S.C. Wack - 15-5-2021 at 12:22

Of course, I’m well aware of CO insidious and very acute toxicity

There are alarming neurological reports in some of the earliest industrial hygiene work for CS2 as well. (BTW it's also responsible for the almost complete 35-year gap in cannabinoid chemistry after CBN was first purified, killing one of the discoverers by fire)

I'd want to ask someone who "found by TG-FTIR" CS2 if they think it's feasible to produce it in that way (I'm guessing no). If pyrite or S and C is too big of a hassle and H2S is OK, wouldn't it be sensible to generate it from methane/ethylene/acetylene and S/pyrite. (strong cooling of the exit gas would be needed)

[Edited on 15-5-2021 by S.C. Wack]

macckone - 15-5-2021 at 21:13

Bryce, W. A.; Hinshelwood, Cyril (1949). 707. The reaction between paraffin hydrocarbons and sulphur vapour. Journal of the Chemical Society (Resumed), (), 3379–. doi:10.1039/JR9490003379

This produces a lot of crud but the three main products are hydrogen sulfide, unsaturated hydrocarbons and carbon disulfide.

Bubbling acetylene vapor through molten sulfur seems like a viable method. Acetylene can be easily produce from calcium carbide. This will reduce the amount of hydrogen sulfide that is produced. 77% of the product is carbon disulfide at 325C. I would expect you could add a catalyst and increase the reaction rate. The paper below tried iron and iodine with little effect. Silicate might help. This is one of those research topics.

http://www.sciencemadness.org/talk/files.php?pid=153376&...

Regardless of method flammability is an issue. I would recommend carbon dioxide as an inert gas as it is easy to produce and far cheaper for the home chemist than argon.

Keras - 15-5-2021 at 22:14

I took a gander at the article but never saw CS₂ mentioned. I just skimmed over it, though.

Bubbling acetylene vapor through molten sulfur seems like a viable method. Acetylene can be easily produce from calcium carbide. This will reduce the amount of hydrogen sulfide that is produced. 77% of the product is carbon disulfide at 325C.

An interesting method, I agree. Well, I have an old stock of calcium carbide that should do the trick. If the method I want to try fails, I’ll try this alternate route.

Regardless of method flammability is an issue. I would recommend carbon dioxide as an inert gas as it is easy to produce and far cheaper for the home chemist than argon.

I happen to have a bottle of argon, that is quite easily found over here amongst welding equipments. Actually, CO₂ would need me to set up another flask, whereas I just have to open the valve of my argon bottle. Besides, in the method I want to try, very little argon is used: once at the beginning to flush out the air present in the apparatus; the rest is performed in a closed environment, where gases can escape but not enter, so no need to provide a continuous flow. At the end, some more is needed, but only in order to ensure that 1. no back flow happens when the heat source is shut off; 2. CS₂ still present in the flask will not ignite when oxygen rushes in.

unionised - 16-5-2021 at 02:17

CO2 is a long way from inert, especially in the presence of hot reducing agents.

garphield - 17-5-2021 at 18:46

Perthiocyanogen decomposes on relatively mild heating to give carbon disulfide, and can be made via the oxidation of thiocyanate. The only problem is that thiocyanates are not super easy to buy and making them requires cyanide. I remember reading a Soviet patent talking about a polymer being generated from urea and sulfur chlorides, maybe the thermal decomposition of that would give some carbon disulfide?

Attachment: SU129816A1.pdf (75kB)
This file has been downloaded 296 times

[Edited on 18-5-2021 by garphield]

clearly_not_atara - 17-5-2021 at 19:05

Suppose you react a thiocyanate salt with S2Cl2:

2 KSCN + S2Cl2 >> 2 KCl + NCS-S-S-S-CN?

Surely the latter must rearrange to CS2 + N2, right?

garphield - 17-5-2021 at 20:43

Maybe. Another option might be the reaction between cyanamide and SxCl2, which might be better since it's apparently available cheaply as a fertilizer (although I have not seen it anywhere, if anyone knows where to order it online it would be nice to know). I know SxCl2 adds across alkenes but it may or may not react in the same way with the triple bond in cyanide/nitriles. Apparently the cyanamide ion in calcium cyanamide (CaCN2) has the structure (-)N=C=N(-), with each nitrogen atom having a charge of -1 and a double bond to the central carbon atom. The immediate product from the reaction of that with SCl2 might be Cl-N=C=N-Cl (plus CaS), N≡C-NCl2, N≡C-S-S-C≡N, or Cl-S-N=C=N-S-Cl. The latter might react with another molecule of CaCN2 to form cyclic -S-N=C=N-S-N=C=N-(bonded to first S), although I'm completely guessing at all of these. Still it might be worth trying, especially when compared to the urea-sulfur chloride polymer (which would be (-NH-CO-NH-S-)n, with no carbon-sulfur bonds while the products of CaCN2 + SxCl2 might have some).

Keras - 18-5-2021 at 12:13

So, I tried the pyrolysis of thiourea today.

Set up was as described in my last message: argon bottle, double-neck RBF with thiourea inside, distillation setup (into a small RBF immersed into ice cold water that was also used for the condenser) with hose leading to a first Drechsel bottle filled with mineral oil (to prevent air for getting in) and then a second one filled with 23% hydrochloric acid to trap the ammonia into ammonium chloride (see picture).

The system was purged 2 min with argon before warming the RBF.

At this point, I should’ve been alarmed by the droplets which condensed on the RBF sides (obviously water).

As soon as the thiourea melted (and maybe a bit before that), there was a slight odour of sulphur? hydrogen sulphide? it was not strong, and transient, on/off like. But as soon as the thiourea reached 200 °C, a very blatant layer of sulphur began to appear on the sides of the still-head. It spread a lot, all along the condenser (didn't clog it, hopefully, there was still space for the vapours to pass in the middle) down to the flask and as far as the first Drechsel bottle, at which point it was nothing much than traces of colloidal sulphur (but the oil became cloudy with it). Meanwhile, there was a strong bubbling and unmistakable mist of ammonium chloride above the hydrochloric acid in the second Drechsel bottle (so ammonia definitely produced). I never detected any ammonia like odour, though, so I suppose it was all trapped into the bottle, though I never witnessed that particular way in which ammonia bubbles disappear in the bottle as the climb up before reaching the liquid interface.

After 30 min or so of constant heating at 210/230 °C, I stopped the heating mantle. I let it cool a bit, then switched the argon source back on, to make up for depressurisation. I then separated the RBF from the still-head, closed it and didn't reopen it until it was way under 100 °C (one never knows).

Almost nothing into the receiving flask, I'd say 1 ml at most. Water probably.

So I suspect I should’ve dried my thiourea before using it (though it was a new batch). Probably all the carbon disulphide reacted with water vapour to form hydrogen sulphide and carbon dioxide which bubbled out with ammonia, explaining why the bubbles never disappeared into solution. Now I have no idea how sulphur was oxidised back to S(0) which deposited on the glass. Does H₂S react with NH₃ to form sulphur and something else (what?)? That seems quite unlikely.

There were, of course, no flames observed.

I had no time to completely boil off the contents of the second Drechsel bottle (NH₄Cl); but starting with 200 ml of 23 % HCl down to less than 50 ml, when cooled there was a lot of white crystals that crashed out. I’ll carry on and plan on weighing the ammonium chloride to get an idea of the yield (at least in ammonia).

So: failure to get any carbon disulphide. But ammonia seems to have made it, and the other positive point was the ability to setup that system and run it under inert atmosphere without any hitch. Next time, I’ll dry the thiourea throughly before beginning to melt it.

PS: Does anyone know how to fully remove mineral oil from the sides of a flask? I washed with toluene and xylene, but there seems to be a sticky film that won’t go away. Also, I had to boil toluene and then to wash with acetone to get rid of all the sulphur inside the distillation system.

[Edited on 18-5-2021 by Keras]

[Edited on 18-5-2021 by Keras]

[Edited on 18-5-2021 by Keras]

[Edited on 18-5-2021 by Keras]

Opylation - 21-5-2021 at 18:34

Piranha solution should do the trick. Any organic substances that don't get washed away with solvents can be cleaned with concentrated sulfuric acid and 30% H2O2

Keras - 21-5-2021 at 22:32

Piranha solution should do the trick. Any organic substances that don't get washed away with solvents can be cleaned with concentrated sulfuric acid and 30% H2O2

Unfortunately, being in the EU, 30% H₂O₂ is not an option. Fortunately, I've found that dissolving sodium percarbonate directly into conc. sulphuric acid was a more than acceptable substitute.

Thanks for that. Since I intend to try the reaction again with a fully dried thiourea, I won’t wash my bottle to the utmost cleanness, but I’ll do that, eventually.

Opylation - 22-5-2021 at 19:42

You can make your own 30% H2O2 using the 3% from the store. It may not be as cheap as getting a jug of 30%, but you can boil the volume of liquid down to 10% it's original volume for close to 30%. I've done it without issue. I usually try to keep it at a simmer or very light boil, not rolling boil. It may degrade a bit during this process, but the result is plenty strong enough for piranha. Trust me

EDIT: also, make sure to do this in some very clean glassware. No metal containers as any bit of iron or other transition metals will catalyze decomposition of H2O2

[Edited on 23-5-2021 by Opylation]

Keras - 22-5-2021 at 22:51

You can make your own 30% H2O2 using the 3% from the store. […] Trust me

Oh, I know! We can even get 12% here, so 1 L can be boiled down to 330 mL 30%+ hydrogen peroxide. If I'd do it, I'd use vacuum to lower the b.p. and avoid most of the degradation.

TBH, sodium percarbonate doesn’t degrade with time, and is very handy to store. The only drawback is that dissolving it into any acid it releases sodium carbonate with tends to neutralise the acid, but with concentrated sulphuric acid it’s a minor issue.

Besides, one of my next experiments will be an hydroboration of styrene to get phenethyl alcohol (something I've been wanting to synthesise for a long time), and the final quenching step, which involves oxidation of the intermediate by hydrogen peroxide under basic conditions, can conveniently (at least I think) be created by adding portions of sodium percarbonate.

[Edited on 23-5-2021 by Keras]

garphield - 30-5-2021 at 21:38

Cyanamide will react with hydrogen sulfide to generate thiourea (which will thermally decompose to generate CS2 among other things, source is attached pdf). What would the products of the reaction between molten sulfur and calcium cyanamide be? Molten sulfur isn't great but it is both easier to produce and less unpleasant than either H2S or SxCl2.

Attachment: thiourea_thermal_decomposition.pdf (169kB)
This file has been downloaded 281 times

rockyit98 - 31-5-2021 at 07:29

You can make your own 30% H2O2 using the 3% from the store. […] Trust me

Oh, I know! We can even get 12% here, so 1 L can be boiled down to 330 mL 30%+ hydrogen peroxide. If I'd do it, I'd use vacuum to lower the b.p. and avoid most of the degradation.

No need to boil I think, what you need to do is freeze. there are videos on the subject on YT.

Alkoholvergiftung - 31-5-2021 at 23:23

Or without loses.You can dry it. one big baker or vakuuexicator filled with h2so4 or other high hygroscopic material and an smaller baker with h202.Stored in an dark place. Some guy had an youtube video too he needs little bit over 70days to reach 75%.
https://www.youtube.com/watch?v=0vcbZQHcPWU&t=151s
only in german

[Edited on 1-6-2021 by Alkoholvergiftung]

[Edited on 1-6-2021 by Alkoholvergiftung]

Keras - 1-6-2021 at 04:36

Quote: Originally posted by Alkoholvergiftung

Some guy had an youtube video too he needs little bit over 70days to reach 75%.
https://www.youtube.com/watch?v=0vcbZQHcPWU&t=151s
only in german

Lol. Genau, das stimmt. :p

But you have to be very patient, or plan it quite in advance!
(Besides, with sulphuric acid now in short supply, that method is going to lose its appeal, I suppose).

Opylation - 4-6-2021 at 10:49

Quote: Originally posted by garphield

Cyanamide will react with hydrogen sulfide to generate thiourea (which will thermally decompose to generate CS2 among other things, source is attached pdf). What would the products of the reaction between molten sulfur and calcium cyanamide be? Molten sulfur isn't great but it is both easier to produce and less unpleasant than either H2S or SxCl2.

That paper is very interesting. The only issue there is how do you get the CS2 to not react with ammonia? Maybe add an alkoxide into the mix to form a xanthate? Or maybe alkylate the Thiourea so that the reaction forms a dithiocarbamic acid so after isolation you can decompose it a lower temp and hopefully distill CS2 out?

Or maybe it’s even easier than that. Feed the gaseous products into hydrochloric acid and then use a sep funnel to obtain carbon disulfide

[Edited on 4-6-2021 by Opylation]

Keras - 4-6-2021 at 22:53

That paper is very interesting. The only issue there is how do you get the CS2 to not react with ammonia?
Or maybe it’s even easier than that. Feed the gaseous products into hydrochloric acid and then use a sep funnel to obtain carbon disulphide

Would CS₂ and NH₃ react?
As I explained, I heated thiourea, and passed the gaseous products produced first into a Drechsel bottle full of mineral oil, then another full of hydrochloric acid. Got what I think is NH₄Cl, though I am not sure, because the crystals I collected after evaporating the HCl are white, but more “fluffy” than I imagined. There was no trace of CS₂ whatsoever, albeit the gases were led to the Drechsel bottles though a cold trap.

Opylation - 5-6-2021 at 00:30

Ammonia and Carbon Disulfide react to form ammonium thiocyanate and hydrogen sulfide. However, what I didn't notice until I was getting this screen grab is that ammonium thiocyanate is in equilibrium with thiourea, which means that any ammonium thiocyanate can just be converted back to the starting material. Even further down, it states that my concerns aren't even warranted as ammonium thiocyanate decomposes to CS2, NH3, and H2S

Also, would the carbon disulfide no dissolve in the mineral oil? They're both non-polar molecules. I would omit the mineral oil and feed it straight into HCl acid to remove the ammonia

ammonia and carbon disulfide.png - 134kB

[Edited on 5-6-2021 by Opylation]

Keras - 5-6-2021 at 22:50

Also, would the carbon disulfide no dissolve in the mineral oil? They're both non-polar molecules. I would omit the mineral oil and feed it straight into HCl acid to remove the ammonia

Ha. Gosh. I didn't even think about that. I was worried CS₂ could somehow react with HCl or water. As you say, I should probably first scrub the ammonia using HCl, then led the exhaust into something that dissolves CS₂. However, I fail to understand why the carbon disulphide did not condense into the receiving flask, despite it passing through a Liebig condenser and that flask, both being cooled by ice cold water.

Opylation - 6-6-2021 at 03:15

I’m not quite sure. I haven’t performed this reaction but have logged it for when I have time to try it. It might require heating in a stream of HCl to lock up the ammonia. The decomposition temperature for ammonium thiocyanate is 200C which is well below the decomposition temperature of ammonium chloride. A short path condenser may alternatively work or even a takeoff adapter from the reaction flask to a gas scrubber filled with HCl acid. It’ll probably need some trial and error unless you can find a good paper describing the exact result you’re looking for.

Also carbon disulfide does hydrolyze but not readily and should be able to handle water for brief periods

[Edited on 7-6-2021 by Opylation]

Keras - 6-6-2021 at 22:25

Unfortunately, I had no other details about the reaction but the fact that, around 220 °C thiourea under argon atmosphere decomposes into ammonia and carbon disulphide. A reaction certainly takes place, and, as indicated by the pictures I took, it also generates elemental sulphur (albeit in very slight proportion) which deposits on the cool walls of the glassware, threatening to block it, and is quite a pain to remove afterwards (needs hot toluene washing).

I’m wondering if that sulphur deposit was not caused by traces of water in the thiourea. If you try that reaction, I'd advise you starting from dried thiourea, i.e. by having it baked in an oven for an hour or so.

garphield - 7-6-2021 at 18:38

The products of the reaction between calcium cyanamide and sulfur might be higher in sulfur and therefore yield more carbon disulfide. Additionally, there would be no hydrogen so you wouldn't have to worry about ammonia, and there would be no need to work with H2S gas. Could someone with more knowledge about chemistry than me predict what the products of that reaction would be?

Junk_Enginerd - 17-6-2021 at 08:47

I made CS2 a while ago in a way I thought was pretty convenient.

I used the fact that charcoal is a pretty decent insulator to my advantage. Filled a 1 liter round bottom flask with a mix of activated charcoal and sulfur. Activated charcoal because it has a convenient and consistent particle size, is generally quite clean and not dusty, and I would imagine a little less hydrogen overall than most charcoal to minimize unwanted H2S production.

Then I simply lowered an electric heating element(100-200 W ish) into the flask and positioned it to be as far from the flask walls as possible(center) and surrounded by the charcoal.

The charcoal insulates very well and protects the flask from getting excessively hot, and allows maintaining the high temperature in the middle quite easily and with relatively low power.

Ideally the power or insulation would be regulated to condense but not freeze the sulfur on the inner walls, allowing it to pool at the bottom of the flask and circulate through the hot zone without getting into the condenser meant for the CS2.

It worked for small amounts of CS2, but there's room for improvements. The heater element is very much a consumable like this. I don't think any metal can survive being red hot in a sulfur atmosphere...
Dipping the element in glass frit to form a viscous glass layer on it helped and could probably be refined further. A halogen bulb as a heat source might work too.

I was hoping the charcoal would be conductive enough to simply use electricity to heat the charcoal resistively but based on my attempts this would require a voltage on the order of 500-1000 V to work reliably, which makes it less convenient and more dangerous...

Another interesting way that works great in theory is to use carbon fibre or graphite rods as the heater element. I tried with carbon fibre but I had problems figuring out a good way for a reliable electrical connection; it kept burning off near the connection or coming undone. An upside with this method is it should keep H2S production very low since afaik there's little to no hydrogen present in graphite.

[Edited on 17-6-2021 by Junk_Enginerd]

clearly_not_atara - 17-6-2021 at 15:43

Quote:

I don't think any metal can survive being red hot in a sulfur atmosphere...

What if it's already yellow?

Fery - 24-4-2022 at 19:26