Sciencemadness Discussion Board - CS2 a Different Way

CS2 a Different Way

Sauron - 21-5-2009 at 11:48

According to Mellor. passing acetylene into molten sulfur produces CS2. I have yet to try to dig out the citation but this is a ,och lower temp. reaction than the usual coke and S vapor at median red heat.

A stab at the stoichio,etry:

(CH)2 + 5 S -> 2 CS2 + H2S

The high temp process also always produces H2S byproduct and Mellor discusses absorbing the gas with slaked lime or ferric hydroxide.

This may conceivably be basis for a lab prep.

panziandi - 21-5-2009 at 12:24

Well I know sulphur vapour and methane react to form CS2 and H2S, as do other alkanes, this is synonymous with combustion of hydrocarbons in oxygen. It is no secret that acetylene is more reactive than saturated alkanes so no doubt this reaction works. Nice work Sauron be good to dig up these references and find out some details.

garage chemist - 21-5-2009 at 17:04

I remember reading somewhere that acetylene and molten sulfur give thiophene. I don't know where this was from, but if one chooses to replicate this experiment, great attention should be paid to the composition of the product, and if necessary, isolation of CS2 from other compounds.

S.C. Wack - 21-5-2009 at 18:03

http://dx.doi.org/10.1039/JR9280002068

len1 - 21-5-2009 at 19:45

Thiophene is apparently only produced in amounts of a few percent, but is easy to separate. H2S is produced in much larger amounts, and as the carbon has nowehere else to go I expect CS2 as well. Im quite suprised by this. Will try to get a complete reference

Sauron - 21-5-2009 at 19:51

Thanks, S.C.

Here is the J.Chem.Soc. paper. The reaction does work, optimally at c.500 C and about 77% of S is converted to an oil brown liquid composed of aboou 77% CS2. Thiophene is a minor product c.5%.

Fractionation affords CS2 of reasonable purity (garlic odor suggestive of allyl sulphide contaminant) good enough for production of CCl4 - a compound Mellor avoids mentioning.

Attachment: jr9280002068.pdf (193kB)
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garage chemist - 21-5-2009 at 20:07

This is a method for CS2 production which only requires standard lab equipment (gas generator, distillation setup, H2S absorber). Excellent discovery, Sauron and S.C. Wack.

The drawback seems to be the careful fractionation necessary for a relatively pure product, which is still contaminated by an unknown impurity.
Also, this process is going to seriously foul up the glassware, in addition to the sulfur, we'll also have carbon deposits.

This has to be tried ASAP. When I am able to get around doing it, I will document it with pictures.

Sauron - 21-5-2009 at 21:37

Yes, g c, precisely why I thought it noteworthy. Most likely my equation is inaccurate, but, maybe not. The rest of products are side reactions with their own stoichiometries.

Previous threads describing attempts at scaling down the industrial process came to little as I recall.

For me, CS2 is just $$, I can buy it, but I know things are tougher in parts of EU so I sat up and took notice.

I found this in Carbon Part II chapter of Mellor Vol VI and I will post the entire 134 page chapter in References very shortly (New Books - Inorganic). Many processes are described, but this is the most bench scale friend;y. Also it is very cheap.

I would think that Norit would be handy ay snagging the trace impurities, wouldn't you reckon Likewise LC.

[Edited on 22-5-2009 by Sauron]

len1 - 21-5-2009 at 23:51

For some reason this process, despite being little known - I havent previously seen any chemical literature references to it despite encountering a lot in my lifetime - appears to be real and high yielding. Apparently all the ingredients to make CS2 at 330C are available in any hardware store, as ethylene it seems can also be used. Here is my reference to it.

http://pubs.acs.org/doi/abs/10.1021/cr60123a002

[Edited on 22-5-2009 by len1]

Sauron - 22-5-2009 at 01:20

Well, acetylene is readily available at any welding supply shop in various sizes of cylinder, and regulators for the appropriate CGA fitting are cheap new. ($50) Or buy calcium carbide, and generate it.

Is ethylene easier than that? I can probably get it from the Thai Industrial Gases people but I bet it costs more per mol than acetylene. I am puzzled about obtaining it at a hardware store. Enlighten me.

Here is the Chem Rev paper, who needs citations? I have Chem Rev onmy desktop.

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Lambda-Eyde - 22-5-2009 at 01:25

Quote: Originally posted by Sauron

Fractionation affords CS2 of reasonable purity (garlic odor suggestive of allyl sulphide contaminant)

This page lists the b.p. of C6H10S as 138-140 C @STP. That means you're using a really shitty fractionating column if you get that together with your CS2 (b.p 46 C).

...Am I the only one concerned with how violent this reaction will proceed? Generating (CH)2 alone might sometimes cause it to self-ignite if you're not proceeding carefully. However, this is easily prevented by flushing the apparatus with nitrogen prior to adding the CaC2.

Sauron - 22-5-2009 at 02:12

The thiophenol need only be present in very low concentration to impart that odor. It's obnoxious stuff.

Lambda-Eyde - 22-5-2009 at 02:50

Make up your mind, are you talking about allyl sulfide or thiophenol?

BTW, commercial acetylene is contaminated with acetone, just so you know.

Sauron - 22-5-2009 at 03:26

Make up whatM Garlic suggested allyl sulfide to me, further reading revealed the culprit to be thiophenol, so my guesswork was wrong.

If you bothered to read the articles posted you would know this already, but instead you just want to blather on.

Yes I am familiar with the way acetylene of commerce is stabilized. I do not see this as a problem. The amount of acetone is small, and I have seen nothing to indicate that acetone will react with S under these conditions. Do you have references to the contrary?

Lambda-Eyde - 22-5-2009 at 03:42

Quote: Originally posted by Sauron

Garlic suggested allyl sulfide to me, further reading revealed the culprit to be thiophenol, so my guesswork was wrong.

If you bothered to read the articles posted you would know this already, but instead you just want to blather on.

I'm sorry, my fault then.

Quote: Originally posted by Sauron

Yes I am familiar with the way acetylene of commerce is stabilized. I do not see this as a problem. The amount of acetone is small, and I have seen nothing to indicate that acetone will react with S under these conditions. Do you have references to the contrary?

No, but I thought I should bring it to your attention since noone else had mentioned it in this thread.
Typically a commercial acetylene cylinder is filled to 50% of its capacity with acetone, however I don't know how much of it escapes when the pressure is relieved. I'm sure it's negligble, like you say, and that it doesn't interfere anyways.

len1 - 22-5-2009 at 04:02

Thanks - it takes me a bit longer to get JCS articles since I dont have a DVD. Ethylene can be produced from alcohol and H2SO4, both available at the hardware store - though prob using bottled acetylene is easier.

The trouble there has already been mentioned - acetylene liquifies at about 4.4 MPa at room temperature - its actually thermodynamically unstable and can explode if compressed much above 2MPa. Acetone dissolves it reducing its vapour pressure, which is about 3MPa and constant in commercial cylinders. Its not liquified like CO2.

How much acetone escapes - the answer is clear, the cylinder is not full of acetone when you return it - so the answer is almost all. How much does it affect the reaction - who knows. Its best removed with a cold trap.

[Edited on 22-5-2009 by len1]

panziandi - 22-5-2009 at 04:18

I suppose you could distill the crude CS2 then double distill it in a fresh set up, however the usual method (and I'm sure people here have heard this before) is to shake with a quantity of mercury which removes most other organic sulphur and leaves the carbon disulphide with a more, pleasant, odour... not that I suggest anyone tries sniffing carbon disulphide! I remember being told once it causes bipolar disorder.

watson.fawkes - 22-5-2009 at 06:10

FYI. Here's an extract of the entry for CS2 in Purification of Laboratory Chemicals:

Quote:

Shaken for 3h with three portions of KMnO4 soln (5g/L), twice for 6h with mercury (to remove sulfide impurities) until no further darkening of the interface occurred, and finally with a soln of HgSO4 (2.5g/L) or cold, satd HgCl2. Dried with CaCl2, MgSO4, or CaH2 (with further drying by refluxing with P2O5), followed by fractional distn in diffuse light. Alkali metals cannot be used as drying agents.

BromicAcid - 22-5-2009 at 09:04

As mentioned before, not only acetylene, but methane, ethane, propane, etc., can all be used to afford this conversion. If there are worries about acetone impuries in acetylene then maybe methane is also worth investigating, it is even more readily avalible for some although honestly I don't know what impuries there are in the stuff coming out of my wall.

Industiral and Engineering Chemistry "Carbon disulfide production: Effect of catalysts on the reaction of methane with sulfur" Feb, 1944 Vol 36, No. 2 Pgs 182-184

One of the remaining references that I have on hand, the following reaction conditions are described:

Silica gel catalyst, temp = 550C, Molar ratio CH4:S2 = 0.5, Space Velocity 825 the conversion is 42.1%. At T=600C the conversion increases to 69.6%.

Quote:

Analysis of the data shows that hte reaction of sulfur with methane under the conditions described proceeds without appreciable side reaction according to the equation: CH4 + 4S ----> CS2 + 2H2S

Just another route for those without ready access to acetylene. Temps are slightly higher however.

len1 - 22-5-2009 at 10:06

Well yes, as you point out this reaction fits into the continuum spectrum of reactions of few carbon atom unbranched hydrocarbons with elemental sulphur.

Yet if this was advertised as a CH4 - S reaction I wouldnt have bothered - its mentioned in almost every text on CS2 I have seen. Trouble is a gas phase reaction, at 600C plus, above the melting point of sulphur, requiring a solid phase/catalyst for interaction, producing a highly flamable gas, is no mean feat even for industrial chemistry, let alone an average small lab.

So the version of the reaction of C2H2 with S at 500C, poses no special interest.

But bubbling C2H2 in LIQUID S at 325C, with no catalyst, IS within the reach of most labs, and that it produces CS2 at 30% yield based on S is mentioned almost nowhere - that is what is most surprising.

BromicAcid - 22-5-2009 at 10:22

Agreed, using liquid sulfur is by far preferred over using gaseous sulfur. It is what separates the industrial process from something readily doable at home. Although running at 325C wouldn't be a walk in the park (plugging of gas inlet tube, subliming sulfur everywhere, possibility of plugging of the apparatus from subliming sulfur, etc.) it is still much preferable.

For the most part I was just pointing out that if you were shooting for the optimal conditions of the reaction anyway (Mentioned by Sauron to be 500C) that there were other gasses that could be utilized.

watson.fawkes - 22-5-2009 at 10:40

Here's another extract from Purification of Laboratory Chemicals:

Quote:

Acetone vapour can be removed from acetylene by passage through H2O, then concd H2SO4

This I find clever because it uses water to scavenge acetone, rather than the typical reverse situation, which is to use acetone to remove water.

len1 - 22-5-2009 at 10:40

Another interesting aspect of this is that the reaction produces voluminous amounts of H2S - a gas poisnous at the scale of HCN. If it was HCN that was the byrpoduct, most people - including myself - would be put off, and again Id forget this reaction. Its funny that H2S doesnt have the same effect. Mainly because its much more common place, and can be detected in much lower concentration - I can not smell HCN at all, and so am freightened of it.

H2S we produced as kids. I remember the experimental chemistry book for children where you burn Fe and S to make FeS, then dissolve it in HCl to make H2S,which you collect over water. Not a word about it being a deadly poison - dont think many people knew- it was just called rotten eggs gas. I think an additional factor is at work here -the parity in the poisnous nature of the two acids does not extend to their salts.

[Edited on 22-5-2009 by len1]

Smells

watson.fawkes - 22-5-2009 at 15:48

Quote: Originally posted by Lambda-Eyde

allyl sulfide or thiophenol

It's almost certainly not thiophenol, since the J.Chem.Soc. paper referenced above mentions that benzene was entirely absent from the products. It does, however, have a smelly thiol group.

Also unlikely is allyl sulfide, named more structurally diallyl thioether, since that contains a pair of 3-carbon groups, and few intermediates and rearrangements would be necessary to make that.

It's likely thiophene, C4H4S, one of the products mentioned by name and formula in the same paper. The Wikipedia page mentions thiophene as having "a mildly pleasant odor reminiscent of benzene". On the other hand the thiophene entry at The Good Scents Company describes the odor as "garlic". I'll believe a scent database over Wikipedia, so this is my guess as to the odor reported in the paper. I'd imagine it forms an azeotrope with CS2. Even an azeotrope at a ratio of 99.99/0.01 would still smell plenty.

The other named product, "thiophten", is an alternate spelling of thiophthene, C6H4S2, a double-ringed aromatic, both rings heterocyclic with sulfur. It's possibly also a culprit. Its two aromatic sulfur atoms are in the same position as that of thiophene, so I'd guess it's got an odor in the same family.

There are two simple thiol modifications of thiophene. One is 2-thiophene thiol, but the linked page describes it as "burnt caramel roasted coffee". The other is 3-thienyl mercaptan, described as "cooked meat". That route doesn't look promising.

It's conceivable that it's the relatively-low-molecular-weight 1,3-butane dithiol, listed as a primary garlic scent. It wouldn't have to be produced at very high concentrations, but it might not be produced at all, given that the other two side products are both aromatics and more stable than a modified alkane.

Please ignore this.

watson.fawkes - 22-5-2009 at 17:38

[Edit:] Oops. Wrong. WRONG. Please disregard.

Quote: Originally posted by Sauron

Here is the J.Chem.Soc. paper. The reaction does work, optimally at c.500 C

Digging into this a little, I've figured out that the temperatures in this 1928 paper must be in Fahrenheit, not Celsius. The boiling point of sulfur is at 444.6 C, so it's doubtful there would be liquid in the reaction tube at that point.

The authors report the best yield at 500 F = 260 C. This is something of a critical point in an extended phase diagram of sulfur (one that takes viscosity into effect) (Mellor, Sulfur, p.44). It's the point where the polymers in the plastic phase of sulfur are all breaking up. No coincidence there, it seems. There's less thermal energy needed to break S-S bonds (for adequate reactivity) and not so much energy as to make ever-more aromatic byproducts.

Very curiously, the lowest reported temperature where the reaction proceeded with any speed was ~ 325 F = 163 C. This is about the temperature where polymerization into plastic sulfur starts. Presumably this is because the S8 rings are being broken by heat. I would conclude that a C2H2 + S8 collision isn't very reactive, and that the first intermediate reaction requires a terminal (in its chain) sulfur atom (or a free one) to get started.

Well, the good news is that the optimal temperature at 260 C is less even that the minimal wrongly-interpreted temperature of 325 C.

[Edited on 23-5-2009 by watson.fawkes]

len1 - 22-5-2009 at 17:45

Quote:

Nice wish, but unfortunately not the case. The temperatures is the article are in Celcius, as the b.p. for CS2 is quoted at 46.5

[Edited on 23-5-2009 by len1]

watson.fawkes - 22-5-2009 at 17:48

Quote: Originally posted by len1

Another interesting aspect of this is that the reaction produces voluminous amounts of H2S - a gas poisnous at the scale of HCN. If it was HCN that was the byrpoduct, most people - including myself - would be put off, and again Id forget this reaction.

I think the big difference is that you can flare off H2S, but not so with HCN. I'd rather have an SO2 problem than an H2S problem. If I were doing this at home, I'd incorporate a little propane bottle and burner orifice just to make sure the flame never went out.

Perhaps the next thing to talk about is a home-brew Claus process plant or a diethanolamine scrubber.

watson.fawkes - 22-5-2009 at 17:52

Quote: Originally posted by len1

Nice wish, but unfortunately not the case. The temperatures is the article are in Celcius, as the b.p. for CS2 is quoted at 46.5

Are those 500 and 650 degree reactions they report being done in the gas phase, then? That's not obvious to me, even on third reading. They seemed to have changed apparatus, though; perhaps that's the sign.

Sauron - 22-5-2009 at 20:37

Acetylene produces half the H2S that methane would, and twice the CS2 (in theory) That's mol for mol. If you equalize the CS2 output by using two mols CH4 you make four times as much H2S. Clearly acetylene is the feedstock of choice for amatuers, on this basis alone. Less H2S is better, safer, less troublesome.

At least two simple methods have been advanced for scrubbing acetone from tank acetylene for lab use: cold trapping and successive washings with water and conc H2SO4. Both easy and both cheap. And it has not even been established that this procedure is necessary. Though I would do it on principle.

Cold trapping will lose some acetylene as solute but who cares?

It is too bad that watson-fawkes was mistajeb re temp scale, as 325 F would have been a nice initiation temp atd 500 F a nice optimum. Frankly I would personally stay at 450 C just so I can work in borosilicate and not have to mess with bycor or quartz. That ought to be close enough for gubbimrnt work, as they say.

The Mellor chapter XXXIX from which this reaction emerged is up in References already. I extracted the section on carbon sulfides but it is too large to post directly here, I will pare it down to the essential pages and do so. He does have some things to say about trapping H2S, and I found all the other prep methods interesting, such as cultivation of an East Javan fungus that on dry distillation releases CS2. All you biotech fans, take note!

For anyone wanting the longer extract:

Pages from CarbonPartII.pdf
http://www.4shared.com/dir/2245331/5a78115f/sharing.html

The forum software barfed twice on the 1.55 Mb shorter extract. So I guess I will have to split it into its two sections, Preparation and Physical Properties, and post them separately.

[Edited on 23-5-2009 by Sauron]

Attachment: CS2 Preparation.pdf (503kB)
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Sauron - 22-5-2009 at 22:49

And Physical Preoperties

Attachment: CS22 Physical Props.pdf (371kB)
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len1 - 22-5-2009 at 23:45

So it looks like we need G. Capelle, Bull. Soc. Chim., (4), 3. 151, 1908.

I have severe doubts that Taylor's electric furnace - the only one illustrated for CS2 preparation actually works well. S leaving the superheating region will not have sufficient heat to raise the temperature of the carbon to the needed 900C, unless it is absolutely flooded with sulphur vapour. The latter condensing in the coke chimney will then choke it. Seems very wasteful

Sauron - 23-5-2009 at 03:05

I am not thirsting for Capelle's paper, but if I were I would be looking to Wiley where I believe it is listed in either the European Journal of Organic Chemistry, or its Inorganic counterpart, as having been incorporated into said journal, along with many others. That being said, I have never found a Wiley biblio page for it.

Not even Gallica, the cyber arm of the French Biblioteque Nationale, has it.

Why do you reckon it is vital> True, it is the sole citation Mellor gave this reaction, but as Mellor obviously regarded this reaction of little merit, that means nothing. we mihght as well chase the paper on thioformamide hydrochloride.

Über Thioformamid
Richard Willstätter, Theodor Wirth
Ber., 42 (1909) (p 1908-1922)
DOI: 10.1002/cber.19090420267

and available only from Wiley.

The Meyer and Sandmeyer paper from Ber. cited in J.Chem.Soc. along with Capelle is quite available since I downloaded alkl of that journal that was available free from BnF and hosted it on 4shared. Note that these are directed toward thiophene and that CS2 appears to have been regarded as a byproduct. Like H1S.

I would not muck around with thw Taylor firnace. If I wanted to try a high temp route I would look at Fe or Cu pyrites and carbon, particularly a carbon form free of hydrocarbons, so no H2A forms, and do it in my tube furnace. But, I see little to gain from this approach since having to deal with twice as hot, or thrice as hot CS2 vapor outweighs the hassle of dealing with taming and destruction of H2S.

Here is the Mryer and Sandmeyer note, a few paragraphis only, sans details. Useless.

Attachment: Pages from Ber_16_2.pdf (55kB)
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[Edited on 23-5-2009 by Sauron]

watson.fawkes - 23-5-2009 at 06:49

I have become interested in the observation that all these CS2 preparations are gas interactions in some way. The Taylor furnace is gas-solid. Of the preparations in the J.Chem.Soc. paper, three are gas-gas phase (one as vapor above liquid, the other two purely gas) and one is gas-liquid (bubbling C2H2 under molten sulfur). In that paper, the two higher temperature reactions proceeded at higher sulfur efficiencies, so presumably at higher reaction rates as well. There are two ways that have occurred to me to increase the gaseous availability of sulfur.

The first is to treat the process, in part, as a vacuum distillation. Removing the atmospheric back pressure will lower the boiling point of sulfur, but more importantly, at any given temperature will increase the rate of sulfur vapor flowing off the surface of the molten sulfur. This technique requires that the H2S be dealt with differently than a simple flare. The easiest way to keep the atmosphere confined, a manostat set at something over ambient atmospheric pressure, creates pressures worse than atmospheric, so that doesn't work. The boiling point of H2S is 213K and the sublimation point of CO2 is 195K, so another technique is to use a cold trap with a lower loop kept full of liquid H2S, dealing with it when it boils on the high pressure side. Since this loop acts as a barometer in the naive setup, there's some question about how you make it tall enough. Perhaps with a pump, a solenoid valves, and a couple of check valves this could be alleviated. A third method is to use a compressor. Getting a compressor not to fail where pumping H2S is something of a lubrication miracle; I don't have any particular ideas on this. If there are sulfuretted compressor oils out there, that might work. A fourth technique is to use a liquid phase scrubber, such as one with diethanolamine. These frequently use other solvents to get the H2S into solution, then neutralize it with the reaction NaCO3 + H2S --> NaHCO3 + NaHS. (Even if you burn off most of your H2S, if you need some Sodium hydrosulfide this reaction might be useful to you.)

A second way of dealing with vapor is to ask the question about the processes in the paper above, "Where did the rest of the sulfur go?" The authors don't say explicitly, but I'm assuming that it went over in vapor phase into the distillate. Temperature-based reaction rates, then, are competing with CS2 formation rates. If you can increase the dwell time of the sulfur vapor in your system, you'll deal with this issue just fine and can be happy with lower reaction rates. So the question becomes "how do I put my sulfur vapor under reflux?" If you operate your apparatus at the boiling point of sulfur, 445 C at ordinary atmosphere, you'll satisfy Sauron's desiderata, which is perfectly reasonable:

Quote: Originally posted by Sauron

Frankly I would personally stay at 450 C just so I can work in borosilicate and not have to mess with bycor or quartz. That ought to be close enough for gubbimrnt work, as they say.

So you need a cold finger operating at 400 C or so. At that temperature, "cold" is perhaps not the most apropos word, but oh well. The point is that you want a cold finger that discriminates between sulfur and the reaction products. The harder question is how to build one. Luckily, the cold finger can operate anywhere in the range 300 - 400 C, so close temperature regulation isn't necessary. The question then become what your working material in the cold finger is. Any pure liquid phase is going to be a real bother to deal with. So don't do that.

It seems that a flash boiler with water would do the trick. As an input, you'd have a metered drip tip allowing water into the end of the finger at some appropriate rate. On the output you have some steam to vent. My view is that if you can deal with H2S, you can deal with a little steam. As for materials, it would seem that this could be built with standard plumbing parts, albeit with hard soldering required at the boiler end (although a crimp might do just as well). Packing the bottom few cm of the pipe with stainless steel shot or ball bearings is an easy way of increasing the surface area at the bottom. You'd need to use distilled water as a medium, because it would be easy to clog the evaporation surface. But you've got very hot steam coming out, and you could use this to run a water still. You'd need a thermocouple at the bottom of the finger, but since the temperature range is so large, manual regulation should be plenty adequate.

Lastly, there's the possibility of raising the temperature in the reaction zone without raising the temperature of the whole apparatus. As Sauron says of the Taylor furnace and other high temperature techniques:

Quote: Originally posted by Sauron

But, I see little to gain from this approach since having to deal with twice as hot, or thrice as hot CS2 vapor outweighs the hassle of dealing with taming and destruction of H2S.

Insofar as building an apparatus that is all hot, I'm with him at the personal scale. If you're going to build something like that, you'll also be making far more CS2 than you really need personally. So if possible, one should heat up only a high-throughput reaction zone. I should point out that this class of techniques is more to increase the output rate. If the output rate is already fast enough, then these won't be needed.

The simplest concept for doing this is just to put a heater element into the chamber. I figure a tungsten rod, such as TIG electrode, would do. There are some engineering questions about how to get the heater element mounted and conductors into the reaction zone. You'd likely need a cooling jacket in the glass around the heater, but this can be run with liquid water. You might be able to do the whole thing inside a Liebig condenser with a largish inner bore. Tricky, because you don't want to condense the CS2, although if you're using a cold finger for the sulfur you don't need to care. You can put the cold finger on the other arm of a Claisen adapter and take off your product there. If doing this, make sure the cold finger is above the heated reaction zone so that the reflux height of the sulfur is adequate.

The second technique uses a susceptor with infrared heating. A standard electric heater is set up and focused upon the susceptor with a lens. Most Fresnel lenses are plastic, but an air cooling system with a pair of glass panels, a frame, and a fan should keep it within operating temperature. Since this is probably a lower temperature process than a heater, air cooling (perhaps forced air, would be adequate). If you were totally clever, a water jacket could be used as a lens in the optical train. The third technique uses a susceptor and an inductive heater. The advantage here is that you don't need internal electrical connections. The disadvantage is that you need a power supply, an induction coil, and its cooling system. Cooling for the coils also will cool the glass in the zone, so there's no need for a separate system there.

On the other hand, if you have an inductive heater, you can use solid carbon rod as the susceptor and reactant all in one. If you make a feed tube, you can simply gravity-feed rod in from the top and have something of a continuous process. Engineering the bottom of the feed tube will tricky and likely involves quartz. If you're going to the expense of obtaining inductive heating, though, you get the advantage of eliminating an H2S disposal problem.

Sauron - 23-5-2009 at 07:48

How about a Vycor immersible heating rod (Corning?

I found a cyber-stash of Bulletin de la Société Chimique de France covering about 50 years, with gaps, but alas 1908 is nit there (1907 is). When I have time I will d/l these and host them on 4shared. Presently they are all jumbled up.

Wiley seems no help.

The JCS authors state quite clearly that 74% of the S charge at 500 C ends up in the liquid condensate and they explicitly state the % composition of the major and minor products so if you calculate the mass balance you will see that is where they get 74%. Accounting for the other 26% of S as H2S is surmise but reasonably so. My guesswork equation is

(CH)2 + 5 S -> 2 CS2 + H2S

Since the acetylene flame is sootym a competing reaction is

(CH)2 + S -> 2 C + H2S

Eq 1 would dictate 80% S conversion to condensate, observed is 6% lower. It is reasonable to attribute that to Eq 2 which must have about 40% of the reaction rate of Eq 1 because if rates were equal, 16% of S woukd be consumed that way and the S accounted for in condensate reduced more than observed.

Or in other words, just assubg the two equations would only reflect 66% of S in condencate (which for simplicity is expressed as CS2 but is really CS2 + thiophene + thiophten) and that is inconsistent with observed result so Eq 2 must be slower.

A better kineticist doubtless can express this more elegantly.

watson.fawkes - 23-5-2009 at 08:54

Quote: Originally posted by Sauron

How about a Vycor immersible heating rod (Corning?

Yeah, that'd work. I had mentioned exposed tungsten because the modal reader here seems to have more time than money.

Quote:

The JCS authors state quite clearly that 74% of the S charge at 500 C ends up in the liquid condensate and they explicitly state the % composition of the major and minor products so if you calculate the mass balance you will see that is where they get 74%. Accounting for the other 26% of S as H2S is surmise but reasonably so.

It's the composition of the other 26% that's in doubt. When the authors report the preliminary experiment, I got a distinct impression that the residue of the first distillation was not 100% thiophthene, since the implication is that the residue (of the first) was not identical to the distillate of the second, and therefore there was some additional, even higher-boiling fraction remaining. Assuming that's some sulfur and higher-MW aromatics (likely some triple rings in there), that would be a candidate for a component of the 26%. I do admit that while my interpretation differs, I don't think there's enough information to make a definitive call between mine and the "it's all in H2S" one. Future experiment will tell. I do suspect the real answer is some mixture of these hypotheses.

At most, it's just a few percent of sulfur in question, not enough to forestall initial trials. For me, the question hinges on the utility of making a cold finger. It seems that the original authors would have been using an excess of acetylene, which is fine for a research experiment, since they want to track the mass balance of the sulfur. For low end production, however, it might prove better to use an excess of sulfur, which is what a sulfur reflux would get you. If you wanted quantitative measurements in this case you'd be tracking the mass of acetylene, which you could estimate with a flow gauge. Time will tell.

The baseline 80% equation seems just fine to me.
i) C2H2 + 5 S --> 2 CS2 + H2S
The other two named products seem to have these reactions:
ii) 2 C2H2 + S --> C4H4S (thiophene production)
iii) C4H4S + C2H2 + 2 S --> C6H4S2 + H2S (thiophthene production)
The competing reaction
iv) C2H2 + S --> 2 C + H2S

Note the H2S production in equation (iii). I don't have time this morning to work out the effect of this on the mass balance, but it's definitely a part of the picture.

Sauron - 23-5-2009 at 10:09

All that looks reasonable. My focus remains making CS2 inexpensively. The identity and quantification of minor products will fall to instrumental analysis, there is nothing there that can hide from GC, MS. IR, NMR, and/or HPLC.

Eclectic - 23-5-2009 at 14:44

I'd think a plain glass or vigreaux riser would work fine for sulfur reflux
You might even need to put some heat on it until it gets up to temp.

[Edited on 5-23-2009 by Eclectic]

Sauron - 23-5-2009 at 17:15

Perfect job for a heating tape. Wrap the Vigreaux and use a simple controller to hold it at 350 C +/- 50.

Sauron - 24-5-2009 at 01:35

An uneconomical route to CS2 is the thermolysis of thioformamide hydrochloride in absence of air/O2.

See paper cited upthread.

Thioformamide is prepared by reaction of formamide with P2S5 and it is the cost of the pentasulfide that is the deal killer.

The attached US patent describes an improvement over the Wilstatter and Wirth method of Ber.41. The improvement is to employ THF rather than Et2O as solvent. Still the yield hovers at 50-60% and I do not yet know how efficient the conversion to CS2 is, this would be expensive CS2 indeed.

Attachment: US2682558.pdf (176kB)
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Sauron - 24-5-2009 at 02:18

It turns out there are two other routes to thioformamide.

The first has pretty much same drawbacks as the classical one as it proceeds through dithioformic acid and NH3: J.Chem.Soc. 1937, 361

The secons is the reaction of HCN and H2S in nonaqueous protic or aprotic colvents and catalyzed by ammonia or trialkylamine. Angew.Chem.Intl.Ed. 8, 278 (1969) see attachment.

Sound like fun?

Attachment: Pages from 8.pdf (119kB)
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[Edited on 24-5-2009 by Sauron]

Sauron - 24-5-2009 at 03:09

I obtained the J.Chem.Soc. paper easily enough (attached FWIIW)

Potassium dithioformate is prepared from chloroform and potassium sulfide. Dithioformic acid, which is unstable, and needs immediate use, is liberated from the salt .

The bad news is that details of both steps are to be found in Levi, Atti R. Accad. Lincei, 1923, 82, I, 569. Looks unlikely to be available, save maybe as a CA abstract.

Attachment: jr9370000361.pdf (323kB)
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watson.fawkes - 24-5-2009 at 06:19

Quote: Originally posted by Eclectic

I'd think a plain glass or vigreaux riser would work fine for sulfur reflux
You might even need to put some heat on it until it gets up to temp.

With this consideration we move from chemical engineering into mechanical engineering. Such a riser would work physically, no doubt. It only might work in practice. The risk is glass breakage; the cause, thermal gradients.

The consequences of losing containment of this operation are pretty large. Starting with the easiest, you have a sulfur vapor plume to deal with. Some of this plume will condense in air, creating yellow smoke. Some will oxidize to SO2, and not in small quantities. Even assuming this is done in a fume hood, you'll likely need supplemental breathing gear, not to watch the pretty smoke, but when you need to reach in a start manipulating equipment. And then there's the explosion risk of hot acetylene suddenly introduced to oxygen. Not to mention the hot CS2 will immediately start burning, as its autoignition temperature is quite low, 90 C.

The practicality here is that *all* the glassware in the hot zone needs to be insulated so that random breezes don't create enough of a strain in the glassware to cause random breakage. It's a real risk, as the instructions in every single glassblowing book I've read talk about the dangers of drafts in the glass shop. The problem is not steady-state strain, but the differential strain that happens when cold air hits one side of a piece of glassware and not the other.

Insulation on a riser tube will generally interfere with its function as a condenser. The job of a condenser is to draw away heat, that of insulation to contain it. You could make the riser tube longer to compensate, of course, but there are architectural limits to that idea. This might work, but at the very least is going to require a lot of experimentation and/or some good engineering calculations.

Left open is the possibility of using air as a working fluid, as the coolant in a condenser. If this is tried, it would need apparatus to keep the air flow smooth and the temperature gradient steady. A shroud around the riser, forming a plenum for the air flow, is the first step. A fan to provide forced air is the second. And I'd add a mixing box between the fan and the plenum to equalize temperature swings in the inlet air to the fan. If you set the fan rate to overcool the riser tube, you could then use Sauron's suggestion of a heating tape to set a fixed point temperature on the glass. This might work. If tried, it would be wise to test it with sulfur reflux only for a while, at least as long as the CS2 reaction run is expected to take.

Another possibility for a glassware system is to use a high boiling point solvent as a coolant in an ordinary condenser. You'd need a closed loop system. If you operated at the boiling point of the coolant, you wouldn't need a pump and could just rely on the circulation afforded by a boiler-condenser system. As serious issue with this idea, though, is finding a solvent with a high-enough boiling point. You need one that's above 260 C, so that you don't condense sulfur vapor into its plastic/high-viscosity state, which would pretty quickly clog the apparatus. This leaves out ethylene glycol, propylene glycol, and DMSO, for example. The mechanical principle here is that you're splitting one temperature gradient across glass into two, crossing a material performance boundary.

Even after I've said all this, I'm still most comfortable with using a metal cold finger. If I had to use glass, I would completely insulate everything. If I were doing it myself, I'd construct a special-purpose metal retort (out of an alloy resistant to sulfide cracking by H2S) to avoid glass risk entirely on the hot side.

Sauron - 24-5-2009 at 10:51

Dithioformic acis is also known as thiolthionic acid and has romula CH2S2. It is not to be be confused with dithiocatbonic acid. It is prepared from chloroform and K2S, my best guess while awaiting details is in a manner analogous to orthoformate ester preps. I would be surprised if CHCl3 will react directly with K2S but first prepare the sodium derivative of chloroform and you have a ball game. The potassium dithioformate salt is stable, the acid is not, it polymerizes steadily.

Liberating the dithioformic acid as needed and reacting it with aq. NH4OH gives a fair yield of thioformamide.

It is unlikely that this method has any real advantage over treating formamide with P2S5 in THF which seems a lot less work.

Yet ANOTHER possibility is to purchase potassium ethyl xanthate, liberate free xanthic acid which readily falls apart to CS2 and EtOH. Xanthates are cheap.

len1 - 24-5-2009 at 14:34

For the home chemist I think it would be most convenient to carry out the reaction between HCN and pressurised H2S using phosgene rather than benzene as solvent.

garage chemist - 24-5-2009 at 16:12

This sunday, I tried out the sulfur-acetylene reaction on a relatively small scale, to find out whether it actually makes CS2, and how the reaction goes.
I made pictures, but I won't post them now since the apparatus I used had a dangerous flaw that was impossible to predict and forced me to prematurely shut down the experiment. I will repeat the experiment soon and correct the apparatus problem.
It was an important learning experience, and I will now closely describe what I did, and what happened.

The apparatus consisted of the following parts, arranged from left to right in this order:
- an acetylene gas generator consisting of a round-bottom flask with pressure-equalized dropping funnel charged with 30g calcium carbide and excess water
- a safety washing bottle arranged in reverse direction
- a washing bottle with gas distribution frit charged with ca. 40ml conc. H2SO4 to dry the acetylene and free it from phosphine
- reaction vessel: a 250ml two-neck round bottom flask with gas inlet tube reaching to the bottom of the flask, charged with 200g sulfur which was being heated by a bunsen burner
- a distillation setup with liebig condenser attached to the RBF, through which cold water was circulating
- a receiver which was being cooled in ice water
- a safety washing bottle arranged in reverse
- a washing bottle with gas distribution frit charged with 100ml of 10% NaOH solution, intended as a H2S absorber.

The receiver and last two washing bottles were located under the fume hood.

I then did the following:
The sulfur was melted with the bunsen burner, and the air in the apparatus displaced by propane to eliminate the risk of an acetylene-air explosion. The propane was being introduced at the top of the water-filled dropping funnel.
It was bubbling through the first washing bottle, the molten sulfur and the last washing bottle, proving that the apparatus was airtight.
The propane gassing was continued until propane was coming out at the end of the apparatus and burned with a yellow sooty flame after ignition, proving that there was no more air in the apparatus.

Now the propane flow was stopped, and the temperature of the sulfur raised beyond the highly viscous phase until it boiled gently, but was not distilling over.

I started adding water to the carbide, and the acetylene flow began.

After a short time, as the acetylene arrived in the reaction flask, dense brown smoke was being generated, and slowly, a dark brown distillate collected in the receiver. I was excited to see this, as this is what the article said would happen.
The smoke continued flowing through the two washing bottles.

The apparatus ran for a few more minutes, and then I saw with shock that suddenly the sulfur in the reactor was flowing back into the washing bottle with the conc. H2SO4.
I immediately extinguished the bunsen burner, opened the washing bottles to let out any pressure, and stopped the acetylene production.
The apparatus was left to cool, and the receiver, which contained ca. 5ml of product, was removed and stoppered.

Upon inspection of the apparatus, I found that the last washing bottle, the one with 10% NaOH, was completely blocked. It had generated backpressure, which pushed the sulfur back into the H2SO4 washing bottle.
The smoke had formed a brown deposit in the gas distribution frit and clogged it.

So, what I've now learned is: don't use a frit washing bottle for the H2S absorber! The smoke (most likely sulfur particles from vaporized sulfur) will block it. Next time I will use a normal washing bottle for the H2S absorber.

The 5ml of raw product were then subjected to a simple distillation in a small still. Ca. 3ml of distillate collected at a gas phase temperature of 46-50°C!
The distillate (smelling strongly of H2S) burned with the bright blue flame characteristic of CS2, depositing sulfur on cold glass surfaces in the flame.

This result is very encouraging. I think it is quite safe to say that I made some CS2.
The distillation residue was high-boiling and became viscous as it cooled down.

Now, what remains to be done is making more raw product with the improved apparatus and subjecting the initial distillate to fractional distillation to separate thiophene and CS2.

len1 - 24-5-2009 at 16:29

Very good - so it works just as promissed - from what you described there is no doubt CS2 is produced in good yield. The generation of CS2 is certainly much easier than with the 900C method we did - which consigns it to the scrap heap. The issue with this method seems to be one of purification, as the product sounds like it contains far more byproducts (S, H2S, compounds of C H and S) than with the former method.

Sauron - 24-5-2009 at 18:48

Purification is not onerous, and extent of purification required varies with application intended for the CS2. You want spectro grade?

All I want is tech or lab grade to turn inro CCl4.

Kudos to g c for being first to carry this out. I am pretty sure no one who read the literature is too surprised that it works.

len1, why the snide remark? The HCN/H2S method is clealr at best an industrial one, thioformamide is a valuable feedstock for thiazoles and the other routes to it are kludgy, that ref popped out of Merck Index, and I included it for sake of thoroughness, but never considered it as suitable. Sarcasm, and phosgenem uncalled for.

It has been obvious since I started this thread that S + C2H2 was going to be the cheapest and friendliest method.

Thioformamide is interesting but a waste of time, unless a cheap supply of P2S5 is at hand.

Decomposition of potassium ethyl xanthate (commercial) will give almost 59% by weight CS2, generates no H2S, requires only dilute mineral acid and byproducts are only ethanol and salt-water (K salt of acid used). Cost depends on what you have to pay for the xanthate.

len1 - 24-5-2009 at 20:13

No snide remark - just a joke, they are OK too sometimes. Well done to GC for checking this and to Sauron for fiding this.

CHCl3 -> CS2 ?? A Suprise

Sauron - 24-5-2009 at 21:45

All in a day's work.

I was wrong about the reaction of chloroform and K2S. The metallo derivative of chloroform is not used, This from p 38 of Volume 4 of THE ORGANIC CHEMISTRY OF BIVALENT SULFUR by Reid.

CHCl3 + 2 K2S -> HCS2K + 3 KCl

The product is potassium dithioformate

According to a recent Japanese paper, the K2S us best prepared from potassium tert-butoxide

http://sciencelinks.jp/j-east/article/200106/000020010601A02...

Convenient Synthesis of Potassium Dithioformate and Formation of Some Esters of Orthotrithioformic Acid from the Dithioformate, Sodium Hydrosulfide, and Each Alkyl Bromide.Accession number;01A0261852
Title;Convenient Synthesis of Potassium Dithioformate and Formation of Some Esters of Orthotrithioformic Acid from the Dithioformate, Sodium Hydrosulfide, and Each Alkyl Bromide.
Author;MURAOKA MOTOMU(Josai Univ., Fac. of Sci.) YAMAMOTO TATSUO(Josai Univ., Fac. of Sci.) TAKAHASHI KENTA(Josai Univ., Fac. of Sci.) AOKI DAISUKE(Josai Univ., Fac. of Sci.)
Journal Title;Nippon Kagakkai Koen Yokoshu

ISSN:0285-7626

VOL.78th;NO.2;PAGE.1191(2000)

Language;Japanese
Abstract;Potassium dithioformate was prepared by the reaction of CHCl3 and K2S obtained from KOBut and H2S at mild conditions in high yield. Formation of precursor, HC(SK)3 was confirmed by isolating trialkyl orthotrithioformate. The orthotrithioformic acid ester was also obtained from HCSSK, NaSH, and RBr. (author abst.)

To reiterate, acidifying this salts liberates dithioformic acid (unstable), reacting trhat with aq a,,onia gives thioformamide in 30% yield, and pyrolysis of thioformamide in inert atmosphere gives CS2 but until I have yje Wilstatter and Wirth paper from Ber.42 I do not know yield.

So this is a method for turning chloroform into CS2 in 304 steps.

Reid section on dithioacids attached.

[Edited on 25-5-2009 by Sauron]

Attachment: Pages from Reid V4.pdf (566kB)
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[Edited on 25-5-2009 by Sauron]

Sauron - 24-5-2009 at 22:49

The Rewards of Scholarship

Recourse to Reid's book also git me a complete cirarion of the Levi paper:

T.G.Levi, Atti Accad. Lincei (5) 32. I, 560-572 (1923)

and more importantly a Chem.Abstracts citation

CA 18, 1114

which may actually be accesible.

Eclectic - 25-5-2009 at 06:49

Just a stray thought: Does anyone know if liquid sulfur reacts directly with calcium carbide? How about H2S and calcium carbide?

[Edited on 5-25-2009 by Eclectic]

Nicodem - 25-5-2009 at 08:04

Quote: Originally posted by Sauron

Recourse to Reid's book also git me a complete cirarion of the Levi paper:

T.G.Levi, Atti Accad. Lincei (5) 32. I, 560-572 (1923)

and more importantly a Chem.Abstracts citation

CA 18, 1114

which may actually be accesible.

The entry was renumbered during the digitalization of Chemical abstracts, and it is now CA 18:8271. Here it is:

Quote:

Dithioformic acid. Levi, T. G.. Atti della Accademia Nazionale dei Lincei, Classe di Scienze Fisiche, Matematiche e Naturali, Rendiconti (1923), 32(i), 569-572. CODEN: AANLAW ISSN: 0001-4435. Journal language unavailable. CAN 18:8271 AN 1924:8271 CAPLUS

Abstract

L. has prepd. various derivs. of HCS2H from the K salt obtained by treating CHCl3 with K2S (cf. Cambi, C. A. 4, 1738) in alc. under definite conditions. The Ag, HCSAg, and Pb salts, (HCS2)2Pb, were obtained as orange-yellow ppts., but could not be prepd. quite pure, owing to their insoly. in ordinary org. solvents; both are comparatively stable; the former blackens at 95-100 and the latter at 130-5. The Zn salt is yellowish white and the Co salt deep brownish red. The disulfide, [(HCS)2S2]x obtained as a yellowish red ppt. by cautious oxidation of the K salt in alc. by means of I, decomps. at above 200 into CS2, H2S, C and S.

The same paper was also published in Gazzetta Chimica Italiana which is a more common journal in the libraries:

Quote:

Dithioformic acid. Levi, T. G.. Gazzetta Chimica Italiana (1924), 54, 395-397. CODEN: GCITA9 ISSN: 0016-5603. Journal language unavailable. CAN 18:18738 AN 1924:18738 CAPLUS

Sauron - 25-5-2009 at 09:06

Thanks, Nicodem. Howeverm it is less helpful than I hopes. While stating that the rxn of CHCl3 w/ K2S in alcohol required definite conditions, it fails to describe these, just bounces us to Cambi, CA 4 1738 - surely also now renumbered after digitization.

Can I trouble you once more, in light of this>

Every bloody paper describing use of this reagent cites Levi and not one gives prep details.

Alternative names for fithioformic acid:

thiolthionmic acid

carbodithionic acid

It apparently is prefered tautomer of initially former orthothioformic acid which like orthoformic acid has not been isolated. Well, not tautomer since orthothioformic acid must lose H2S to for, HC(=S)SH.

Since pyrolysos>200 C breaks potassium dithioformate down to CS2, H2S, S and C, then no need for liberation of the acid, and no need for the lossy comversion to thioformamide hydrochloride just tp pyrollize it instead. Therefore this is a two step conversion of chloroform to CS2 via KCS2H.

Gazzetta Chimica Italiana is yet another defunct jornal that Wiley boasts has been incorporated into its European Journals of Chemistry but unlike Ber. and Ann. I have not been able to locate an index page for it at Wiley's site.

There is sporadic coverage of this journal on Internet Archive via Googlebooks. But sadly nothing after 1908.

[Edited on 25-5-2009 by Sauron]

Sauron - 25-5-2009 at 23:27

Potassium Ethyl Xanthate is pricier than I remembered, so its decomposition, while facile and efficient, is not an economical way to obtain CS2.

The conversion of chloroform to potassium dithioformate and hence to CS2 requires K2S, and therein lies the rub. The Japanese use tBuOK to make this from H2S. tBuOK is not expensive, c.$30/mol. K2S is very air and moisture sensitive and so this is a glove box procedure. The commercial solutions of tBuOK are too bloody expensive.

The classical prep of K2S in from the elements in liquid NH3, see Brauer.

Nicodem - 25-5-2009 at 23:50

Quote: Originally posted by Sauron

Thanks, Nicodem. Howeverm it is less helpful than I hopes. While stating that the rxn of CHCl3 w/ K2S in alcohol required definite conditions, it fails to describe these, just bounces us to Cambi, CA 4 1738 - surely also now renumbered after digitization.

I can not find that one. There is only one CA entry for Cambi as author in the years corresponding to vol. 4 of CA. In the other references of this author I found only one mention of the reaction of Na2S with CHCl3 but the product is used in further reaction without its isolation.

Quote:

Thiohydroxamic Acids. Cambi, Livo. Lab. Chim. Farm. R. Inst. Studi Sup., Firenze, Atti. acad. Lincei (1910), 18(I), 687-90. Journal language unavailable. CAN 4:9654 AN 1910:9654 CAPLUS

Abstract

Thiobenzohydroxamic acid, PbC(: NOH)SH, obtained by adding to an aqueous soln. of PhCS2K a soln. of NH2OH-HCl containing the equiv. amt. of K2CO3, acidifying with dil. H2SO4 and extracting with Et2O. Its salts are obtained from the H2O soln. by evaporating to dryness in vacuo and extracting with EtOH. Both the acid and its salts easily lose S, the free acid giving chiefly PhCN, S and H2O. Benzyl ester, PhC( : NOH)SCH2Ph, from the K salt and PhCH2Cl in EtOH, m. 120-2; benzoyl derivative, PhC( : NOBz)SCH2Ph, from the ester and BzCl in pyridine, m. 135. Dibenzoyl derivative, PhC(:NOBz)SBz, from the K salt in H2O with excess of BzCl and subsequent addition of NaOH, m. 90-2 : stable towards dil. HCl, even on boiling, but saponified by not too conc. alc. KOH, partly into PhC( : NOH)SH and BzOH, but chiefly into PhC(: NOH)OH and PhCOSH.

Action of Hydrogen Sulfide on Fulminic Acid. Cambi, Livio. Milano, lab. elet. chim. r. ist. techn. sup., Gazzetta Chimica Italiana (1911), 41(I), 166-73. CODEN: GCITA9 ISSN: 0016-5603. Journal language unavailable. CAN 5:15314 AN 1911:15314 CAPLUS

Abstract

cf. C. A., 4, 1738. When a suspension of Hg fulminate in H2O is treated in the cold with H2S, the filtered soln. gives an intense violet-blue color with FeCl3 changing to blue on the addition of dil. HCl or H2SO4 and after standing a while, especially on heating, the soln. gives the reaction of HSCN. With excess of NaCO3 is obtained an amorphous sodium salt, sol. in alc., which, with PhCH2Cl, gives benzyl thioformhydroxamate, C7H7SCH : NOH, elongated prisms, m. 144-6, decomp. by boiling in the presence of HCl into C7H7SH, HCO2H and NH2OH. The same Na salt can be obtained by treating CHCl3, in alc. with 2 mols. Na2S and, after removal of the NaCl and alc., adding NH2OH.HCl satd. with Na2CO3 to the H2O soln. of the residue. With the heavy metals it gives amorphous ppts.: Cd, white; Pb, yellowish; Hg, yellowish; Ni, reddish brown; Co, brown; Cu, black; Ag, light yellow rapidly blackening. The free thioformhydroxamic acid could not be isolated, as it decomp. into HSCN; the Na salt in soln. gives NH3 and when heated in the dry state forms NaSCN. Similarly, the benzyl ester, heated with Bz2O at 80-5, gives C7H7SCN.

Sauron - 26-5-2009 at 02:31

Many thanks. That must be it.

watson.fawkes - 26-5-2009 at 09:21

FYI. There's another drawing of the Taylor furnace in the 1921 book A course in general chemistry, p 885. This drawing looks a little more realistic. In particular, it shows a clean-out port on the bottom of the sulfur retort. Fine coke/charcoal dust settles to the bottom, clearly.

Sauron - 26-5-2009 at 10:16

If anyone wants to know why you can't just bubble H2S into aqueous KOH to obtain K2S, the answer is

KOH + KSH equilibrates with K2S and H2O

So the anhydrous salt has to be made by union of the elements in NH3 (not convenient) or by psdding H2S into s nonsqueous solution of tBuOK (such as in THF).

THF and tBuOH are removed and the K2S dissolved in anhydrous ethanol for reaction with chloroform.

HCS2K precipitates, ethanol and KCl are removed.

Potassium dithioformate. so simplr yet so obscure.

Incidentally Aldritch lists K2S but ONLY AVAILABLE IN JAPAN. WTF?

No one else I know of even offers it.

[Edited on 26-5-2009 by Sauron]

Sauron - 26-5-2009 at 17:31

There's a metos for preparing anhydrous K2S by adding molten KCl to liquid S. Sulfur chlorides form and volatilize, leaving K2S. Does this sound like fun?

len1 - 26-5-2009 at 18:21

That cant be right - at the melting point of KCl the S would long ago have volatalized.

watson.fawkes - 26-5-2009 at 19:20

Quote: Originally posted by len1

That cant be right - at the melting point of KCl the S would long ago have volatalized.

I imagine that the molten KCl is dripped in, forming volatile and extremely reactive S vapor. I would also guess that sulfur chlorides are a by-product.

Also, apropos of high-temperature processes. The following article is supposed to discuss a fluidized bed reactor for the production of CS2, should the footnote be accurate: Johnson, H.S., Reactions in a fluidized coke bed with self resistive heating, Can, J. Chem. Eng., 39, 145, 1971.

Sauron - 26-5-2009 at 23:08

It's in Mellor.See attached extract, Vol II, p.623 top of page. The citation is p 628m para 2 of References,

Bemelmans, L., D.R.P. 49628, (1894)

I have not read the German patent but apparebtky the nobject was to make sulfur chlorides, and K2S is a byproduct. As for my criticisms of the process as described by Mellor:

Mp KCl is 770 C, bp S is 444 C. I think it's a joke since K2S if formed will react with excess S to form polysulfide. Reversing the addition would be better, but the S melt to be fluid rather than viscous would have to be at lower end of its range, and unless the KCl melt can be agitated local excess of S is still possible. In short is seems a good way to make crappy livers of sulfur rather than nice clean well defined K2S anhydrous.

The section is pretty much a recitation of such bad methods.

Only the union of S with a solution of K in liq NH3 of all the methods mentioned will produce clean dry salt, which has to be rigirously protected from air and oisture.

The Japanese method, H2S and tBuOK, is of course nowhere in Mellor. It is IMO more convenient, and avoids handling K. but does involve H2S.

Attachment: Pages from DjVu Document.pdf (796kB)
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[Edited on 27-5-2009 by Sauron]

Sauron - 27-5-2009 at 03:24

There is also a prep from the elements conducted between a solution of S in toluene or napthalene, and metallic K. I will have a look at this one. I reckon if the solvent us dry and free of dissolved air amd the whole thing conducted in inert atmosphere, this could fly.

J.Locke, A.Austell, Am.Chem.Journ. 20, 592 (1898)

[Edited on 27-5-2009 by Sauron]

Attached. Unfortunately this is one of those times when Mellor got it all wrong.

[Edited on 27-5-2009 by Sauron]

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S.C. Wack - 27-5-2009 at 17:26

Quote: Originally posted by Sauron

If anyone wants to know why you can't just bubble H2S into aqueous KOH to obtain K2S

I want to know why, especially since all the old references say you can. This method for it and Na2S are mentioned a lot in Mellor, for instance. Dehydration seems unchallenging.
http://dx.doi.org/10.1039/CT9007700753

garage chemist - 27-5-2009 at 17:45

Aqueous K2S will always be an equilibrium mixture of KOH and KHS (with mainly K2S), no matter how it was prepared. Na2S is exactly the same, it's made by saturating NaOH solution with H2S forming NaHS, then adding an equimolar amount of NaOH.

t-BuOK, come on, let's keep it affordable and OTC. CS2 preparation in a home lab becomes somewhat pointless if it requires chemicals from a chemical supplier. If someone can buy stuff like t-BuOK readily made, he could just as easily buy CS2.

Quote: Originally posted by Eclectic

Just a stray thought: Does anyone know if liquid sulfur reacts directly with calcium carbide? How about H2S and calcium carbide?

[Edited on 5-25-2009 by Eclectic]

I have tried this myself, both with molten sulfur and calcium carbide, and with sulfur vapor and calcium carbide at 440°C (the carbide was put into a Hempel column, which was connected to a flask in which sulfur was being boiled). There is no reaction at all.
H2S and calcium carbide would probably give acetylene and calcium sulfide, as H2S is a proton acid like water, an even stronger one actually, and will protonate the acetylide anion in CaC2.

S.C. Wack - 27-5-2009 at 18:14

Quote: Originally posted by garage chemist

Aqueous K2S will always be an equilibrium mixture of KOH and KHS (with mainly K2S), no matter how it was prepared.

I am not talking about an aqueous mixture, this would be useless for the reaction with chloroform! I am talking about crystalline hydrates, of either potassium or sodium.

[Edited on 28-5-2009 by S.C. Wack]

Sauron - 27-5-2009 at 19:26

There's a procedure for dehydrating sodium sulfate nonahydrate by vacuum dessication over H2SO4. It takes `4 days and the material still contains 4% water. To drive that off takes high heat.

No such technique is mentioned for K2S.

The problem with Mellow on this substance is that sometimes his entries are too elliptical and either misrepresent the actual report, if you read the original, or are ambiguous.

A case in point is the one I cited above, Locke and Austell. Mellor states they made Na2S from the elements by boiling in toluene or napthalene. I obtained the volume 20 of Am Chem J.
The paper says they wanted and expected Na2S but insteas obtained a higher oligosulfide estimated to be Na2S3. A ;ot of unreacted Na remained.

In short the prep failed, but Mellor reported it as a success. Mellor is sometimes quicksand. In many instances, the original full trxt is well nigh impossible to obtain soone must decide whether to rely on Mellor alone.

It IS known and reported in Mellor as well as the lit. old and new that the same reaction of the elements works fine in liq NH3 - this is also descrived in Brauer. This produces rigorously pure white crystalline K2S quaantitatively from stoichiometric quantities of K and S. But it is inconvenient.

Garage chemist, I am NOT advancing the reaction of tBuOK amd H2S in EtOH as a way to make K2S cheaply enough to exploit the chlorodorm reaction to make CS2.

ou have the economical route to CS2 from S and acetylene. Be happy. There is no other unless you want to return to scaling down S + C.

S.C., it is clear that many if not all of the old methods did NOT produce inambiguous K2S, but mixtures of that with polysulfode. KSH, thiosu;fate, KOH, carbonate etc. Preparing K2S in air will have that result regardless of method. It might be possible to isolate, quite laboroously, purer crystalline material from such mixtures. But if you want pure unambiguous K2S colorless crystals, not yellow and with correct constants, there are only two methods. From the elements in liq NH3, or metathesis of tBuOK and H2S in ethanol.

Sauron - 27-5-2009 at 20:20

I read the paper you cited, S.C., and saw prep of K2S.2H2O but no anhydrous materiall.

S.C. Wack - 27-5-2009 at 20:24

Quote: Originally posted by Sauron

There's a procedure for dehydrating sodium sulfate nonahydrate by vacuum dessication over H2SO4. It takes `4 days and the material still contains 4% water. To drive that off takes high heat.

No such technique is mentioned for K2S.

So? There are numerous examples of dehydrations of the hydrates of both that make it clear that it is not difficult as long as air is excluded.

Quote: Originally posted by Sauron

The problem with Mellow on this substance is that sometimes his entries are too elliptical and either misrepresent the actual report, if you read the original, or are ambiguous.

A case in point is the one I cited above, Locke and Austell. Mellor states they made Na2S from the elements by boiling in toluene or napthalene. I obtained the volume 20 of Am Chem J.
The paper says they wanted and expected Na2S but insteas obtained a higher oligosulfide estimated to be Na2S3. A ;ot of unreacted Na remained.

In short the prep failed, but Mellor reported it as a success. Mellor is sometimes quicksand. In many instances, the original full trxt is well nigh impossible to obtain soone must decide whether to rely on Mellor alone.

Jesus! All I did was mention the dudes name since he, among other sources available, gives several references, and you are familiar with the chapter (20) apparently. None of that has anything to do with the validity of papers where the authors unambiguously state that they have monosulfide as whichever hydrate.

For all we know, this is what whoever used with chloroform instead of a commercial sulfide.

Quote: Originally posted by Sauron

S.C., it is clear that many if not all of the old methods did NOT produce inambiguous K2S, but mixtures of that with polysulfode. KSH, thiosu;fate, KOH, carbonate etc.

Says who? About what? Who says that Bloxam, for instance, is full of it?

Sauron - 27-5-2009 at 21:07

Quicksand.

The Italian papers unavailable.

So you are right, we do not know whether or not anhydrous K2S was used, or a hydrate.

My assumption is anhydrous, the Japanese method unambiguously used anhydrous K2S to prepare potassium dithioformate. Preps of K2S in liquid NH3 cited by Mellor predate Cambi's paper and Levi's paper. C.Hugot, Compt.Rend 129, 388 (1899) so method which is still the main method for anhydrous K2S was known a decade before Cambi and 24 years before Levi.

However, it is possible that cambi and/or Levi used a hydrate.

We simply do not have enough information to be sure.

If I did not find Mellor useful I would not be scanning his emtire series. HOWEVER he is not perfect, he makes mistakes, and is also capable of ambiguity. My point is that the original cited texts need to ve read whenever possible. The same goes for Beilsteinm and Gmelin. All of them are just reviews of the work of others. Vast and precious but imperfect.

[Edited on 28-5-2009 by Sauron]

Nicodem - 27-5-2009 at 23:01

Quote: Originally posted by Sauron

However, it is possible that cambi and/or Levi used a hydrate.

We simply do not have enough information to be sure.

As far as I understood they used ethanol as solvent in which case it is irrelevant if the sodium or potassium sulfide is hydrated or not. It is not even important if it is sulfide or hydrogen sulfide since both nucleophiles lead to the same compound at the end (the "orthotrithioformates" are intermediates anyway). In cases where you are in doubt it is always a good idea to check the reaction conditions, such as the solvent used, and compare it to the up to date theory.

Sauron - 28-5-2009 at 00:07

True, but we have on one hand Levi's prep as repeated by the authors of the JCS paper cited upthread, in which the yield of potassium dithiocarbonate was something less than 40%; and on the other hand the Japanese abstract from year 2000 which claims a high but unstated yield using K2S prepared from H2S and tBuOK. That prep of K2S seems to be sole novel feature of the procedure. So perhaps the matter of how the K2S is prepared and whether or not it is hydrated, is not so immaterial to the outcome of the dithioformate prep.

The Japanese also ran that reaction in alcohol, presumably ethanol.

While Na2S hydrate is cheap, and the anhydrous salt available but expensive (Alfa $5/gram in 100 g pack) potassium sulfide seems unanailable either hydrated or anhydrous, save Aldrich which has its sale restricted to Japan only. What the hell is that all about?

Sauron - 28-5-2009 at 01:27

There's quite a bit of lit. on sodium dithioformate as source of dithioformic acid used in thiazole preps.

Nicodem suggests that in ethanpl. hydrated monosulfide ought to work fine, so Na2S.9H2O which is common ACS reagent ought to convert CHCl3 to sodium dithioformate and thus to CS2, about 50% or more on NaCS2H basis by thermolysis.

If so we may have a whole new ball game.

Nicodem says KSH ought to work. Na2S.9H10 is $80/Kg,

NaSH is much cheaper, <$50 for 2.5 Kg and <$150 for 10 Kg.

According to a paper I post a few posts fown, KSH worksand quite exothermically so then major cost will just be the chloroform.

[Edited on 28-5-2009 by Sauron]

Eclectic - 28-5-2009 at 04:10

Hummm.....We have a procedure that needs acetylene and produces H2S, and another that needs H2S....

Might be worthwhile to see what the products are from passing H2S into a suspension of calcium carbide in ethanolic KOH.

Eclectic - 28-5-2009 at 08:04

Well you need a scrubber for the H2S produced by the acetylene process....Might as well use it to make NaSH, Na2S, KSH, or K2S in ethanol, with CaC2 possibly removing the water and making more acetylene.

Nicodem Is Correct

Sauron - 28-5-2009 at 08:08

A paper has turned up from Bull.Chem.Soc.Japan which details:

-- prep of KSH from methanolic KOH and H2S
- reaction of that solution with chloroform (vilent!)
- Use of crude potassium dithioformate solution in thioformylation without isolation

Paper is attached.

This is my modofocation of their procedure.

A solurion of KSH is prepared by mixing a solution of 46 g KOH (820 mmol) in 150 ml MeOH with 150 ml Ml MeOH which has been saturated with H2S and containing a further 46 g KOH. Total KOH 1.64 mol. The solution is transferred to a 3 L flask equipped with a Dimroth condenser. The solution is heated to 50 C and 30 g chloroform (250 mmol)added. The flask id dhaken constantly. An exothermic reaction sets in leading to violent boiling. After ten minutes the mixture is cooled, KCl filtered off, and filtrate rotavaped to dryness. The crude potassium dithioformate can be thermolyzed to CS2 without further purification.

CHCl3 + 4 KSH -> HCSSK + 3 KCl + 2 H2S

Obviously do this in hood.

The oversize flask and efficient reflux condenser are necessary to avoid loss of reactants and solvent. The violence of the reaction may present a scaleup problem.

Discontinue heating when the exotherm kicks in.

Attachment: BCSJ.pdf (538kB)
This file has been downloaded 967 times

[Edited on 28-5-2009 by Sauron]

Sauron - 28-5-2009 at 20:33

The theoretical yield from above ought to be 250 mmol CS2. That is 19 g. CS2 d 1.26 so about 15 ml.

To do this on a 1 mol CHCl3 basis use a 12 L flask, 120 g chloroform, and prepare KSH by mixing:

184 g KOH in 600 ml MeOH

and

184 g KOH in 600 ml MeOH saturated with H2S.

You will need one or more large efficient reflux condensers. Dewar condensers, double surface coil condensersAllihn condensers and Friedrichs condensers all come to mind.

[Edited on 29-5-2009 by Sauron]

len1 - 28-5-2009 at 22:40

Quote:

The crude potassium dithioformate can be thermolyzed to CS2 without further purification.

Is there a yield for this? Many compounds of C S and H have some CS2 in the decomposition - it does not sound like a high yielding process

Nicodem - 28-5-2009 at 23:34

At the moment I do not have the time to read the whole thread and find the citation regarding the dithioformate thermolysis, but Len's question makes good sense, because if it is just a thermolysis and no other reagent (or air) is used, then you obviously need a redox reaction to transform HCSSK to CS2.
Since it is unlikely for this compound to decompose to CS2 and KH or any such thing, then the formation of CS2 must be the consequence of a disproportionation at the carbon. This means the stoichiometry between HCSSK and CS2 is 2 : 1. In practice this means half less CS2 than one would hope for, and also the formation of the other disproportionation product(s). I have no idea what is the other disproportionation product, but one option could be K2S and trithiane, or KSCH2SK (depending on which form is thermodynamicaly more favourable). Or maybe it's just a dirty thermolysis with CS2 just as one of the many products?

Sauron - 29-5-2009 at 00:39

According to Reid, in extract posted upthread, potassium dithioformate melts with decomposition at 193 C. Dithioformic acid, solid, liberated from above, melts with decomposition at 55-60 C. and as no pesky K is present we can work out the reaction:

HCSSH -> CS2 + H2

One mol HCSSK MW 116

gives one mol HCSSH 78 g

which theoretically gives 76 g CS2

The free acid must be decomposed as soon as prepared or it will polymerize.

Details on acidolysis are in Levi but are very likely to be similar to those for dithiocarbonate -> xanthis acid, which are in Mellor on Carbon Part II. Cold dil aq mineral acod is used. Dithioformic acid is insoluble and ppts out as a buff colored powder.

Potassium ethyl dithiocarbonate (xanthate or xanthogenate)

KSC(=S)OEt

liberates xanthis acid HSC(=S)OEt

which falls apart in minutes at RT to CS2 and EtOH

The structural analogy to dithioformate and dithioformic acid is clear

HC(=S)SH

Under oxidizing conditions one might get COS and H2S but at 55 C? To be certain, do it under N2.

[Edited on 30-5-2009 by Sauron]

Sauron - 29-5-2009 at 04:42

Reid also states that CCl4 can be used in place of CHCl3 in this reactuin. He cites Leci again, same paper and a follow on.

I wonder what the stoichiometry is?

A completely different way

watson.fawkes - 29-5-2009 at 05:11

This method applies mostly only to Sauron, but I'm sure that the Thai Rayon plant in Angthong Province has CS2 on premises. If they sell it officially, it's likely to be by the drum.

Sauron - 29-5-2009 at 11:13

This review has popped up, right on point.

Ita states that purer HCSSK is obtained when the reaction is run in mrthanol tather than ethanol and that the reaction should be done under an inert atmosphere.

Read along with me,

Attachment: Pages from rc45.pdf (645kB)
This file has been downloaded 1544 times

Sauron - 29-5-2009 at 23:18

References 15 abd 16 from the Russian reciew above are actually two short notes back to back on a single page from same authors, covering potassium dithioformate and dithiofomic acid respectively.

Naturwissenschaften, 58 (1), p.53 (1971)
G. Gattow, M. Dräger and R. Engler
Über Dithioformiate

Attached.

Apparently the salt dimerizes ar its mp to sulfide KCSSSCK rather than fragmenting. That is (KC(=S)2S. This is a liquid d 1,74.

I am stiill reading so the other shoe has not dropped yet. What does duthioformic acid do at the mp?

In the free acid mass spec peaks correspond to the diner, the monomer, and CS2, which sounds promising. HCSSH is monomer, (HCSS)x the polymer, the dimeric disulfide x=2. Not the same as with the salt. In both cases H2 is extruded as I surmised.

Note that the polymerization is fairly rapid at RT, see Todd's paper (JCS 1837 for timeline of MW growth over 24 h. So the acid is stored as the salt, liberated w/dil HCl and thermolyzed at once at 60 C or whatever gives best yie;d of CS2.

Attached

Attachment: nat71.pdf (137kB)
This file has been downloaded 535 times

[Edited on 30-5-2009 by Sauron]

Formatik - 31-5-2009 at 15:36

Since the work of Levi [1] concerning dithioformates, these salts have been only briefly mentioned in the literature [2, 3]. By the described method of preparation one obtains however a strongly impure product.

K[HCS2] (Mp. = 196°C, decomposition under formation of K2CS3; d20/4= 1.737 g/cm³) was prepared by the reaction of chloroform with a methanolic solution of K2S under the absence of air; from the concentrated down solution, the substance crystallized out in the cold. - The tetraphenylarsonium salt [.phi.4 As] [HCS2] (Mp. = 191-193°C, dec.: d20/4 = 1.39 g/cm³) results as a barely soluble precipitate by the addition of [.phi.4 As]Cl to an aqueous solution of K[HCS2], and it can be recrystallized from water.

Aqueous solutions of the [HCS2](-) ion show visible and UV-range absorptions at 225, 331, and 386 nm.

IR spectrum of K[HCS2]: .... [spectral data for the potassium and tetraphenylarsonium salt]

Dithioformates and dithioformic acid, whose existence was first pointed out by Levi [1], are only briefly mentioned in the following literature [2,3].

Dithioformic acid was precipitated from an aqueous solution of K[HCS2] by the addition of dilute hydrochloric acid. It is as opposed to formic acid and monothioformic acid [5], barely soluble in water. Aqueous sodium hydroxide as well as tetrahydrofuran and dioxane solubilize the free acid well; in ether, chloroform and benzene it has only low solubility.

The weak yellowish, X-ray amorphous acid [HC(S)(SH)]x melts at 55-60°C. In the mass spectrum, aside the molecule peak at m/e = 234 (x=3), there are peaks at m/e= 156 (x=2), 78 (x=1) and 76 (CS2). Solutions of dithioformic acid in diethyl ether show in the UV-range absorption at 295 nm.

IR spectrum of dithioformic acid: [spectral data concerning the said compound and its ethyl and methyl ester briefly]

Sauron - 31-5-2009 at 15:50

Thank you, formatik.

Sauron - 1-6-2009 at 22:34

To recap, it appears that potassium duthioformate does not thermolyze to CS2 but that dithionic acid might well.

Also treating the acid with aqieous NH3 obtains thioformamide, and Wilstatter and Wirth (posted above) report that the hydrochloride of that compound, prepared with HCl gas in ether, does thermolyze to CS2.

Is this competitive with acetylebe and S? That depends on circumstances. Os chloroform readily available and cheap? If so then just maybe this is a viable alternative.

Rattata2 - 13-2-2010 at 02:40

Quote:

[quote=Sauron]Potassium Ethyl Xanthate is pricier than I remembered, so its decomposition, while facile and efficient, is not an economical way to obtain CS2.

This post gave me an idea..how about the decomposition of cellophane or rayon to obtain CS2? That stuff is dirt-cheap and used everywhere..and is a xanthate