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len1
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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.
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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]
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watson.fawkes
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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.
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watson.fawkes
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| 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
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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.
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Sauron
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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
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And Physical Preoperties
Attachment: CS22 Physical Props.pdf (371kB)
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len1
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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
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Sauron
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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]
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watson.fawkes
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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.
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Sauron
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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.
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watson.fawkes
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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.
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Sauron
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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.
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Eclectic
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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]
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Sauron
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Perfect job for a heating tape. Wrap the Vigreaux and use a simple controller to hold it at 350 C +/- 50.
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Sauron
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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
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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]
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Sauron
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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
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| 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.
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Sauron
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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.
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len1
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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.
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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.
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len1
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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.
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Sauron
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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.
Sic gorgeamus a los subjectatus nunc.
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len1
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No snide remark - just a joke, they are OK too sometimes. Well done to GC for checking this and to Sauron for fiding this.
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Sauron
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CHCl3 -> CS2 ?? A Suprise
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)
This file has been downloaded 835 times
[Edited on 25-5-2009 by Sauron]
Sic gorgeamus a los subjectatus nunc.
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Sauron
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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.
Sic gorgeamus a los subjectatus nunc.
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