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Sauron
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CS2 a Different Way
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.
Sic gorgeamus a los subjectatus nunc.
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panziandi
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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.
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garage chemist
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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.
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S.C. Wack
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http://dx.doi.org/10.1039/JR9280002068
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len1
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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
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Sauron
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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)
This file has been downloaded 1021 times
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garage chemist
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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.
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Sauron
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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]
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len1
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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]
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Sauron
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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.
Attachment: cr60123a002.pdf (270kB)
This file has been downloaded 1320 times
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Lambda-Eyde
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Quote: Originally posted by Sauron  |
Fractionation affords CS2 of reasonable purity (garlic odor suggestive of allyl sulphide contaminant)
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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.
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Sauron
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The thiophenol need only be present in very low concentration to impart that odor. It's obnoxious stuff.
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Lambda-Eyde
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Make up your mind, are you talking about allyl sulfide or thiophenol? 
BTW, commercial acetylene is contaminated with acetone, just so you know.
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Sauron
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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?
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Lambda-Eyde
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| 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.
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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.
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len1
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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]
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panziandi
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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.
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watson.fawkes
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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. |
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BromicAcid
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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.
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len1
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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.
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BromicAcid
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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.
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watson.fawkes
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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.
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len1
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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]
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watson.fawkes
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Smells
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.
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watson.fawkes
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Please ignore this.
[Edit:] Oops. Wrong. WRONG. Please disregard.
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]
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