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BromicAcid
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[*] posted on 3-11-2003 at 10:00
Carbon Disulfide Preparation


Quite some time ago I produced a very small amount of carbon disulfide (like 2 or 4 ml) and it completely fouled up my glassware. I'm gonna make another run at it soon and I wanted some feedback. Industrially and pretty much the only way carbon disulfide is made is the reaction between coke and sulfur at high temperatures. My reaction was with air float charcoal and sulfur in glass distillation apparatus.

I am going to replace the airfloat with activated carbon ground down to a coarse texture, more surface area should yield a higher reaction rate. From what I've read the reaction rate is hampered by excess carbon because of its low heat conductivity.

Next up I'm doing the reaction in an all galvanized iron contraption. It's a two inch diameter pipe six inches long with an end cap and the otherside having a pipe of approximately 3/4 inch diameter going down at a 45 degree angle for 18 inches ending in another 45 to put it straight down where it will end up underwater where the carbon disulfide will hopefully collect. I'm assuming that this is sufficient distance to condense out some of the carbon disulfide and I may fill the water with ice to trap any additional fumes.

Heating will be accomplished by an acetylene torch on low with a broad flame. Reaction temperature is supposed to be around the boiling point of sulfur so it should not be too hard to maintain.

One thing that I'm wondering, molten sulfur is supposed to be highly corrosive and texts recommend treatment with aluminum powder at high temperature to develop a layer of aluminum on the inside of the reaction vessel to reduce corrosion. The pipe is 1/16 thick and is zinc plated steel. I do not believe that burn through will be a problem with only one batch and not too excessive temperatures but I may be wrong.

The whole reason I'm posting is for feedback. I need carbon disulfide as a solvent for sulfur and phosphorus. The hazardous material charges through the mail are killer and the initial price is none to inviting either.

Thanks in advance.
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[*] posted on 3-11-2003 at 10:12


Quote:

From madscientis in the chlorate thread:
I really can't think of a simple solution to the inert-gas problem. Most of the time, the gas is going to get heated to autoignition temperature; however it will not combust if there is no oxygen present. Try boiling hexane and using the hexane vapors as an inert gas; higher autoignition temperature. I have produced CS2 which has an autoignition temperature of around 90C by heating sulfur and carbon in a pyrex flask at temperatures more like 500C; never have had any trouble. Mostly, just make sure you're using a flask, and do it outside (it case the gas being used does ignite). If it ignites while outside, it really doesn't matter.




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[*] posted on 9-2-2004 at 14:48


Yup, I agree with madscientist whole heartedly about not worrying about autoignition due to there being no oxygen present. The only thing that I worry about is the notoriously low yields from the reaction of sulfur with carbon. Today I came across a variation of the reaction from the 1800's. Mix sulfur with sugar, heat, sugar decomposes and creates an activated form of carbon intimately mixed with the sulfur, yields are supposedly much higher from this reaction. Sounds like a fun past time, anyone else hear about this method?



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[*] posted on 9-2-2004 at 15:56


I have read somewhere or other that sugar charcoal does not work as well as coke for this reaction, simply because sugar charcoal is too pure. The addition of a small amount of alkali carbonate is supposed to have a catalytic effect.



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[*] posted on 18-9-2004 at 17:14


Tried another run today to make carbon disulfide from carbon and sulfur, a mixture of 64 g sulfur and 12 g charcoal powder were mixed and the mixture heated with a torch. The resulting gasses from the reaction mixture were lead though a condenser and into water, I've read that carbon disulfide is occasionally condensed under water. This water bath was in a sealed vessel and the exit gasses were lead into the flame heating the reaction mixture.

The reaction vessel was not sealed well and SO2 kept leaking out, a constant stream of flammable gasses came out as exit gasses and they quickly attacked the galvanized layer on the pipe giving it four distinct circles of color. Problem was my yeild was 0. I got nothing. Carbon disulfide is slightly soluble in water so some may have been formed but solvated. The reaction was hot enough as sulfur boiled off and went into the condenser. It just didn't work well, just like the first time.

Maybe I'll add some carbonate tomorrow or the next day and try again like Polverone said.

Anyway, the main question that I have. The last time this method succeeded my carbon disulfide was highly contaminated with sulfur. Upon cooling sulfur crystallized out of it, and sulfur is fairly soluble in it to begin with. So how to get the sulfur out of the resulting CS2?

Redistillation, especially fractional with a nice column packed with glass beads would work well. However think about how fouled the glassware would get. I was thinking about shaking the resulting carbon disulfide with a metal that forms suflides easily, like iron shavings, aluminum turnings, steel wool, etc. Thereby removing the sulfur, but does anyone think it might reduce the CS2 itself, or possibly get too much out of control and start a fire?




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[*] posted on 18-9-2004 at 22:49


Sounds like you need two condensers, with the second being a long and cold one. In the reading that I've done, it seems that the S and C are never intimately mixed, either S vapor is passed over charcoal in a tube heated to medium red heat, or there are two layers/compartments in a retort, the S on bottom of course.

It can be freed from the last of the S by distilling with aq. hydroxide a few times. Ca(OH)2 suspension would probably be a good choice.

Have you read the Thorpe article (in Carbon) from his Dictionary of Applied Chemistry? It's a long article and good reading. My copy was not from the FTP, so I'm assuming that it's there intact. Starts on page 74 of the proper volume.

Fe, Al, and Zn all play nice with CS2.
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[*] posted on 19-9-2004 at 04:41


I also think that you shouldn't mix carbon and sulfur and heat the mass, because the carbon needs to be at 800°C to react!
The apparatus should consist of a sulfur vaporizer which produces gaseous sulfur. This should then be lead into a pipe filled with charcoal that is maintained at read heat.
You'll need two burners! Or the charcoal pipe is led through a furnace that uses charcoal as the fuel. You also need ice water to condense the fumes!
Charcoal doesn't only consist of carbon but it also contains some hydrogen. Be careful as H2S could form!

Also, I don't understand why you would need a fractional column to redistill CS2 from the contaminating sulfur. Sulfur boils at around 400°C so a simple distillation should yield nearly sulfur- free CS2.


I strongly recommend you to try the production of disulfur dichloride instead. It is a GREAT solvent for sulfur (sulfur dissolves up to 67% in it) and it also dissolves white phosphorus well.
It has a high boiling point, so no problem to condense the fumes.
Also, only a simple distillation apparatus with a claisen adapter with the chlorine feeding tube inserted airtight into it is needed.

I have heard from other amateur chemists which also tried to produce CS2. They all failed.
S2Cl2 production was a success the first time I tried it.
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[*] posted on 19-9-2004 at 06:26


Garage chemist, you said yourself that when you made S2Cl2 that your vessel had a thin coating of powdery sulfur all along the inside. It does the same when distilling the carbon disulfide, the sulfur will carry over with it at the boiling point and deposit on things. Even though it is a considerably lower boiling point then the sulfur itself. So I was trying to think up something that I could add to the impure carbon disulfide obtained so I could distill it from there and it would take up the excess sulfur.

I know the design of industrial CS2 production systems but originally it was just made in a metal retort with the beak dipping below water with a 2 fold excess of sulfur and heated to red heat. That is what I was going for. But I had better success when I did it last using a different kind of charcoal and heating on a hot plate.

And yes, I got H2S, but I was burning my exit gasses, I only smelled burning sulfur until I actually opened up the vessel.




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[*] posted on 19-9-2004 at 06:41


Fill your iron vessel with pieces of irontubing, length 1,25xdiameter, this will give a better heatdistribution through the mass and should solve the problem of the charcoal being an insulator. A vertical piece of irontube as outlet filled with SS-wool as air-cooled condensor will recondense most of the sulfur, withdraw the CS2 from top and condense in a second iron condensor cooled with water, Liebig style.
Depending on the volume of your reaction vessel I would suggest the iron-Raschig rings to have a diameter of 0,5-2 cm.
You will need a stronger flame/heatsource though.

Just suggestions, but with some background..... :D




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[*] posted on 19-9-2004 at 21:48


The iron tubing idea sounds promising, fill a larger iron pipe with tubing in the bottom to make it honey comb like and fill with mixture, it should conduct heat well.

Aside from alkali carbonate does anyone know any other possibly catalytic additives? I could see that many metals might prove to be catalytic, forming initally the sulfide then being reduced but I can't find any literature on anything specific.

I found lots and lots of patents on the reaction between molten sulfur and both low weight and high weight hydrocarbons, this method gives much lower reaction temperatures, 350 - 700 C but of course the main by product is H2S, however with me burning the exit gasses this is not a problem, my attempt at CS2 yesterday mad lots and lots of H2S but I didn't smell any of it till I opened up the vessel later because I was burning it as it was formed. So using propane or butane might prove interesting.

Another thing that I found, only one patent relating to it was the reaction between chlorinated hydrocarbons and sulfur. The main example being hexachlorohexane.

2C6H6Cl6 + 3S8 ---> 12CS2 + 12HCl

If the ratio of hydrogen to carbon was not equal it warned, sulfur chlorides would be formed. It also stated that the reaction above, and others were exothermic. I've got some hexachloroethane laying around, maybe mix some together and see if I get a pyrotechnic mixture.

4C2Cl6 + 5S8 ---> 8CS2 + 12S2Cl2

Purificatoin would be simple, agitate the mixture with water, however I would have to be sure to wash my exit gasses, maybe agitate the mixture with basified water.




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[*] posted on 20-9-2004 at 14:20


I found my hexachloroethane but I figured it was too chlorine rich and I used some dichlorobenzene to make the chlorine to hydrogen ratio equal as the patent suggested:

C2Cl6 + 3C6H4Cl2 + 5S8 ---> 12HCl + 20CS2

I need to run this reaction in a closed vessle because sulfur burns on it's own, I just tested to see if it would burn, and it does, barely, the flames were not sooty though as they are when either of the chlorinated hydrocarbons do when burned alone, so that's a good sign. I will try this in a closed vessel latter this week hopefully.




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[*] posted on 21-9-2004 at 16:47


I've decided it might be time to scale down the industrial process.

Propane is readily available and when passed over hot sulfur and reacted with sulfur vapor at a relatively low temperature (450 - 600C) for 3 - 6 seconds 99% conversion to CS2 is obtained. Propane, being readily available isn't a problem to work with, sulfur isn't the problem.

I probably need to pre heat the gas entrance tube, pack it with something inert to just make sure the gas is nicely heated. Then I need a strong heat source on the sulfur, and yet another heat source on the final reaction tube, packed with something like alumina, ceramic pieces, or iron pyrite.

Scaling down an industrial process, I fell like Axehandle ;)




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[*] posted on 21-9-2004 at 17:29


That would have to be with the careful exclusion of air, of course. If sulfur vapor can reduce propane to CS2, the hydrogen in it (or at least some of it) would probably form H2S. (Hydrogen and hydrogen polysulfides are also likely.) This is deadly poisonous as well as inflammable, and is something you will have to take care of when condensing out the CS2.

John W.
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[*] posted on 21-9-2004 at 17:43


Very true John, it's actually a total combustion of the propane molecule with the sulfur instead of oxygen, so quite a bit of H2S formed.

5S8 + 4CH3CH2CH3 ---> 12CS2 + 16H2S

Although I saw a reaction where the first step was the reaction of the hydrocarbon with the catalyst for the purpose of making finely divided carbon:

CH3CH2CH3 ---> 3C + 4H2

But that is not my goal, the simple oxidation of propane with sulfur and basically every other hydrocarbon seems to be time tested, from what I saw they switched over to it around the turn of the last century and have been developing variations of it ever since, the lower weight hydrocarbons, with 3 or less carbons show the highest reactivity and being that they are so easily acquired I might give them a shot.

The plan is to run the exit gasses though a condenser to condense out the CS2 and bubble though water to keep the CS2 covered, the gas then exits and goes into the burner to burn it off.

However when I tried CS2 production recently..... there was a lot of SO2 generated from my burning of H2S containing exit gasses, so maybe I'll just precipitate some metal sulfide simultaneously so I don't have to put up with SO2.

US patent 4,073,868 is a prime example of this sort of reaction although I have seen it in many other places. Here is an exerpt:

Quote:
In this Example a narrow stream of propane heated to 425.degree. C is injected countercurrent concentrically into a much wider stream of sulfur preheated to 700.degree. C to react substantially adiabatically at a pressure of about 40 psig. The flow rates are so controlled that the amount of sulfur is about 34% in excess of that required for the stoichiometric reaction with the propane to form carbon bisulfide. The reaction is effected in a short reactor, the residence time therein being 0.61 second, and the reaction mixture is then immediately quenched, first in a vessel at 140.degree. C (thereby condensing the sulfur in the reaction mixture). The non-condensed gases, including carbon bisulfide, then pass through a pressure-regulator (set to provide a back pressure of 3.7 atmospheres, i.e., 40 psig) from which the gases are passed to a condenser at 0.degree. C and under pressure to condense carbon bisulfide; non-condensed gases are vented at atmospheric pressure. The propane is injected through a 0.318 cm diameter circular orifice into the stream of sulfur flowing in a circular pipe having an internal diameter of 2.09 cm.

The calculated mixing temperature is about 675.degree. C.

The propane is converted substantially quantitatively (over 99%) and analysis of the condensed carbon bisulfide indicates that it has a purity of 99.89%, about 0.02% of benzene, about 0.09% of thiophene and no toluene. The condensed sulfur contains only traces of carbonaceous material.


[Edited on 9/22/2004 by BromicAcid]




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[*] posted on 21-9-2004 at 20:15


Quote:

Scaling down an industrial process, I fell like Axehandle

I hope you're more competent than me though, I'm STILL stuck at the sulfur burner. And trust me, you don't want to feel like I :P

Incidentally, we do have one common problem: Your propane stream preheater has almost the exact requirements my SO<SUB>2</SUB> stream preheater has. I'm considering a copper pipe filled with copper wool. Should carry the heat nicely.


[Edited on 2004-9-22 by axehandle]




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[*] posted on 22-9-2004 at 19:30


Quote:
I'm considering a copper pipe filled with copper wool.

At such high temperatures could the copper wool reduce the sulfur dioxide? For my reaction pipe I am considering iron pyrite chunks, they behave as a semi metal, able to conduct heat and they should not react with the sulfur, H2S, and they can only react with the hydrocarbons passing though to make more CS2.

Asking something more on topic, does anyone have a good washing solution that would eliminate most of my hydrogen sulfide so I don't have to burn as much and therefore have to deal with less SO2 gas?

My plans for the whole reaction apparatus are almost finalized, I just have to receive $36 funding to continue :)
BTW: After looking up more information at the library in newer text books, specifically one on rayon manufacture, it turns out the method for producing CS2 from lower weight hydrocarbons reacting with gaseous sulfur has completely displaced the retort method which was just about the only promenade method till around the 1960-70's I think.




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[*] posted on 23-9-2004 at 01:10


It should also be possible to replace the carbon by PE (polyethylene). PE can be used instead of carbon in many reactions I cannot see why it shouldnt work here.

More suitable foe a batch-process though.




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[*] posted on 23-9-2004 at 07:19


You will be the center of attention on the day you make the Carbon Disulfide and H2S! :-))!! Only a trace of the H2S will stink up a whole neighborhood, so expect your neighbors to come sniffing around. I would recommend a windy day. Even if you could burn "all" the H2S into SO2 you will still be able to smell the H2S. BTW, your nose will fatigue of smelling the H2S and you will not be able to smell it long, but it will be just as poisonous, and annoying to your neighbors. Have you checked the solubility of SO2? You might be able to dissolve it into a big container of water, or just draw the SO2 carrying air through a water aspirator, the one used to empty your water bed mattress or aquarium. All that said, it's a very tempting procedure, and I think I'll give it a try... someday.
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[*] posted on 23-9-2004 at 15:10


Mr. Wizard, when I tried my CS2 procedure last week it made corpious quantities of H2S, but I never smelled a bit of it while the reaction was going, although the SO2 smell became over powering from burning my H2S. However when I opened up my apparatus to clean it out, boy oh boy did I small the H2S. And the water was saturated with it (tried to collect my CS2 under water). The SO2 produced turned my galvanized pipes all manner of color, although some of it was just ZnO doing its color change at high temperature thing that it does.

To draw the SO2 though water though would mean an enclosed burning apparatus along with pulling of the gasses though water. That's why I was looking for a way to take care of the H2S before hand, so I just have to find some commonly available salt that forms an insoluble sulfide. Maybe just leading H2S though alkaline permanganate, the permanganate oxidizing it and the alkaline component reacting with the SO2 and keeping it in solution.

Looking though patents today I found some catalysts for this process although it works without them. There are of course forms of silica and alumina that help, but in addition to these two, nickel/aluminum alloys are good catalysts too.

But there are probably many others.




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[*] posted on 23-9-2004 at 15:34


Maybe this is too obvios but can you not absorb most of the H2S in a slurry of sodium hydroxide?
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[*] posted on 23-9-2004 at 15:37


Yes, that was too obvious, I'd be left with a slurry of sodium hydroxide/sodium sulfide right? I didn't know H2S would react with sodium hydroxide appreciable at STP.



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[*] posted on 23-9-2004 at 16:31


I'm under the impression that sodium hydroxide solution will absorb H2S. I'm not sure what level of H2S will remain, almost certainly enough to be a hazard and I'm also unsure if it will go to completion (efficiant use of sodium hydroxide). I'm covering my bases a bit by thinking that Na2S is probably a lot less soluable than sodium hydroxide, so by making a slurry of NaOH beforehand it shouldnt take a high concentration of H2S to start kicking out Na2S as a solid and so long as solid hydroxide remains to replace the hydroxide lost this should continue. You'll know how well this is working if you get suckback. What you make the bubbler and suckback trap out of might require rather more thought. Uncondensed CS2 might also do odd things in the hydroxide solution as well as dissolve the container if its made out of plastic, and if made out of glass, the sodium hydroxide will eat that.

I'm not convinced the exit gas would be breathable (read ventable), even at low flow rates, but I think it should get rid of the molar amounts of gas produced if that makes any sense. It needs to be a very mobile slurry or gas will produce preferential gaps in the mixture and keep them open, it needs to circulate easily (the gas and the slurry).

Ok, now Ive taken a really simple answer and turned it into 2 paragraphs of unjustified waffle most of which youve probably thought of allready but if its going to work, I think this has the best chance.
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[*] posted on 23-9-2004 at 17:04


From the Microscale Gas Chemistry Site hydrogen sulfide in the experiment is neutralized by putting solutions containing it into a 1 M NaOH solution. However a 6 M solution is recommended for disposal of the pure gas. The reaction with 6 M NaOH is supposedly quite exothermic.

H2S is slowly oxidized by water so nearly any oxidizing agent should destroy the compound, if I needed H2SO4 I might experiment with what it does when mixed with H2O2. But like I said, alkaline permanganate should work.

As for using the H2S to precipitate a sulfide, maybe just use CuSO4 as it is widely available.

But the NaOH method sounds pretty good right now, will experiment with it later.

And I agree with you Marvin, these exit gasses will not be breathable/ventable. I just want to get rid of a majority of the H2S before I incinerate them due to the large amount of SO2 generated during my last attempt. So this is just my attempt to cut down on that specific emission.

[Edited on 9/24/2004 by BromicAcid]




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[*] posted on 24-9-2004 at 07:17


Black Copper sulfide, which should form with Copper Sulphate is very insoluble. Would this solution get more acidic as the Copper was removed, forming Sulfuric Acid, and shifting the reaction back the H2S? It forms directly on copper when a soluble sulfide such as calcium poly-sulfide reacts with the copper on a US penny. If you were to make a very dilute NaOH solution and put a copper wire or copper pot scrubbers in it, you should have an effective H2S trap. When all the copper is black, you might have to agitate the copper to expose more metal. Copper sulfate solution would be the easiest, but copper wire is more available.
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[*] posted on 20-10-2004 at 00:37


If you want to use copper sulfate for this, make it dilute, because the high concentration of H2SO4 will eve ntually stop the reaction due to the equilibrium:
CuSO4 + H2S <=> CuS + H2SO4
being shifted to the left. Similarly with copper chloride, acetate should present no such problem. An ammonia complex would work fine also.




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