Sciencemadness Discussion Board - Carbon Disulfide Preparation

Carbon Disulfide Preparation

BromicAcid - 3-11-2003 at 10:00

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.

Blind Angel - 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.

BromicAcid - 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?

Polverone - 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.

BromicAcid - 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?

S.C. Wack - 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.

garage chemist - 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.

BromicAcid - 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.

Organikum - 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.....

BromicAcid - 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.

BromicAcid - 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.

BromicAcid - 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

JohnWW - 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.

BromicAcid - 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]

axehandle - 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

Incidentally, we do have one common problem: Your propane stream preheater has almost the exact requirements my SO2 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]

BromicAcid - 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.

Organikum - 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.

Mr. Wizard - 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.

BromicAcid - 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.

Marvin - 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?

BromicAcid - 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.

Marvin - 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.

BromicAcid - 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]

Mr. Wizard - 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.

Theoretic - 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.

BromicAcid - 18-12-2004 at 17:37

Quote:

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.

Well, Organikum I've kind of come back to this advice of yours. I was thinking about the lab preparation of H2S by mixing sulfur and paraffin in a test tube with glass wool. By heating the test tube H2S is evolved. This reaction is anologus to the reaction between propane and sulfur to form hydrogen sulfide and carbon disulfide, I actually have to wonder how much CS2 would be evolved from that lab scale prep of H2S. So, I figured that maybe the reaction between paraffin and sulfur would have a higher yield of CS2 if one could get the temperature higher, which would mean making the paraffin go to a higher temperature.

Going beyond paraffin’s which are classified as one web page said as having boiling points between 250 and 350 C and chains of carbons between 24 and 36 carbons in length, you come to asphalt, boiling point in excess of 550C, carbon chains greater then 40 molecules in length. These will allow for higher temperatures and hopefully when mixed with sulfur, a decent yield of CS2, while cutting down a little on H2S production compared to propane. Asphaltum is available readily and cheaply OTC for artistic projects, 400g for $1.25 as one web site sells it. I think the asphaltum will afford a lower reaction temperature then carbon alone and as such make the reaction more feasible. I’ll probably just use the H2S evolved to make some hydrobromic acid or something simultaneously then burn the exit gasses. Still though, this is a long ways off, it’s winter here, and it’s cold.

Re: propane and sulphur

Thomas Winwood - 21-12-2004 at 19:55

Here's the result of a little brainstorming:

The propane/sulphur reaction isn't difficult:
5S8 + 4C3H8 ---> 12CS2 + 16H2S

So what do you do with the hydrogen sulphide? I thought maybe you could oxidise it with something like a permanganate:
8MnO4- + 8H2S -> 8MnO2 + 8H2O + 4O2 + S8

(Stoichiometric quantities of everything would be useful here to make sure you have enough permanganate to oxidise all the H2S and so on.)

The sulphur could, if you set it up right, be passed back through to be reused. (The intricacies of how you pass back the sulphur I haven't considered. It was a thought that only just came to me while writing this post. My initial mental apparatus called for a solution of KMnO4 through which the H2S was bubbled. If you want to collect the sulphur you'll need something more fancy.)

[Edited on 22-12-2004 by Thomas Winwood]

chloric1 - 23-12-2004 at 14:14

Quote:

Originally posted by Polverone
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.

Polverone,
I am interested in your carbonate catalyst statement. I contemplate the mechanism is based on momentary formation of alkaline polysulfides which would contain"active" sulfur atoms stripped of some of their valence electrons. Anybody agree with my thinking? Can anyone dispute it? I am curious for sure.

Some other thoughts concerning preparation: Activated charcoal for fish tanks could be charged with a potassium carbonate solution and then baked at 200 or 300 fahreinheit to drive off water. Then maybe tis charcoal could be loaded in a cast iron pipe in a charcoal fire to red heat and sulfur could be melted and boiled to release vapors through the pipe to a cooling coil and into scrubers.

[Edited on 12/23/2004 by chloric1]

[Edited on 12/23/2004 by chloric1]

BromicAcid - 8-10-2005 at 21:34

CS2 again.... and nothing good to report. I made a metal retort from iron pipes and attached a 24/40 ground glass adaptor into the mess so I could hook it to my distillation setup. I put 5 hollow aluminum rods into the metal pipe to help with heat conduction, they didn't fill it by a long shot but they made a little pentagon in it. 9 grams of airfloat carbon and 48 grams of sulfur along with a pinch of sodium carbonate were well mixed and put into the pipe, not filling it much. This was hooked to a standard distillation setup and the exit gasses were run into water basified heavily with KOH. The reaction mixture was then heated with the aid of a propane torch.

Within five minutes there were whisps of gas appearing in the still head. They entered into the condenser without condensing and the temperature in the still head rose slightly (5 C or so). The gas condensed slightly on the glassware, it may have been sulfur. Some of it made it all the way though the apparatus and passed through the was without being absorbed noticeably, I didn't notice a smell to them either.

Eventually thick yellow/white gas made its way out of the metal retort and into the still head condensing as a solid sulfur. More would come up and travel further down the condenser, and more and more, the condenser had sulfur in it about halfway down and still the still-head was at only about 40C.

About thirty minutes into things, things started to get interesting. Inside the still head I noticed that the sulfur was dripping back down into the flask. I felt the glass and although it was hot it was not near the necessary temperature to melt sulfur. Looking closer I could easily differentiate white drops and yellow drops. The identity of the white drops was confirmed a few minutes later when some began to merge within the condenser, and flow, leaving behind sulfur. CS2 huzzah (see attached picture)! White from dissolved sulfur they cleared up once they started to move and travel down the condenser, appearing faintly yellow.

This went on for some time, with barely any liquid coming over, a drop every 5 minutes or so. Meanwhile I was getting some odd smells, maybe some H2S mixed with some SO2 so I got curious about the washing. I went to my shed and grabbed some powdered copper sulfate. Figuring if I added it to the mixture I would either get black copper sulfate or copper metal or something. I poured the powder in and the soluiton immediately went black. Dark black. Hum.... I added some KOH pellets to the water to keep things basic and pondered that while I watched the insulation burn off my apparatus.

When I got back to the wash bottle I found that it had settled and there were bands of black on top and copper metal shiny on the bottom, neat! The reaction had been going about an hour and the gas evolution had stopped completely despite the torch being on high. So I gave up.

I turned off the heat and allowed to cool for an hour. Opening up the inside of the pipe I found absolutely no sulfur condensed on the inside walls and a small amount on the inside of the 24/40 adaptor, must have been too hot for it. The mixture in the bottom had solidifed and I couldn't remove it. Overall I got maybe 2 - 4 ml of CS2 a really patetic yield. Maybe next time I will use activated charcoal and hope for better results.

sulfurcondenser2.jpg - 75kB

Eclectic - 9-10-2005 at 12:16

Has anyone tried running sulfur vapor into a combustion tube full of yellow-hot charcoal and condensing the result? Or putting the sulfur-charcoal mix in the bottom with a layer of pure charcoal on top and heating from top down? It seems like just heating a mix of charcoal and sulfur will drive the sufur off before you get any significant reaction with the charcoal.

Oh! Nevermind. This was mentioned a while back...

[Edited on 9-10-2005 by Eclectic]

Carbon sulfide - CS2

al - 18-10-2005 at 04:17

Here is a very good method to get CS2.
Berzelius wrote about it .

Distillate wax or paraffine + sulfur in a flask.
CnH(2n+2) + (3n+1) S -----> n CS2+ (n+1) H2S

You get an escape of H2S and CS2
vapors you may condense in water.
CS2 heavier than water has a yellow color, you need to distillate it carefully in order to get the colorless CS2.

Other way to get H2S

I heated 1/2 hour to 700° C a mixture of plaster of Paris an carbon : 200 g plaster for 50 g of wood carbon (powder)
I got a grey compound : CaS. It reacted a little with water giving H2S, but with hydrochlorhydric acid it gave a very strong escape of H2S.
To determine the yield , you may use the calcimeter flask : the weight difference gives H2S , ... so the yield.

BromicAcid - 18-10-2005 at 05:35

Al, see my first post on this page. Heating parafin with sulfur and glass wool is a time tested way to make hydrogen sulfide, but I've never actually heard of CS2 being produced, I figured I was just speculating, like I say in that post though, I don't think the reaction would go to completion as the parafin would volitize away leaving the sulfur eventually, although the glass wool would help. That is why I was thinking about using asphaltum.

I found a nice little review on carbon disulfide manufacture in pdf format, I could have swore it said that there was no catalytic activity of alkali carbonate even though erlier sources said there might be, although that may have been in a different document.

PDF on Carbon Disulfide Production Methods

garage chemist - 18-10-2005 at 11:21

@ al: Can you give us the original report from Berzelius, which contains correct procedure or at least ratio of reactants (if different than stochiometric ratio according to the equation, which suggests 1 mol paraffin and 3 mol sulfur)?
That would be nice.

I'll try out this procedure with paraffin and sulfur. Sure enough, it is normally a procedure to make H2S, but the ratio of sulfur and paraffin seems to be very imporant here.
I think that first the hydrogen of the paraffin completely reacts with the sulfur, leaving finely divided carbon (this is confirmed by the production of a coke deposit from this reaction) to form H2S.
Then excess sulfur (the extra 2 moles) reacts with this special form of carbon to form CS2.

H2S and CS2 could be produced one after the other by first heating gently producing H2S and then stronger to produce CS2.

H2S is absorbed best by bubbling through NaOH solution, forming Na2S which is highly useful.
I tested this with pure H2S (from ZnS and HCl) and the H2S is absorbed completely and very vigorously (be careful of suckback! Use a safety bottle or use a washing bottle with a pipette-tip, this is efficient in preventing suckback as I've found out).
If an excess of H2S is used, NaHS will form, but this can be turned into Na2S by adding an equimolar amount of NaOH.

A very promising alternative to paraffin would be polyethylene. It doesn't boil but melt and react completely with the sulfur.
I have some old PE canisters that I'll cut up and heat with sulfur in the molar ratio 1:3 (PE:sulfur).

ordenblitz - 18-10-2005 at 16:34

Skylighter sells polyethylene powder for $5.35 a pound, in case you didn't want to mess with snipping up containers.

http://www.skylighter.com/

garage chemist - 2-11-2005 at 12:45

Today I tried to make CS2 from polyethylene and sulfur.

I cut up 1,4g (0,1 mol) of HDPE (from an old canister) into very small pieces and added 10g (0,33 mol) of sulfur.
This was put inside a 100ml round- bottom flask.
Upon heating, lots of H2S was first produced, as expected (I worked in my fume hood), and the mixture became black.
Yellow smoke billowed out of the reaction vessel and coated the inside of the apparatus with fine elemental S.
I ran icecold water through the liebig condenser the whole time.
As the H2S production slackened, it was obvious that the rxn vessel now contained some kind of coke and the unreacted sulfur.
No CS2 distilled over, not a single drop, despite heating the vessel until the sulfur boiled.
10 extra grams of sufur were added and heated again, still nothing.

After this experiment I cleaned the flask by boiling some NaOH solution in it, this was very effective (If such a good cleaning method wouldn't exist I wouldn't have even thought about using my precious ground- glass equipment for this!).
A residue of coke was left when all the sulfur had dissolved, it was porous and brittle.

I then tried out paraffin + sulfur (1,4g + 15g), in hope that Berzelius actually improved the process of CS2 production over the charcoal/sulfur process with that.

Again the same, lots of H2S and no CS2, but this time even nastier mess inside the flask due to the paraffin partially evaporating and coating the inside with tar.

Carbon disulfide has so far resisted every attempt at its preparation (I tried charcoal + sulfur before, with and without catalyst, and also heating a test tube with sulfur in the lower part and charcoal in the upper part from the top down- still nothing).

This drives me crazy... it just can't be possible that charcoal + sulfur at over 800°C or methane + sulfur at 900°C are the only methods of preparation!

Maybe the carbon needs to be in a special active form to react to an appreciable degree at 400°C...

My new idea is to heat calcium carbide with sulfur, this should theoretically yield calcium (poly)sulfide and CS2. But since CaC2 doesn't react with oxygen even at 1000°C (only the nitrogen in the air reacts, forming calcium cyanamide), my hopes are low for this. I'll try it anyway.

Acetylene and molten sulfur react to form thiophene, sadly.

Any other unconventional ideas?

Maybe something with hydrogen persulphide... this is a strong oxidizer like H2O2 (it also bleaches hair and tissue like H2O2), but with sulfur instead of oxygen.
Its preparation is easy, boil Na2S solution with sulfur to form polysulfide, then drip this solution into HCl, the hydrogen persulphide separates as an oil, care must be taken as it can spontaneously and violently decompose in the presence of bases or other catalysts.
Maybe this is able to oxidise charcoal or some organic compound to CS2...

In Hollemann- Wiberg, hidden in some other article, I found the reference that an early experimenter accidentally prepared "sulfur alcohol", as he called it (it was actually CS2) by heating pyrite (FeS2) with charcoal.

Now, where can I get pyrite? Can I prepare it by reacting sulfur with FeS or directly two moles of S with one mole of Fe?
Does anyone know if it is possible to prepare artificial pyrite or where to get the mineral?

[Edited on 2-11-2005 by garage chemist]

12AX7 - 2-11-2005 at 15:48

Doesn't particularly suprise me... what temp does carbon burn with oxygen?

Just use a heated, sulfur-resistant tube loaded with charcoal. Pass S vapor through it and condense yon CS2.

Pyrite? That would be interesting. Is carbon electronegative enough with respect to iron to remove sulfur, or is only the extra S reduced? (2FeS2 + C = CS2 + 2FeS, which is stoichiometric.) You can make it by direct synthesis, yes. I don't know of any source of the mineral, but eBay might be worth a try (unlikely since the only specimens would be top quality masses of crystals, but who knows, maybe someone has it in bulk).

Tim

[Edited on 11-2-2005 by 12AX7]

garage chemist - 5-11-2005 at 07:16

Today I went to the university chemistry library and looked up, among other things, the production of CS2.

My reference was the Gmelin Handbook of Inorganic Chemistry. With about 5000 Volumes it contains nearly the entire chemical knowledge of humanity on inorganic chemistry.
There was a 250- page volume dedicated exclusively to compounds of carbon with sulfur!

Covered were the different industrial processes to effect the reaction of carbon with gaseous sulfur, methods employing electrical heating by current flowing through the charcoal bed were also described. Same for the various processes employing hydrocarbons and sulfur vapor (all requiring superheated sulfur vapor above 600°C, sadly).

The reaction of charcoal with pyrite was also described, and nearly quantitative yield of CS2 was stated!
The reaction goes to completion, producing elemental iron, which means that FeS made by the well- known reaction between iron filings can also be used.

FeS2 + C ----> Fe + CS2

No temperatures were stated though. This could be the biggest problem, since the high temperature required is also the main problem in phosphorus production.

For the issue of producing high temperatures, I have found that a propane- oxygen torch is capable of softening quartz, more than just a tiny bit.

If a chemical mixture (like for example bone ash + charcoal + SiO2 !) has to be heated to over 1000°C, I would put it in a quartz test tube (available from Ebay for less than 10$)
and heat it with a propane- oxygen torch.
Bright yellow heat is very easy to achieve.

[Edited on 5-11-2005 by garage chemist]

chloric1 - 5-11-2005 at 12:38

Mind the ignition point. CS2 ignites at about 100C. If there is a leak in your aparatus you could have a flash explosion.

garage chemist - 6-11-2005 at 09:56

5,6g Fe filings and 3,2g sulfur were mixed and ignited to give FeS. The identity of the product was checked by adding some of it to HCl, H2S stink was produced at once.

The FeS was powdered, mixed with 1,2g charcoal powder and put in a Duran test tube.
It was heated with a bunsen burner.
Soon the evolution of gas was apparent and the powder even was thrown around a bit.
Colorless drops of liquid condensed in the upper part of the test tube.
The liquid was soaked up with a strip of blotting paper and tested for flammability.

It didn't burn.

It was just moisture from the charcoal.

I heated the test tube with a propane- oxygen torch until it was glowing yellow and the glass got a blister because it started to melt.
NOTHING.

garage chemist - 6-11-2005 at 11:05

OK, I did some patent research and can now see the superiority of the sulfur+propane process when done right.
The temperature of boiling sulfur is high enough to effect the reaction with propane.
Bubble propane through boiling sulfur.

The problem is that the CS2 can not be condensed completely from the exit gas. Too much H2S is present, the partial pressure of CS2 is too low to allow its complete condensation.

The unreacted sulfur first needs to be removed from the gas stream.
Due to the tendency of sulfur to not condense on the walls of a condenser, but rather as a fine yellow smoke of solid sulfur, I would use a gas washing bottle stuffed with glass wool for the removal of unreacted S.

Then a condenser in which most of the CS2 condenses.

The next stage, the CS2 absorption unit, would consist of a gas washing bottle filled with heptane or just simple petroleum.
CS2 dissolves into it and can later be recovered by distillation.
Choose a petroleum type of high enough boiling point in order to facilitate the separation. If necessary distill your petroleum before use and discard everything coming over below 100°C.
And don't use too little petroleum, fill your washing bottle up to its designed capacity, the absorption capacity of petroleum for CS2 is limited.

Then the most important step, absorption of H2S.
Alkali is the way to go. A washing bottle with NaOH solution would do the job nicely, I speak from experience.
However, the problem is suckback. H2S is absorbed rather violently.
Maybe a column with soild NaOH (or Ca(OH)2 if it works) is better here since it eliminates this risk.

The whole apparatus will need to be located under a good fume hood and operated after midnight, when nobody is outside, due to the inevitable smell evolution. Getting a H2S apparatus smelltight is next to impossible, I know that by experience. Silicon and PVC hoses hold up great, but diffusion of H2S through the walls is very noticeable.

[Edited on 6-11-2005 by garage chemist]

neutrino - 7-11-2005 at 14:12

You have some interesting ideas there. A few thoughts about this process:

Is CS2 stable in basic conditions? If it is, wouldn’t it be easier to remove the hydrogen sulfide in the gas stream and then condensing the CS2?

If you want a high bp petroleum, try mineral oil. It is available pure at any pharmacy and has a bp well over 250*C

garage chemist - 7-11-2005 at 14:27

CS2 reacts instantly with alkali sulfide solutions (such as that produced by reaction of the H2S and NaOH), producing a solution of trithiocarbonate. So better first condense and wash out the CS2 and then dispose of the H2S.
Another solution to the H2S issue would be burning of the H2S, e.g. feeding into the air input of the bunsen burner. The SO2 creates far less stink than the H2S, which even I with my tolerant neighbors cannot blow out of my fume hood in such amounts.

I have some doubts if propane and sulfur will actually react rapidly enough and in the desired manner. Propane is said to give problems due to coke formation.

Anyway, today I went to the pharmacy and asked if they could order 250ml of carbon disulfide for me. The short conversation was interesting. Full report here when they have called and informed me if they will sell it to me.

garage chemist - 8-11-2005 at 10:28

From "Mercks Warenlexikon", it describes the construction of a small- scale carbon disulfide furnace fired with charcoal:

Schwefelkohlenstoff (Schwefelalkohol, Kohlenstoffbisulfid, Sulfokohlensäure, lat. alkohol sulfuris, carboneum sulphuratum; frz. sulfure de carbone; engl. sulfuret of carbon). Aus zwei festen Körpern, Schwefel und Kohle, entsteht, wenn sie durch Glühhitze zur Verbindung gezwungen werden, der S., oder wie er von seinem ehemaligen Entdecker, Lampadius in Freiberg, genannt wurde, Schwefelalkohol. Bei Bereitung desselben kommt es darauf an, Schwefel in Dampfform durch glühende Kohlen streichen zu lassen. Die Apparate hierfür haben eine, in einem Ofen stehende, von unten zu beheizende Retorte von Thon oder Eisen, im letzteren Falle innen mit Thon ausgekleidet. Sie wird mit Holzkohle in haselnußgroßen Stückchen gefüllt und angefeuert, bis die Kohlen hell glühen. Es wird nun nach Bedarf ganzer Schwefel eingeworfen durch ein Rohr, das von außen in die Retorte und bis nahe an deren Boden führt und immer rasch wieder geschlossen wird. Der unter den Kohlen schmelzende und verdampfende Schwefel durchzieht dieselben, sättigt sich mit Kohlenstoff und die neue Verbindung zieht oben dampfförmig durch ein Knierohr ab, das durch eine Kühlvorrichtung geht. In einem zwischengelegten Gefäße schlägt sich erst der unverbunden mit fortgegangene Schwefel nieder; das übrige geht weiter, verdichtet sich tropfbar und sammelt sich in einer Vorlage unter Wasser, denn der S. ist schwerer als dieses und mischt sich nicht mit Wasser. (...)

I'll translate it when I have time.

I can also get a drawing of the apparatus, it's in my school and I have to copy it.

[Edited on 8-11-2005 by garage chemist]

chloric1 - 8-11-2005 at 18:00

Quote:

Originally posted by garage chemist

Anyway, today I went to the pharmacy and asked if they could order 250ml of carbon disulfide for me. The short conversation was interesting. Full report here when they have called and informed me if they will sell it to me.

I would not be surprised if they sent the feds to your place. But maybe Germany is more friendly to intellectuals than the USA? Curious to see if you get any static at all.

chemoleo - 8-11-2005 at 20:11

Quote:

CS2 reacts instantly with alkali sulfide solutions (such as that produced by reaction of the H2S and NaOH), producing a solution of trithiocarbonate.

Well, shouldn't it then be possible to revert this reaction, by i.e. adding a strong acid to get back CS2? Similar to getting CO2 from Na2CO3 and acid?

Maybe you ought to order trithiocarbonate from the pharmacy instead.

What exactly did you tell them to convince them to sell CS2?

I remember once purchasing 1 l of 80% phenol from the pharmacy. It took a bit of convincing, but I said I wanted to make bakelite. The pharmacist loved me for my enthusiasm!

[Edited on 9-11-2005 by chemoleo]

chloric1 - 8-11-2005 at 20:21

Yes, I can see already things are different. Most "pharmacists" I have meet in the US of A actually hate chemistry or are indifferent to it. It seems they are in it only for the money. Some of the prissy girls would instantly be suspicious if I asked for anything like that.

Don't get me wrong a few have a passion for the central science. But most just want the green.

garage chemist - 9-11-2005 at 10:02

Hmm, no luck here, despite the pharmacist being willing to order some CS2 for me.

When I first asked for CS2, the immediate first answer was: "We don't sell chemicals because we don't know what you are going to do with them! For what do you need carbon disulfide?"
"As a solvent for sulfur and phosphorus", I answered.
"And then?"
"I am a hobby chemist!", I said.
She said that she will phone her supplier and ask what this (250ml) will cost.

The next day I was told on the phone that the minimum amount I could buy was 1 Liter, and that it would cost 99€.
Of course I was not willing to pay this astronomic price, I opened Mercks CD catalogue and printed out the product site for 1 Liter Carbon disulfide, the price being 35,75€.
With this I went into the pharmacy again, she said "That's very interesting!", because the supplier who got his chemicals from Merck was the one who charged that ridiculous price.
Today there was no phone call, I called the pharmacy myself and was told that another supplier only sells at least 3 Liter of CS2, the price for it being 90€ (plus tax, plus shipping and plus the money the pharmacy wants to make, mind you!).
The pharmacist faxed Merck directly and asked if CS2 is available in amounts of less than one liter, and in lower purity grade than "Extra Pure" (technical grade would suffice for me). The answer to this fax has not arrived yet.
I was also given the adress of a local chemicals redistributor, but they were closed already because it was too late today.
Tomorrow I won't have time, sadly...

So I'll have to wait. But atleast I know that I would have got my CS2 if there had been a supplier who sold small amounts.

That is the actual problem today: the big companies refuse to sell in small amounts.
"Too much hassle for too little profit", they think.

My only hope now is this chemical redistributor.

A few tips when buying chemicals from pharmacies:

The most important thing is to tell them that you are a hobby chemist.
Most of the chemicals we want have absolutely no use for the average citizen, and the pharmacists know that.
They also can NOT, I stress this, send the police to your house. The worst they can do is refuse your order. And there are a lot of pharmacies in cities. New pharmacy, new luck.
It is of course very unwise to order listed drug precursors, this could actually get you into trouble.
But CS2, for example, is no drug precursor, and also no explosives precursor. The only problem is its toxicity and flammability.

Also, know the price that big chemical companies charge for your wanted substances (search in their online catalogue), because the pharmacists don't search for the cheapest seller. This is important to avoid paying ridiculous prices.

If you don't try to get your substances in the pharmacy, you can't complain about chemicals being hard to get.
It is one thing to save some money or avoid some hassle by using OTC chemicals or synthesizing chemicals yourself, but another thing to spend months trying to make a stupid SOLVENT. It will eventually be much more expensive than ordering it, even if the price may seem high, and much more work.
The example above of the pharmacy not being able to order the desired small amount is actually an exception, I got quite a bit of chemicals in the pharmacy before (it was another pharmacy, the one in which I asked for CS2 had never seen me before).

garage chemist - 24-12-2005 at 02:44

Today, on december 24th, seven weeks after my first inquiry about the carbon disulfide and after numerous other calls to chemical suppliers and visits to the pharmacy, the pharmacy phoned to tell me that my carbon disulfide had arrived.

I went there immediately, paid the price and thanked them for their cooperation.
The long delay was the fault of the supplier who didn't have the CS2 in stock and had to order it first.

The 1 L bottle of "Carbon disulfide, extra pure" from Schuchardt (re-sold by Köhler) has now found a place on my chemical storage shelf.

What a wonderful christmas present!

neutrino - 24-12-2005 at 06:24

Very nice. What was the price of this bottle?

garage chemist - 24-12-2005 at 08:50

The list price was ca. 36€, but after adding taxes, enormous shipping fees and the amount the pharmacy wanted to make the price was at 56€.
It was still OK since I am going to give some of this to other hobby chemists in exchange for other chemicals. The entire liter would be too much for me.

garage chemist - 24-12-2005 at 16:23

I did a few simple experiments with the CS2.

When added to water, it forms a separate layer below the water without any noticeable dissolution.

I added some CS2 to some sulfur in a test tube, swirled it and the sulfur dissolved in a matter of seconds! Surprisingly large amounts of sulfur could be dissolved in the CS2.
When poured into a shallow dish, the CS2 evaporated, leaving behind some more or less clearly rhombic sulfur crystals.

I poured some of the CS2 into a small beaker, held a metal wire into a bunsen flame for a short period and inserted it into the beaker (it didn't glow anymore). Whomp! made the beaker and the CS2 burned with a bright blue flame, giving off choking clouds of SO2.

My carbon disulfide has a pleasant aromatic smell, not unpleasant at all. It reminds me of butane for lighters (the kind that doesnt contain odorants).
I smelled it only briefly, since there's a "toxic" warning label on the bottle and the toxicity on inhalation is stressed in the warnings. I also conducted the experiments in my fume hood.

Fleaker - 25-12-2005 at 17:30

Odd, I've always remembered it and described it as having a very unsavory odor that some consider the smell of rotten cabbage, but you say it had a pleasant aromatic smell?

Merry Christmas by the way.

neutrino - 25-12-2005 at 19:08

Is it possible that you had some thiol contamination in your CS2? Thiols are generally considered to be very smelly, for example CH3CH2SH is listed in the Guinness book as the worst smell in existence.

Fleaker - 27-12-2005 at 10:29

That's a definite and probable possibility although it was ACS reagent grade and should not have had such impurities. Well if that is so, then garage chemist can be proud to have made a better product than Fisher!

Endimion17 - 30-5-2012 at 07:36

I'm briefly ressurecting this thread as it's the one with the most suitable title.
Check this out.

It's about hydrogen sulphide production methods, but there's carbon disulphide, too.

Methane. Cheap and easily available in some areas as natural gas in pipes. The temperature of the synthesis does not pose an unsolvable problem. Has anyone tried this?

[Edited on 30-5-2012 by Endimion17]

DJF90 - 30-5-2012 at 07:47

If you look on Lambdasyn (and possibly Versuchschemie too!) you'll see a preparation for schwefelkohlenstoff, a.k.a. carbon disulfide. These guys are doing it by passing acetylene gas through molten sulfur at about 320-380*C if memory is good. Calcium carbide is cheap and readily available if you dont have an acetylene cylinder, and this method has the significant advantage that it can be performed in borosilicate glassware (600-650*C will be too high for this!!).

Links: http://www.lambdasyn.org/synfiles/schwefelkohlenstoff.htm
http://www.versuchschemie.de/topic,12871,-Schwefelkohlenstof...

p.s. the latter is by our very own GarageChemist! And the former may be also...

[Edited on 30-5-2012 by DJF90]

Endimion17 - 30-5-2012 at 08:12

Calcium carbide unfortunatelly isn't readily available, and neither is cheap. It used to be available just about everywhere because it was used for welding, but today it's quite hard to find. The more one country is developed, the less chance is that CaC2 is lying around cheap. It's simply not used extensively as in the past.

I found the info on methane in one old book, so I placed it here for the sake of accumulation of information.
Regarding the borosilicate glass, there are plenty reactions done in ceramic tubes and metallic pots. Glassware is for low temperatures, and ~350 °C is about the highest temperature such glassware was supposed to normally endure without problems.

garage chemist - 30-5-2012 at 10:52

Yes, I did this synthesis together with a friend some time ago.
The idea came from Sauron in this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=12267
The process that Sauron posted works exactly as advertised.
I found that the temperature must be very close to the boiling point of sulfur for the reaction to occur. 430-440°C were necessary for all the acetylene to react. At less than 420°C, much acetylene passes through the sulfur unchanged. The offgas must be checked often for its acetylene content.
A thermocouple temperature sensor inside a glass tube must be used inside the reaction, otherwise there's no control over the temperature, and yield may drop to zero because the reaction is just a little too cold.
Protective gas must be used to flush all air out of the apparatus before beginning to heat the sulfur, because otherwise the CS2 vapor will autoignite and explode, possibly breaking the reaction flask and scattering molten burning sulfur. I used propane as protective gas because I had nothing else back then. It was quite a dangerous synthesis and I wouldn't like to do it again.

ScienceSquirrel - 30-5-2012 at 11:43

Quote: Originally posted by Endimion17

Calcium carbide unfortunatelly isn't readily available, and neither is cheap. It used to be available just about everywhere because it was used for welding, but today it's quite hard to find. The more one country is developed, the less chance is that CaC2 is lying around cheap. It's simply not used extensively as in the past.

I found the info on methane in one old book, so I placed it here for the sake of accumulation of information.
Regarding the borosilicate glass, there are plenty reactions done in ceramic tubes and metallic pots. Glassware is for low temperatures, and ~350 °C is about the highest temperature such glassware was supposed to normally endure without problems.

The most common use of calcium carbide was in lamps. Bicycle, motocycle, car and hand lamps all used carbide as a fuel source.
Even in the seventies you could still buy it in 1lb tins in cycle shops in Britain.
The best place to get it today is in caving shops.

Endimion17 - 30-5-2012 at 11:59

Quote: Originally posted by ScienceSquirrel

We still often use carbide for caving purposes because it liberates more heat, which can come very handy in the cold depths, but there aren't specialized shops. People have their "dealers".

I know only one store that sells it legally, and that's one pyrotechnic store which sells it for a lot of money. The purpose of it is to create detonations in barrels for fun. Wikipedia mentions carbide shooting in Netherlands, as if it's their national phenomenon, but it's something people do pretty much everywhere in Europe, or at least they were doing it when carbide was cheap and readily available.

ScienceSquirrel - 30-5-2012 at 12:23

There was at least one shop in the UK that used to sell it to cavers until quite recently. I tried to buy some from them but they would only ship to the UK mainland.
Maybe it has since slipped further off the market.
I bought a jar some time ago from eBay and I only use a few grams at a time so it may outlast me.

ScienceSquirrel - 30-5-2012 at 13:16

These people will sell it to you in the mainland UK;

http://www.caving-supplies.co.uk/cgi-bin/psProdDet.cgi/26012...

BromicAcid - 30-5-2012 at 14:10

Endimion17 the first page details some of my first research on using low weight hydrocarbons as a carbon source, what I had found was propane was used industrially. Due to the ease of dispensing propane at a lab scale vs. methane (without special hooksups) it called to me however it never got past the drawing board due to the large volume of H2S that would be produced.