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Author: Subject: Elemental sulfur from sulfates
Sedit
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[*] posted on 26-7-2009 at 07:25


You do realize if the person finds success they are more then likely going to poison them selfs with Carbon monoxide due to lack of knowledge about what there working with. Be careful Devilinajolie and read more into the process and products before ever attempting this please.




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[*] posted on 26-7-2009 at 08:09


Of course Sedit , danger is always present by the way ..
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[*] posted on 26-7-2009 at 20:27


CO burns in air with a thin blue flame producing harmless CO2 in high yield.



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[*] posted on 26-7-2009 at 23:04


With such high temperatures you must be absolutely sure that no air can enter into the apparatus. The gaseous CO/S mix will burn violently when it comes in contact with air at 1200 C, producing CO2 and SO2. So, if you attempt to do this process you should have some cooling tube in which the S/CO mix is passed. The S will condense on the inside of the tube and the CO will pass through as gas.



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[*] posted on 27-7-2009 at 02:11


I wonder what a controlled amount of air would do. Would CO or S burn, leading to S + CO2, or CO and SO or SO2?

Oh, and don't forget the extraordinarily robust retort which must contain this mix. It should probably be disposable, since you aren't likely to melt all the slag out of it.

Tim




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[*] posted on 1-8-2009 at 11:19


I've heard that carbon dioxide reacts with moist calcium sulfide to yield hydrogen sulfide gas...does this actually work? It makes sense: the CaS would become calcium carbonate. Anyway, it might be easier to convert H2S to elemental sulfur.
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[*] posted on 1-8-2009 at 12:45


Sure. Or HCl if you want it to go faster. H2S is a lot more dangerous though.

It can be burned in a controlled environment, yielding H2O and S (as vapor, I suppose).

Tim




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[*] posted on 1-8-2009 at 18:36


Another way to do it is to get SO2 and H2S to react. One could use an apparatus similar to the one shown below seen in Lehrbuch der anorganischen Chemie by Hugo Erdmann. The explanation of the apparatus: into the bulb (A) first H2S is led in which comes from the side through the tube from apparatus (C) in a moderate stream. The apparatus (B) is a SO2-generator (here SO2 is generated using Hg and H2SO4, Cu should also work in place of Hg). The excess gases are lead through the outlet tubes (D) into a fume hood canal (an old fashioned way of doing it). The reaction between the two begins very soon and in a short amount of time, the inner wall of the bulb is covered with intensely yellow sulfur. I'm pretty sure there are simpler ways of forming sulfur from hydrogen sulfide.

SfromH2S.png - 266kB
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[*] posted on 1-8-2009 at 19:21


Quote: Originally posted by Formatik  
Another way to do it is to get SO2 and H2S to react. One could use an apparatus similar to the one shown below seen in Lehrbuch der anorganischen Chemie by Hugo Erdmann. The explanation of the apparatus: into the bulb (A) first H2S is led in which comes from the side through the tube from apparatus (C) in a moderate stream. The apparatus (B) is a SO2-generator (here SO2 is generated using Hg and H2SO4, Cu should also work in place of Hg). The excess gases are lead through the outlet tubes (D) into a fume hood canal (an old fashioned way of doing it). The reaction between the two begins very soon and in a short amount of time, the inner wall of the bulb is covered with intensely yellow sulfur. I'm pretty sure there are simpler ways of forming sulfur from hydrogen sulfide.


This reaction is moisture catalyzed. This is not a great quality video, but illustrates the reaction: http://www.youtube.com/watch?v=ZCR1HAad4ww

SO2 is probably easier to make by heating a small amount of acid and a bisulfite or metabisulfite for a gas generator.




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[*] posted on 1-8-2009 at 23:31


Always good to see it demonstrated. Thanks. I didn't know it was moisture-catalyzed, but then neither did Erdmann. Though heating the H2SO4 under such a flame as in the set-up was bound to give off some slightly moist SO2 which would have given the impetus for the reaction.
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[*] posted on 26-11-2012 at 12:23


Uberbump! :D
Does anyone know of a chemical way to prepare sulfur, preferably from a sulfate (or even sulfuric acid)?
Where I live, there is absolutely no 'flowers of sulfur', nor sulfur-related garden products (to quote a long-ago post, oh, the capriciousness of chemical availability). We have Epsom salt, calcium sulfate, and sulfuric acid, however.
Also, I'm not sure about a good apparatus for the method involving H2S and SO2, as that has to be carried out at high temperatures, likely involving steel reactors, etc.
You'd think there'd be a good chemical displacement method...




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[*] posted on 26-11-2012 at 13:38


Hydrogen sulphide (danger!) is oxidised by bleach to elemental sulphur. High temperature glowing of sulphate with carbon gives (dirty) sulphide (react with dilute acid to get hydrogen sulphide). Not easy though.

CaSO4 'thermited' with Al powder gives CaS and Al2O3:

CaSO4 + 8/3 Al === > CaS + 4/3 Al2O3. React with dilute acid to get H2S.




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[*] posted on 26-11-2012 at 21:50


Which bubbles out of solution and into a solution of bleach?
What about equivalents with, for example, MgSO4? Or is CaSO4 more favorable?
This seems somewhat easier than the other method, thanks.




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[*] posted on 26-11-2012 at 23:25


Sure hypo, if thats academic interest, but
the need for furnaces hasnt been dismissed, it sounds confortable to keep the things in the retort.

The equations, not balanced (not now)
CaS + CaSO4 --> CaO + SO2

SO2 + O2 --> SO3
SO3 + H2O --> H2SO4
H2SO4 + CaS -> H2S + CaSO4
2 H2S(g) + SO2(g) → 2 H2O(l) + 2 S(s)
My fault: It has to get outta the retort!
CaSO4 is almost trash, no way the precious sand should be wasted with this. furthermore you have to ball mill CaS + CaSO4 for a decent yield.
With a stream of water vapour CaS should yield H2S directly.
Dont mess with the reaction
CaS + 3/2 O2 --> CaO + SO2 its favorable at all temperatures.
CaS + 3/2CO2 -> CaO + SO2 + C will never occur

CaS + 3CO2 -> CaO + SO2 + 3CO favorable past 2200K
CaS + 3CO -> CaO + SO2 + C favorable past around 350K

This last equation if funny, actually.

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[*] posted on 27-11-2012 at 06:27


I suspect that a good approach would be to run the old Leblanc soda process with either sodium or potassium sulphate.

http://en.wikipedia.org/wiki/Leblanc_process

Lixiviation would remove the alkali carbonate leaving you with the calcium sulphide, dilute acid would liberate hydrogen sulphide which would then be oxidised to sulphur.
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[*] posted on 27-11-2012 at 06:54


Quote: Originally posted by ScienceSquirrel  
I suspect that a good approach would be to run the old Leblanc soda process with either sodium or potassium sulphate.
Sodium sulfide is an intermediary in this process, which can directly generate hydrogen sulfide. No need to convert to calcium sulfide first.
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[*] posted on 27-11-2012 at 07:41


Reduction of carbon monoxide back to soot is very unlikely, as evidenced by the accompannying reactions:

16C + 8SO2 --> S8 + 16CO
Which is spontaneous @ 500K
&
8C + 8SO2 --> S8 + 8CO2
which only occurs @ 3000K

So get rid of the furnace idea!
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[*] posted on 27-11-2012 at 08:25


So, a 'thermite' of sorts with sodium / calcium sulfate and carbon, then reacting with a weak acid to generate H2S, and then bubbling through a solution of bleach to generate sulfur... Sounds like a plan.
Could I use a sep funnel to drip dilute acid onto the CaS, or would I just have to mix the stochiometric quantities of the dilute acid and the CaS and bubble whatever bubbles out?
Would charcoal work as a source of pure (>95%) carbon?




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[*] posted on 27-11-2012 at 08:34


Quote: Originally posted by watson.fawkes  
Quote: Originally posted by ScienceSquirrel  
I suspect that a good approach would be to run the old Leblanc soda process with either sodium or potassium sulphate.
Sodium sulfide is an intermediary in this process, which can directly generate hydrogen sulfide. No need to convert to calcium sulfide first.


That sounds like the way to go.
Hopefully it would be possible to mix the raw sodium sulphide with a dilute acid peroxide solution and precipitate the sulphur that way.
I suspect it would not be very neighbour friendly chemistry.
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[*] posted on 27-11-2012 at 08:45


Quote: Originally posted by elementcollector1  
'thermite'

i don't believe the carbothermic reduction of sulfate occurs as readily as the word 'thermite' might suggests; strong heating is required (although probably not as strong as in smelting iron). the good news is that in this kind of process the carbon doesn't have to be very pure (IIRC even flour was used as a source of carbon for reducing phosphates), although you want to keep the hydrogen to a minimum. workup of the incomplete reaction products is necessary anyway.
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[*] posted on 27-11-2012 at 09:24


I read in an old chemistry book that sulphuric acid can be electrolyticly reduced to elemental sulphur or hydrogen sulphide depending on acid concentration, amp density and temperature. Unfortunately, after hundreds of hours of searching, I could not find a similar electrolytic reaction to make elemental phosphorus from phosphoric acid. Can anyone explain why sulphuric acid can be electrolyticly reduced to elemental sulphur, but phosphoric acid cannot be reduced to elemental phosphorus by electrolysis. ?
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[*] posted on 27-11-2012 at 09:26


So, does the reaction maintain its own heat after it starts or is constant external heat required?
What temperature does this reaction favor?




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[*] posted on 27-11-2012 at 09:54


Roasting pyrite, FeS2, liberates free sulfur at lower temperatures than required for roasting sulfates and carbon to produce sulfide. Use a U-shaped sealed retort made out of black iron pipe. Draw a vacuum on it, optionally, to reduce losses and noxious SO2 fumes when opening the retort. Even an L-shaped retort might work, depending on the operating temperature gradients in the retort. You should be able to scrape sulfur crystal off the side of the retort after disassembly. Other sulfides can work, though I don't recall details offhand.
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[*] posted on 27-11-2012 at 10:04


Does anyone know how a lead acid battery sometimes generates hydrogen sulphide?
Electrolytic hydrogen sulphide reacting with electrolytic made sulphuric acid from sulphates, seems an easy route to elemental sulphur from sulphates.
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[*] posted on 27-11-2012 at 10:24


I don't think I have enough pyrite.
Electrolytic hydrogen sulfide... Does this make sense to anyone? Not to diss shannon, but it doesn't seem feasible, I'd need a source and some explanation.




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