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Author: Subject: Elemental sulfur from sulfates
guy_bourgogne
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[*] posted on 29-11-2005 at 21:08
Elemental sulfur from sulfates


Out of curiosity, I have been researching on whether or not elemental sulfur can be easily isolated from a sulfate (e.g., the sulfates of magnesium, calcium, or potassium).

I have not found much, except some vague references to Germany's use of gypsum (calcium sulfate) during WWI. That reference suggests something like reducing CaSO4 to CaS by using carbon and 700 degree C heat. Then, by heating CaS and CaSO4, free sulfur might result. There is no mention of pressures required or anything else. Does anyone know anything about this?

I know sulfur is easily purchased; my question is more academic than practical.
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[*] posted on 29-11-2005 at 21:20


I've heard of the first reaction a few times although I was under the impression that it took place at even higher temperatures, approaching 1100C or so, but I have no distinct recolection on the subject. Actually even my chemistry dictionary lists reduction of calcium sulfate with charcoal to give calcium sulfide.

As for the second reaction you mention, heating calcium sulfate with calcium sulfide... how would this work... I could see some SO2 forming maybe, but that is about it... however I do remember something from a trip to the library when I was looking up things that could be made with calcium sulfate. Anyway, someone else might be able to help you with the second reaction, but there are other ways to sulfur from calcium sulfide then more heating.




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12AX7
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[*] posted on 29-11-2005 at 22:17


Hum...

CaSO4 + 4C = CaS + 4CO
(Some hot, say, yellow temp, perhaps with NaSO4 catalyst, for which I assume Na2S is soluble in, CaSO4 partially soluble, and CaS insoluble, not to mention providing a liquid rather than solid medium for the reduction.)

Then:
3CaSO4 + CaS = 4CaSO3
(CaSO3 is stable and insoluble, IIRC, although in spite of that it probably hydrolyzes a bit.)
CaSO3 + CaS = CaSO3 + S ... um, I don't think this will balance.
On the other hand:
CaS + 4CaSO4 = S + 5CaSO3
...No, it can't ever balance, because all sulfur ions have two bonds, equal and opposite of calcium. For it to balance, something has to displace it. A controlled oxidation could certainly displace it:
2CaS + O2 = 2CaO + S2
(Probably a gaseous displacement, but who knows, it might work at S8 temperatures. Not that I frankly care any about balancing an equation to molecules (O2, S8, P4O10, blah) when it only represents a proportion!)

A similar method is used industrially to process a byproduct of petroleum, H2S. It's burned catalytically (with extra heat input mind you, so it's not a cost-effective method for the producers, only healthier than farting into the atmosphere), so that H2S + O = H2O + S(g) occurs. Sulfur is condensed and sold.

(Off topic, in case you were wondering, you can carry the carbon reduction even further, to CaC2 + CO + S products in total. Though carbide is about as electronegative as sulfur, it'll displace it by evaporating sulfur. CaC2 probably forms in the 2000C+ range, so I'm just mentioning this as a curiosity. ;) )

If I wanted to extract elemental sulfur from sulfates, what I would do is start with sodium sulfate, dehydrate it, reduce to Na2S with C, dissolve and then add an oxidizer such as Cl2, as cost-ineffective as that is. Any mild oxidizing agent (that stops at elemental S) will do it. Say, does anyone know if Na2S solution is oxygen-sensitive? I'd think it'd be a good reducing agent, able to reduce atmospheric oxygen automatically, possibly with the help of lower pH conditions, hum, which would produce H2S...*choke*.

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[*] posted on 30-11-2005 at 01:01


Quote:

If I wanted to extract elemental sulfur from sulfates, what I would do is start with sodium sulfate, dehydrate it, reduce to Na2S with C, dissolve and then add an oxidizer such as Cl2, as cost-ineffective as that is. Any mild oxidizing agent (that stops at elemental S) will do it. Say, does anyone know if Na2S solution is oxygen-sensitive? I'd think it'd be a good reducing agent, able to reduce atmospheric oxygen automatically, possibly with the help of lower pH conditions, hum, which would produce H2S...*choke*.

Solutions of Na2S are quite sensitive to aerial oxidation. When you dissolve Na2S in water, then you get a colorless solution. This quickly turns yellow, due to aerial oxidation. This oxidation gives sulphur, which in turn reacts immediately to form polysulfides with the remaining sulfide. You do not get a precipitate of sulphur.

When such a yellow solution is acidified, then H2S is evolved (even bubbling, if the solution is sufficiently concentrated) and then the liquid also becomes turbid, due to very finely divided sulphur.

I personally do not think this is the best way to prepare sulphur. It is very messy, the result is very impure and the yield is low.




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guy_bourgogne
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[*] posted on 30-11-2005 at 08:39


Thanks for the suggestions.

Here is something from "A Comprehensive Treatise on Inorganic and Theoretical Chemistry". It suggests:

CaSO4 + 3C --> CaS + CO2 + 2CO (@ 700 deg C)

Later on it mentions a Calcium Sufite reaction:

4CaSO3 --> CaS + 3CaSO4 along with
CaS + CaSO4 --> 4CaO + 4SO2 both together at 600 deg C

Suppose I used the first reaction (CaSO4 + 3C --> CaS + CO2 + 2CO) to get the CaS. Then I heated the CaS with some more CaSO4 to get SO2 as described in the second reaction. This eliminates the need for any CaSO3. Does anyone think this might be acheivable with simple home equipment?

Then the SO2 could be dissolved in water, used to make H2SO4. Maybe then we could get sulfur somehow from sulfuric acid.
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[*] posted on 30-11-2005 at 09:00


Pffbt... just find/make FeSO4.7H2O. Dehydrate to FeSO4, then pyrolyze: FeSO4 > FeO + SO3. (Or Fe(III) sulfate.) It decomposes at a low enough temperature that SO3 = SO2 + O doesn't happen much.

CaSO4 can be decomposed directly but you'll need a high refractory, flux-resistant retort to do it in. At those yellow heats, you'll get a lot of SO2.

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[*] posted on 30-11-2005 at 19:28


I had a very old bottle of slightly used photographic fix solution ( a thiosulphate usually) that I found had a lot of elemental sulphur floating about in it (before I threw it out). Does this stimulate any ideas?



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[*] posted on 30-11-2005 at 19:48


Quote:
Originally posted by Twospoons
I had a very old bottle of slightly used photographic fix solution ( a thiosulphate usually) that I found had a lot of elemental sulphur floating about in it (before I threw it out). Does this stimulate any ideas?

I do believe that you could get sulfur from sodium thiosulfate by adding some HCl.
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[*] posted on 1-12-2005 at 10:35


"Pffbt... just find/make FeSO4.7H2O. Dehydrate to FeSO4, then pyrolyze: FeSO4 > FeO + SO3. (Or Fe(III) sulfate.) It decomposes at a low enough temperature that SO3 = SO2 + O doesn't happen much.

CaSO4 can be decomposed directly but you'll need a high refractory, flux-resistant retort to do it in. At those yellow heats, you'll get a lot of SO2.

Tim"
Great- but not what was asked for so ...?
Anyway


Reduce CaSO4 to CaS with C at high temp.
Add acid to get H2S (might as well use H2SO4 and recyle the Ca)

Burn two thirds the H2S to give SO2
React them in solution
2 H2S + SO2 --> 3S + 2H2O

I'm not sure, but I think that
H2S--> H2 +S
at high temps which is even easier.

BTW,
3CaS + CaSO4 --> 4CaO + 4S
seems to balance well enough to me and, at high temp the sulphur boils off forcing the reaction to the right. I'm not too certain about this but it looks plausible that they made sulphur exactly the way they said they did.
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[*] posted on 1-12-2005 at 11:07


I've been told H2O2 will oxidize the H2S (maybe even CaS) to sulfur
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[*] posted on 1-12-2005 at 12:09


Could one prepare a H2S generator, and bubble this though ~30% H2O2? Using a bubbler would help the reaction proceed, if needed. One could do this in a hood or recycle the H2S.

I am not sure how vigerously the reaction would proceed, if needed the concetrations and heat of the H2O2 solution could be varied to suit these needs.




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[*] posted on 2-12-2005 at 03:43


The reduction of sulfates with carbon will probably yield sulfur much more easily if you add SiO2. Because the CaO produced when CaSO4 and CaS conproportionate would react with the SiO2 and evolve energy, the reaction would go faster at a lower temperature and would require less heat. In this way sulfur can be directly produced from the C reduction of Na2SO4:
Na2SO4 + 3C + SiO2 => Na2SiO3 + 3CO + S.




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[*] posted on 2-12-2005 at 09:40


Ah, there 'ya go! Adding an acid like that will make things work much better :)

Same process as phosphorous, notice.

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[*] posted on 2-12-2005 at 15:27


Thanks for sharing that.

Air will also oxidise H2S in solution to sulphur and is easier to get tha H2O2.
IIRC H2S can be oxidised to S by combustion with a limited air supply.
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[*] posted on 6-12-2005 at 20:16


I have heard of using SiO2 to aid the decomposition of CaSO4, although I cannot remember where. It sounds reasonable, and there's plainly plently of SiO2 around.

I also like the idea of using an acid to generate H2S, by which we can exploit the simple reaction of it with SiO2 (it does seem a little dangerous though). What acids would decompose CaSO4 to H2S? I'm guessing the mineral acids (nitric, hydrochloric, and sulfuric) will work, but what about weaker acids like acetic.

Thanks for the suggestions; I found them all very interesting.

Edit: Also, I read recently that Magnesium Sulfate and Carbon, heated together at 600 or 700 deg C, will yield some fraction of elemental sufur. That seems like a pretty easy reaction to do. MgSO4 is not as abundant as CaSO4 (which is in drywall, so it's everywhere), but it's fairly common.

[Edited on 7-12-2005 by guy_bourgogne]

[Edited on 7-12-2005 by guy_bourgogne]
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[*] posted on 7-12-2005 at 08:25


Water is acidic enough to displace (hydrolyze) H2S out of alkaline earth sulfides (CaS, MgS, Al2S3, etc.).

Carbon plus a sulfate is going to work but obviously, it's only ever going to work in low yields. A mixture of charcoal, sulfate and silica (or boric oxide for a lower reaction temperature) under a bed of charcoal (to ensure SOx is reduced) would suffice to distill sulfur.

Tim




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guy_bourgogne
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[*] posted on 7-12-2005 at 10:40


Now that sounds interesting. Will the charcoal, sulfate, and silica reaction work even with CaSO4 at temperatures one could obtain in a simple lab (say 800 deg C or so)? It seems to me that CaSO4 is difficult to decompose.

And should the reaction occur in an oxygen poor environment? or does the CO yielded from burning the charcoal in oxygen cause the reduction of SOx to Sulfur?
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[*] posted on 7-12-2005 at 13:39


It's going to be making CO in the first place anyway. Calculate somewhere around CaSO4 + 3C + SiO2 = CaSiO3 + 3CO + S.

I don't know if CO burns preferentially to S vapor, or if it can reduce SOx. I'm sure gas phase reduction sucks. You have to condense the vapor sans oxygen anyway, so I don't know what you're getting at.

Count on using a high refractory retort (lime is a strong flux), probably composed of a neutral or basic oxide such as Al2O3, MgO, spinel, chromite, etc.

Caveat: if you use an excess of carbon, especially as a cover layer, you'll probably get a lot of CS2 as well.

Tim

[Edited on 12-7-2005 by 12AX7]




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[*] posted on 29-1-2007 at 08:49


"A mixture of charcoal, sulfate and silica (or boric oxide for a lower reaction temperature) under a bed of charcoal (to ensure SOx is reduced) would suffice to distill sulfur."
Good point. The reduction will have to happen with an excess of C, or otherwise the side reaction CaSO4 + C => CaO + SO2 + CO will happen (first CaSO3 forms then decomposes).
CO does reduce SOx, this is an industrial (catalytic) process, however this is probably too slow for amateur uses.




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[*] posted on 26-7-2009 at 00:29


Hi everyone


I have a question about this reaction CaSO4 + 3C + SiO2 = CaSiO3 + 3CO + S.


How can you separate sulfur from the mixture in the end ?
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[*] posted on 26-7-2009 at 00:42


It boils out.. Then you condense it and try not to kill yourself with the CO byproduct..
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[*] posted on 26-7-2009 at 00:42


It is gaseous and bubbles off with the CO.

Tim




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[*] posted on 26-7-2009 at 00:50


Or you might be able to dissolve it in CS2. BTW I think that there are also anaerobic bacteria that reduce sulfates to elemental S, as their source of oxygen; this would be the origin of S deposits in sedimentary rocks.
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[*] posted on 26-7-2009 at 01:03


Thanks a lot , so if S is boiling it can be easy to separate . Whats the exact temperature needed for the reaction between Charcoal , sand and gypse ?
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[*] posted on 26-7-2009 at 05:16


Over 1200C.

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