Sciencemadness Discussion Board - Elemental sulfur from sulfates

Elemental sulfur from sulfates

guy_bourgogne - 29-11-2005 at 21:08

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.

BromicAcid - 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.

12AX7 - 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*.

Tim

woelen - 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.

guy_bourgogne - 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.

12AX7 - 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.

Tim

Twospoons - 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?

Darkblade48 - 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.

unionised - 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.

stygian - 1-12-2005 at 11:07

I've been told H2O2 will oxidize the H2S (maybe even CaS) to sulfur

Swany - 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.

Theoretic - 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.

12AX7 - 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.

Tim

unionised - 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.

guy_bourgogne - 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]

12AX7 - 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

guy_bourgogne - 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?

12AX7 - 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]

Theoretic - 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.

Devilinajolie - 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 ?

497 - 26-7-2009 at 00:42

It boils out.. Then you condense it and try not to kill yourself with the CO byproduct..

12AX7 - 26-7-2009 at 00:42

It is gaseous and bubbles off with the CO.

Tim

JohnWW - 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.

Devilinajolie - 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 ?

12AX7 - 26-7-2009 at 05:16

Over 1200C.

Tim

Sedit - 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.

Devilinajolie - 26-7-2009 at 08:09

Of course Sedit , danger is always present by the way ..

12AX7 - 26-7-2009 at 20:27

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

woelen - 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.

12AX7 - 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

bilcksneatff - 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.

12AX7 - 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

Formatik - 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.

UnintentionalChaos - 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.

Formatik - 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.

elementcollector1 - 26-11-2012 at 12:23

Uberbump!

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...

blogfast25 - 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.

elementcollector1 - 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.

Poppy - 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.

ScienceSquirrel - 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.

watson.fawkes - 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.

platedish29 - 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!

elementcollector1 - 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?

ScienceSquirrel - 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.

tetrahedron - 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.

shannon dove - 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. ?

elementcollector1 - 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?

watson.fawkes - 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.

shannon dove - 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.

elementcollector1 - 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.

shannon dove - 27-11-2012 at 10:35

My explanation is a rotten egg smell coming from a lead acid battery.
but this might be caused by lead sulphate getting hot with reduced lead. Usually when my car battery starts making hydrogen sulphide smell, its going bad.

elementcollector1 - 27-11-2012 at 10:44

Empirical evidence is always a good start.
I couldn't find any evidence of car batteries emitting H2S, but perhaps I haven't been looking hard enough. It would be a useful method, however (although my stomach turns at using my conc. sulfuric acid, when poor chemists in other states don't have any acid).
The only related thing I could find is "sulfation", when stable, crystallized lead sulfate collects in a car battery instead of turning back to lead metal or oxides, leading to the loss of useable lead.

shannon dove - 27-11-2012 at 11:37

I can't find the book right now, but I can assure you that it said that both hydrogen sulphide and elemental sulphur can be made by electrolysis of sulphuric acid. I read it over and over again just to make sure I was understanding correctly. If my memory is right, the sulphur and hydrogen sulphide formed on the anode.

elementcollector1 - 27-11-2012 at 11:44

I'll have to try that when I have the time! Incidentally, what if anything did it say about other sulfates?
Now is a great time for me to figure out a gas-capturing electrode...

chornedsnorkack - 28-11-2012 at 01:47

Quote: Originally posted by shannon dove

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. ?

Because phosphorus is a stronger reducer.
From:
http://www.webelements.com/sulfur/compounds.html
in acid solution, the electrode potential of HSO4/S is not expressly shown, but easily computes as +0,39 V. The electrode potential S/H2S is smaller but also positive at +0,14 V.
From
http://www.webelements.com/phosphorus/compounds.html
in acid solution, the electrode potential of H3PO4/P is not expressly shown either but also easily computes, as -0,412 V. The highest reduction potentials of H3PO4 are actually the expressly shown H3PO4/H3PO3 at -0,276 V and the not expressly shown but easily computed H3PO4/PH3 at -0,281 V.

So. In acid solution, cathode can reduce sulphate first to sulphur and then to hydrogen sulphide.

Under the same conditions, cathode cannot reduce phosphate because it reduces protons first, and ALL forms of phosphorus save phosphoric acid are unstable against reducing hydrogen from acids.

elementcollector1 - 1-12-2012 at 18:02

Can we get it to reduce just to sulfur and not H2S? This would probably require specific conditions, but give the end-product with less reactants involved.

AJKOER - 1-12-2012 at 18:41

Note, the following reactions (see http://books.google.com/books?id=25qJbzc1wMEC&pg=PA83&am... ):

2 HI + H2SO4 --> SO2 (g) + I2 (s) + 2 H2O

6 HI(g) + SO2(g) --> H2S(g) + 3 I2(s) + 2 H2O

So heating an excess of Oxalic acid (or Tartaric or Citric acid) with an excess NaI with CaSO4, for example, may produce H2S which can be reacted aqueous NaOCl (or SO2) to produce Sulfur.

[Edited on 2-12-2012 by AJKOER]

elementcollector1 - 1-12-2012 at 18:51

I liked the electrolysis of sulfuric acid because it required so few reagents, just sulfuric acid, bleach and electricity. It would have been even better to use a sulfate such as magnesium or calcium.
NaI is hard for me to procure, the closest analogue would be KI which I could probably extract from tincture of iodine.
Oxalic acid? I just ran out of that...

shannon dove - 15-12-2012 at 16:32

I have already said this before, but I will talk about it again because it pertains to this. Bacteria can reduce sulfates to hydrogen sulphide. Landfill gas contains a little hydrogen sulphide. Some of it comes from decomposing plant and animal, but some of it comes from bacteria reducing sulphates especially calcium sulphate wall board in construction debris.