Sciencemadness Discussion Board

Elemental sulfur from sulfates?

Junk_Enginerd - 8-9-2019 at 00:00

Hey

Sulfur is starting to become quite a chore to get a hold of where I'm at. It's not impossible yet though, so note I don't need suggestions on how to find it.

Mostly out of curiosity, I was thinking of ways it could be synthesized from what I do have. My main sulfur containing cheap bulk chemicals would be iron sulfate and gypsum.

I've read around a little here and done some random googling and found a couple of ways from FeS2 to sulfur, but none from sulfates...

There's got to be a route for this right?

I suppose I could just outright calcine the iron sulfate to SOx and pipe it through hot activated charcoal to reduce it, but that'll be labourous and messy as fuck. I don't know how I'd separate charcoal from sulfur either. Maybe melting the sulfur... What other ways might exist?

An electrochemical process would be neat for me, if that's a way. I've got half decent equipment in that area, but not the chemical know how to figure out how to use it here.

Deathunter88 - 8-9-2019 at 00:16

Well to separate sulfur from impurities a great way is to dissolve it in toluene or xylene. Not sure it would be practical to calcinate large amounts of most sulfate salts though.

Junk_Enginerd - 8-9-2019 at 03:29

Quote: Originally posted by Deathunter88  
Well to separate sulfur from impurities a great way is to dissolve it in toluene or xylene. Not sure it would be practical to calcinate large amounts of most sulfate salts though.


Allright, good tip. I have xylene so that works.

wg48temp9 - 8-9-2019 at 03:59

Calcium sulfide can be produced by the carbon reduction of the sulphate I guess that iron sulfide could be produced by the same process.

Sulfides can be used to produce hydrogen sulfide which can be burnt in a limited supply of air or oxidized in solution to produce sulfur.

Sulfur can be separated from carbon very easily as carbon is not volatile at any reasonable temperature while sulfur boils at about 450C and also sublimes at low temperature. meaning it can be distilled or sublimated from the carbon.

Alkoholvergiftung - 8-9-2019 at 04:33

You dont Need to burn the H2S Gas you only Need it to bubble througt MgSO4 solution or FeSO4 or over moist Gypsum. You get Metall Sulfid and lot of Sulfur.

unionised - 8-9-2019 at 10:52

Quote: Originally posted by Alkoholvergiftung  
You dont Need to burn the H2S Gas you only Need it to bubble througt MgSO4 solution or FeSO4 or over moist Gypsum. You get Metall Sulfid and lot of Sulfur.

Really?
How?

AJKOER - 8-9-2019 at 13:40

The action of highly electropositive zinc metal with H2SO4 does apparently create some H2S gas! Source: See for example, https://chemiday.com/en/reaction/3-1-0-985 . See also comments at https://books.google.com/books?id=Dv_F03cdKPUC&pg=PA429&... .

Shaking an aqueous solution of H2S with H2O2, or some chlorine gas, or add a small amount of hypochlorous acid, or even chlorine bleach (but not in excess), will liberate elemental sulfur. Reactions:

H2O2 + H2S --> 2 H2O + S (s)

Cl2 (g) + H2O = HCl + HOCl

H2S + HOCl --> S (s) + H2O + HCl

H2S + NaOCl (aq) --> S (s) + H2O + NaCl (aq)

But, do avoid an excess of, for example, HOCl owing to:

S + 2 HOCl --> SO2 + 2 HCl

And, in further excess, even back to sulfate!

[Edited on 8-9-2019 by AJKOER]

teodor - 9-9-2019 at 02:49

Quote: Originally posted by AJKOER  
The action of highly electropositive zinc metal with H2SO4 does apparently create some H2S gas!


Thank you AJOKER for this information! Because I always thought the smell of H2S from this reaction is because of some impurities, like with Fe. Also the "Inorganic Reactions in Water" contains another interesting information.

Bedlasky - 9-9-2019 at 03:51

Quote: Originally posted by teodor  
Quote: Originally posted by AJKOER  
The action of highly electropositive zinc metal with H2SO4 does apparently create some H2S gas!


Thank you AJOKER for this information! Because I always thought the smell of H2S from this reaction is because of some impurities, like with Fe. Also the "Inorganic Reactions in Water" contains another interesting information.


I also noticed smell of H2S when I reacted zinc and H2SO4 and it surprised me.

Quote: Originally posted by AJKOER  
The action of highly electropositive zinc metal with H2SO4 does apparently create some H2S gas! Source: See for example, https://chemiday.com/en/reaction/3-1-0-985 . See also comments at https://books.google.com/books?id=Dv_F03cdKPUC&pg=PA429&... .


In this book is answer to my recent question about reaction of thiocyanate with zinc/nascent hydrogen. Thank you very much!

[Edited on 9-9-2019 by Bedlasky]

teodor - 9-9-2019 at 04:20

So, is it possible to reduce, say CaSO4 by Zn and heating? Because there is another thread about converting CaSO4 to something useful and couple of new threads about getting CaS. It seams that Zn could be an options for all of them. Considering the melting point of Zinc it could be much better route than reducing with carbon.

[Edited on 9-9-2019 by teodor]

Alkoholvergiftung - 9-9-2019 at 05:31

unionised
They used the waste Sulfides from the leblace Prozesse for the H2S Generation.They lit it throught pipes filled with Gyps at red hot temperature
CaSO4 + 4H2S = CaS + 4H2O + 4S.

If you use Ironsulfate it will go at roome temperature.

Weldon used an suspention of Ironoxyde or Manganoxyde and bubbled H2S and air inside.
Calciumsulfide was also made by mixing gyps and charcoal and formed briks out of it and burned it like clay briks. 1g gyps on 0,15 if it had crystall water in it and 0,2 without.

MrHomeScientist - 9-9-2019 at 06:50

CaSO4 and aluminum is used as a heat booster in some thermite reactions (I did this for silicon thermite). That reaction produces calcium sulfide, which reacts with water to release copious amounts of H2S. So zinc would probably also work, if you can find powdered Zn.

Boffis - 9-9-2019 at 07:53

Historically barium hydroxide was produced by the "Black ash process" where barium sulphate was reduced with coke dust and the resulting crude barium sulphide leached with hot water. Air was then drawn through the hot, filtered barium sulphide solution. Oxidation of the sulphide solution gave the desired barium hydroxide solution and a precipitate of sulphur. If carried out close to boiling point the sulphur clumps together and was easily removed by filtration. I see no reason why you couldn't do this with baked plaster of paris and charcoal instead of baryte and coke.

Bedlasky - 9-9-2019 at 08:15

Quote: Originally posted by Alkoholvergiftung  

If you use Ironsulfate it will go at roome temperature.


At room temeprature ferrous sulfate react with H2S only at basic conditions to produce FeS.

[Edited on 9-9-2019 by Bedlasky]

hodges - 9-9-2019 at 16:17

Not from sulfates, but from sulfites (including the readily available "hypo", or sodium thiosulfate). This might work, but I have not seen it described before.

I remember once long ago someone challenged me to prove that carbonates have carbon in them. So I reacted some NaHCO3 with a dilute acid to produce CO2. Then I lowered a burning magnesium ribbon into the CO2. The magnesium reduced the CO2 to C, and I was left with carbon (along with magnesium oxide).

So maybe the same thing would work with SO2, which could be produced by treating sulfides (as well as the readily available "hypo" - sodium thiosulfate) with a dilute acid.

Equation would be
SO2 + 2Mg -> S + 2 MgO

The enthalpy of formation of CO2 is -394 KJ/mol, whereas for SO2 it is -297 KJ/mol. So I would guess that the SO2 would be easier to reduce than CO2. But I don't have the chemicals handy to test this myself at this time.

Certainly, the sulfur produced would be vaporized. It would likely sublime onto the walls of the container of CO2.

Probably not real practical, but the original author did mention curiosity, so thought I would chime in with this idea.


Bedlasky - 9-9-2019 at 23:16

Quote: Originally posted by hodges  
So maybe the same thing would work with SO2, which could be produced by treating sulfides (as well as the readily available "hypo" - sodium thiosulfate) with a dilute acid.


If you acidified sodium thiosulfate you'll get sulfur in solution.

S2O32-+2H+--->SO2+S+H2O

But reaction is more complex. Sulfates, polythionates and hydrogen sulfide are also formed in this reaction.

teodor - 10-9-2019 at 02:09

I am curious, is it possible to disproportionate SO42- and OH- somehow to get S2O32-? Because we have reaction like this

4S+8NaOH = Na2SO4+3Na2S+4H2O

(dry process)

and the equilibrium possible is driven by evaporation of H2O. So, if we start with Na2SO4 and Na2S (or NaOH) and H2O can we get some oxysulfate anion (between SO42- and S, like S2O32-)?

Alkoholvergiftung - 10-9-2019 at 06:44

hodges
I ve read an early production way was to bubble H2S throught H2SO3 yields immedatly fine Sulfur powder.

Bedlasky Ive searched again and it was Iron III sulfate with the Sulfid Sulfur reaction.

teodor - 10-9-2019 at 06:50

Actually there are sooo many ways to get a sulfur powder from H2S, actually when you try to get H2S sometimes you get just a sulfur powder ... Many oxidizers will work on H2S but I never saw H2S as something affordable that I can use if I have no access to sulfur.

So, I propose routes in which you mix (poly)sulfide with some oxidizer and acid and you get S directly without bothering with H2S which requires either descent equipment or huge optimism to work with.

draculic acid69 - 11-9-2019 at 00:47

Trying to get elemental sulfur from sulfates sounds like a long way around to get something you can get at any hardware store.in my opinion it's a bad starting point.
And anything utilizing h2s is soooooo dangerous.so unnecessarily dangerous

teodor - 11-9-2019 at 03:16

Oh, I forgot about the most obvious route to get sulfur from sulfates. From wikipedia:

"Native sulfur is synthesised by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes."

And I think with a proper culture of bacteria making sulfur at home should be not more complex than making alcohol or bread.

Also, if you like to travel: "Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire; such volcanic deposits are currently mined in Indonesia, Chile, and Japan."

So, if it will disappear from your hardware store one day I believe you know what to do :)

[Edited on 11-9-2019 by teodor]

Alkoholvergiftung - 11-9-2019 at 06:36

Ive read that the chemical factory Griesheim made Soda via Sodiumsulfate and straw and piss.The waste product was elemental sulfur If you want to go the bacterial route. :)

[Edited on 11-9-2019 by Alkoholvergiftung]

walruslover69 - 11-9-2019 at 08:57

I think sticking with a SO2 method instead of H2S is the safer and better way to do. You could generate SO2 rather consistently by hot sulfuric acid and copper on a regular hot place. Then reduce the SO2 over carbon, zinc or magnesium. sulfur should be easy to extract using xylene or toluene. It does sound laborious though.

draculic acid69 - 11-9-2019 at 17:38

If you can find the actual bacteria that they use that would be the best way but how hard is it to get?is it even available to the public? A lot of companies use strains of bacteria they've spent many years breeding and modifying.

Sulaiman - 11-9-2019 at 17:56

Sodium thiosulfate plus hydrochloric acid is a simple method to produce sulphur,
but
sulphur should be available in massive quantities in any country that refines crude oil.

rockyit98 - 11-9-2019 at 21:13

sulfur is dirt cheap like 50 Cents per pound.Do your local market a visit.

AlbertaSulfurAtVancouverBC.jpg - 205kB

Tsjerk - 11-9-2019 at 23:21

You don't want to work with sulfate reducing bacteria/archaea... They are facultative anaerobic. If you open the jar or whatever you are growing them in, they die.

teodor - 12-9-2019 at 00:02

Is it the same as making alcohol on vinegar depending on the conditions?

Really I will not use them because I still have access to sulfur and probably should bother more about my health to live long enough until the day it will disappear :) but, just to discuss something interesting, how do you think, can we use the same kind of bacteria to make, say, Se or some other elements we unable to buy . A good thing is that they reproduce himself unlike chemical compounds we want to get.

And we use today a lot of organic compounds and most of them are made by bacteria. They are very potential chemical factories.

And about getting bacteria who specialized on sulfur. Just put a thiosulfate solution in open container, after some time you will find a sulfur on the bottom. It is work of bacteria. So, you know where to get it.

[Edited on 12-9-2019 by teodor]

[Edited on 12-9-2019 by teodor]

Tsjerk - 12-9-2019 at 03:24

No, it is not like alcohol / vinager, yeast happily grows with our without oxygen but will produce either product depending on whether oxygen is present. Sulfate reducing microorganisms literally die when oxygen is present.

How do you imaging producing Se? From selenium oxides I assume? Yes, that is possible, but also these species are anaerobic.

Do you have a reference for your thiosulfate in a bucket bacteria sulfur claim? I would be happily surprised if you manage to find these bacteria as, said before, they don't tolerate oxygen. They grow meters deep in anoxic mud meters under water.

teodor - 12-9-2019 at 03:37

Quote: Originally posted by Tsjerk  

Do you have a reference for your thiosulfate in a bucket bacteria sulfur claim?


I red it in one really old book but I usually trust to the information it contains. Treadwell & Hall, Analytical Chemistry, Volume II. 7nth edition. p 551 (part Volumetric Analysis, chapter "Preparation of Sodium Thiosulfate Solution").

"A solution of pure sodium thiosulfate in "best" water will keep very well but thiosulfate solutions usually deposit sulfur on standing and the titer changes until the decomposition brought about by impurities is complete. The principal cause of the decomposition is bacterial action. Sterile solutions , free from carbon dioxide, keep indefinitely"

Well, it could be that the other form of bacteria produce some "impurities" like CO2 mentioned which decompose the thiosulfate and it is not sulfur-producing bacteria, so I can be wrong.

[Edited on 12-9-2019 by teodor]

Tsjerk - 12-9-2019 at 04:01

Hmmm, I think it is the CO2 that forms an acidic solution that decomposes the sulfate. But I don't think the CO2 is produced by microorganisms. There is nothing for them to grow on.

Sterilization also drives of dissolved CO2, which would explain why a sterile solution does not decompose.

Metacelsus - 12-9-2019 at 04:41

Quote: Originally posted by Tsjerk  
You don't want to work with sulfate reducing bacteria/archaea... They are facultative anaerobic. If you open the jar or whatever you are growing them in, they die.


The term is "obligate anaerobic." Facultative means that they can either use oxygen or grow anaerobically.

Tsjerk - 12-9-2019 at 09:03

Quote: Originally posted by Metacelsus  


The term is "obligate anaerobic." Facultative means that they can either use oxygen or grow anaerobically.


I noticed, it was a bit too early for me to post stuff. I hoped no one would notice.

teodor - 13-9-2019 at 01:44

http://digicoll.library.wisc.edu/cgi-bin/JCE/JCE-idx?type=tu...

"Aqueous solutions of sodium thiosulfate exposed to the air undergo slow decomposition due to oxidation by dissolved oxygen and, ocassionally, to the growth of sulfur-consuming microorganisms (thiobacteria)".

Tsjerk - 13-9-2019 at 22:54

Quote: Originally posted by teodor  
http://digicoll.library.wisc.edu/cgi-bin/JCE/JCE-idx?type=tu...

"Aqueous solutions of sodium thiosulfate exposed to the air undergo slow decomposition due to oxidation by dissolved oxygen and, ocassionally, to the growth of sulfur-consuming microorganisms (thiobacteria)".


Check where they base that claim on... reference 10. It is the book I already thought to be incorrect.

Apparently there are species that can tolerate oxygen, but no metabolism was observed while being oxic. Apparently this was novel enough to be published in an impact 4 journal in 2004.

The pathways needed to reduce sulfate just aren't compatible with oxygen. You can compare it with carbon monoxide in the human body. It interacts with the system, quite well, but the system can't deal with it.

teodor - 14-9-2019 at 05:26

Thiosulfate reacts with oxygen, so may be its level in the solution is much lower than in usual water. Especially in stoppered bottles.

[Edited on 14-9-2019 by teodor]

Junk_Enginerd - 17-9-2019 at 01:46

Quote: Originally posted by teodor  
So, is it possible to reduce, say CaSO4 by Zn and heating? Because there is another thread about converting CaSO4 to something useful and couple of new threads about getting CaS. It seams that Zn could be an options for all of them. Considering the melting point of Zinc it could be much better route than reducing with carbon.

[Edited on 9-9-2019 by teodor]


Well, I tried it and it surprised me a bit. I was certainly expecting some sort of exothermic reaction, but it was more vigorous than I would have thought. Quite close to thermite, though not as insanely hot.

It took a whole lot of heating. I would estimate ignition occurred at perhaps 900-1000 deg C based on the black body color. Once initiated it jumped to probably 1400 deg C and burned vigorously.

At this point I was tired and my experiment progressed a little sloppy. I noted things that indicated I was approaching sulfur at least. Wifts of SO2 certainly, and a yellow tinge that suggested sulfur.

I ground the result into a powder and added sulfuric acid. It reacted quite exothermically for a few minutes and I think there was sulfur suspended in the liquid. Then I realized sulfuric acid was quite a poor choice because I didn't have a sulfuric acid tolerant filter and boiling it off didn't seem like a very good option either.

Then I thought I might neutralize the acid to filter it off as a salt instead, and I added some sodium bicarbonate. I think I may have ruined the result at this point because the sulfur yellow color dissapeared and then I just had a brown sloppy goo.

Either way, the reduction part certainly worked.

Junk_Enginerd - 17-9-2019 at 01:51

Quote: Originally posted by walruslover69  
I think sticking with a SO2 method instead of H2S is the safer and better way to do. You could generate SO2 rather consistently by hot sulfuric acid and copper on a regular hot place. Then reduce the SO2 over carbon, zinc or magnesium. sulfur should be easy to extract using xylene or toluene. It does sound laborious though.


I have an SO2(+SO3 I guess) generator based on pyrolysis of iron sulfate, so that's no problem. I only have a limited quantity of sulfuric acid purified from lead acid batteries, so I don't have a reliable supply and would rather not waste any more than necessary.

Could you add some details regarding the reduction of the SO2? Does it have to be very hot carbon/zinc/mg? Or would room temperature work? 100 degrees C? 500? Just a ballpark...

Junk_Enginerd - 17-9-2019 at 01:58

As another experiment I tried synthesising sodium sulfite as an alternative route, by bubbling SO2 through NaOH. It seemed successful, but I would like to further refine it to sodium thiosulfate. From what I can gather it should be simple, but I can't find a clear enough description of the process to understand it. Does anyone here know?

teodor - 17-9-2019 at 12:53

Quote: Originally posted by Junk_Enginerd  

.
Either way, the reduction part certainly worked.


Wikipedia says that reduction with C also can go this strange way - CaSO4 acts as an oxidizer upon CaS:

3 CaSO4 + CaS = 4 CaO + 4 SO2

The same way you probably can get SO2 from oxidation of CaS by concentrated H2SO4. So, for the purpose to test the substance on presence of CaS it is important to use diluted (1M) sulfuric acid (H2S!).

Presence of CaO or even CaO2 you can check just with water (also H2S from CaS).

Quote: Originally posted by Junk_Enginerd  
As another experiment I tried synthesising sodium sulfite as an alternative route, by bubbling SO2 through NaOH. It seemed successful, but I would like to further refine it to sodium thiosulfate. From what I can gather it should be simple, but I can't find a clear enough description of the process to understand it. Does anyone here know?


As far as I know you need either S or alkali polysulfide to disproportionate Na2SO3 to Na2S2O3. 2 atoms of sulfur in thiosulfate have different valencies, by this reason I hardly can imagine how to get it from only one sulfur compound.

UPD. I was probably too tired yesterday to realize that in a case of the reaction:

4S + 6 NaOH = 2 Na2S + Na2S2O3 + 3 H2O

we get 2 different oxidation states of sulfur starting from only one. Probably by understanding how it happens (in a water solution) we can plan other experiments (I think the reaction goes at least in 2 steps). The proportion Na2S/Na2S2O3, by the way, according to my own observation, can be probably shifted if we reflux the water solution with additional components, like alcohol.



UPD UPD. I did some silly experiment of boiling CuSO4 with S. Definitely some reaction happens but not with the rate of SO3{2-} with S. Also I was unable to identify a small quantity of black compound I've got as a result.

[Edited on 18-9-2019 by teodor]

woelen - 17-9-2019 at 23:04

The disproportionation reaction of S with NaOH is not very special. It is an example of a broader class of reactions. Other elements show similar behavior:

- chlorine gives chloride and hypochlorite (and on standing or heating the hypochlorite gives more chloride and chlorate).
- bromine and iodine give similar reactions
- white phosphorus gives phosphine and hypophosphite
- selenium gives polyselenides and selenite (there is no analog of thiosulfate for selenium)

In all cases you get from the element (oxidation state 0) a species with positive oxidation state and a species with negative oxidation state. The reaction is not always clean in the sense that exactly two species are formed. The resulting solutions can be quite complicated. This is especially the case with S, Se and P.

walruslover69 - 18-9-2019 at 09:25

Quote: Originally posted by Junk_Enginerd  

I have an SO2(+SO3 I guess) generator based on pyrolysis of iron sulfate, so that's no problem. I only have a limited quantity of sulfuric acid purified from lead acid batteries, so I don't have a reliable supply and would rather not waste any more than necessary.

Could you add some details regarding the reduction of the SO2? Does it have to be very hot carbon/zinc/mg? Or would room temperature work? 100 degrees C? 500? Just a ballpark...


Thermodynamically zn/mg should work at about any temp. The problem at room temp is going to be reaction rate, but if you have it any hotter than 388C you will boil any sulfur formed. Thinking about it now i think you would run into problems doing it above 100C too because the molten sulfur would probably initiate a reaction with the Zn/Mg producing the sulfide. Getting the reaction conditions might be very tricky.

teodor - 12-5-2020 at 14:16

Quote: Originally posted by Tsjerk  
Quote: Originally posted by teodor  
http://digicoll.library.wisc.edu/cgi-bin/JCE/JCE-idx?type=tu...

"Aqueous solutions of sodium thiosulfate exposed to the air undergo slow decomposition due to oxidation by dissolved oxygen and, ocassionally, to the growth of sulfur-consuming microorganisms (thiobacteria)".


Check where they base that claim on... reference 10. It is the book I already thought to be incorrect.

Apparently there are species that can tolerate oxygen, but no metabolism was observed while being oxic. Apparently this was novel enough to be published in an impact 4 journal in 2004.

The pathways needed to reduce sulfate just aren't compatible with oxygen. You can compare it with carbon monoxide in the human body. It interacts with the system, quite well, but the system can't deal with it.


OK, I eventually found the answer in the old book "Standard Methods for examination of Water, Sewage, and Industrial Waste" 10th edition.

There are 2 types of bacteria - Sulfur and Sulfate-Reducing. I attached scan how to grow them. Probably they both can be detected in water samples.
That one which grows in thiosulfate actually oxidise sulfur. It is probably Thiobacillus thioparus. And the sulfate-reducing specia, you are right, has 2 possible way to live - consume organic matter or consume hydrogen. The book claims it can grow in iron pipes consuming sulfates and hydrogen from H2O + Fe reaction. Also, from wikipedia on Desulfovibrio: "Like other sulfate-reducing bacteria, Desulfovibrio was long considered to be obligately anaerobic. This is not strictly correct: while growth may be limited, these bacteria can survive in O2-rich environments. These types of bacteria are known as aerotolerant."

To catch this bacteria you can use this information (also from wikipedia): " Desulfovibrio species are commonly found in aquatic environments with high levels of organic material, as well as in water-logged soils, and form major community members of extreme oligotrophic habitats such as deep granitic fractured rock aquifers".

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