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Author: Subject: Elemental sulfur from sulfates?
rockyit98
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[*] posted on 11-9-2019 at 21:13


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

AlbertaSulfurAtVancouverBC.jpg - 205kB
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Tsjerk
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[*] posted on 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.
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[*] posted on 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]
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Tsjerk
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[*] posted on 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.
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teodor
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[*] posted on 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]
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Tsjerk
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[*] posted on 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.
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[*] posted on 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.




As below, so above.
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Tsjerk
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[*] posted on 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.
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teodor
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[*] posted on 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)".
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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.
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[*] posted on 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...
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[*] posted on 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?
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[*] posted on 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]
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[*] posted on 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.




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[*] posted on 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.
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