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Author: Subject: Preparation of Dilute H2SO4 from FeSO4
AJKOER
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[*] posted on 22-1-2012 at 10:28
Preparation of Dilute H2SO4 from FeSO4


It occurred to me that a thread of preparation of H2SO4 directly from FeSO4 (cheap and available) is probably worthy of mention. But, before I even attempt to list possible paths to H2SO4, I should mention that in some cases, the direct employment of FeSO4 as an H2SO4 substitute should be mentioned.

For example, the reaction of an excess of FeSO4 and bleach (NaOCl/NaCl), appears to effectively produce Chlorine. However, the mechanics of this route, appears to be less clearly understood. My speculation is as a by product of the reaction of Ferrous hypochlorite (unstable) and Ferrous chloride in a multi-step reaction:

FeSO4 + 2 NaOCl ---> Fe(OCl)2 + Na2SO4

FeSO4 + 2 NaCl ---> Na2SO4 + FeCl2

FeCl2 + Fe(OCl)2 ---> 2 Cl2 + 2 FeO

2 FeCl2 + Cl2 ---> 2 FeCl3

There are, however, other possible paths:

1. 2 NaClO --Ferrous ion--> 2 NaCl + O2

2. FeSO4 + O2 + 10H2O --> 4 Fe(OH)3 + H2SO4

3. H2SO4 + NaCl --> NaHSO4 + HCl

4. H(+) + NaOCl --> Na(+) + HOCl

5. HCl + HOCl <----> Cl2 + H2O

6. 3 HCl + Fe(OH)3 ---> FeCl3 + 3 H2O

Now, as to the main topic how to produce H2SO4 from FeSO4. First thermal decomposition of Ferrous Sulfate, which commences (per Wkipedia) at 480 C:

2 FeSO4 ---Heat---> Fe2O3 (s) + SO2(g) + SO3 (g)

Capturing gases and further oxidizing (via H2O2 or with O2 with some MnSO4, acting as a catalyst):

SO2 + H2O2 --> H2SO4

SO3 + H2O --> H2SO4

Method 2. By reacting with HCl (or other strong acid):

FeSO4 + 2 HCl ---> FeCl2 + H2SO4

Method 3. This reaction has been noted previously:

4 FeSO4 + O2 + 10 H2O --> 4 Fe(OH)3 + 4 H2SO4

Reaction speed and dilution may be issues.

Reference: "Physicochemical simulation of calcite (dolomite)-FeSO4-H2O open systems" by I. P. Kremenetskaya, O. P. Korytnaya, T. N. Vasil’eva, A. T. Belyaevskii, G. I. Kadyrova and S. I. Mazukhina. LINK:

http://www.springerlink.com/content/h24543g24x366822/

Note, with an excess of FeSO4 (should be avoided), the reaction with H2SO4 and oxygen could just form Fe2(SO4)3 and no H2SO4 as:

4 FeSO4 + 2 H2SO4 + O2 ---> Fe2(SO4)3 + 2 H2O

Comments and other ways (like electro chemical) to employ FeSO4 are welcomed.


[Edited on 22-1-2012 by AJKOER]
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weiming1998
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[*] posted on 22-1-2012 at 16:04


For method 3, wouldn't the H2SO4 formed immediately react with the Fe(OH)3?
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[*] posted on 22-1-2012 at 16:28


If I remember correctly, the dry distillation of iron sulfate actually makes fuming sulfuric, or oleum.

But is there some reason that you're going through all these contortions to make H2SO4?

Is it academic interest? Or what? I hate to use the T word, but your posts kinds of have that aroma.

Some chemicals are just not worth making. I make hydrochloric acid just because the hardware stores no longer carry it and I kind of enjoy making chemicals.

But making sulfuric acid just seems to be so 17th century. If you don't know where to buy it, I can tell you.
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[*] posted on 22-1-2012 at 16:48


Quote: Originally posted by AJKOER  
Method 2. By reacting with HCl (or other strong acid):

FeSO4 + 2 HCl ---> FeCl2 + H2SO4



*sigh*
If only it were that easy. Do you have an enthalpy of reaction value for that reaction?

Have no fear, I'll calculate it for you: -46kJ/mol (give or take).

On paper it seems like it would work, but in practice, why would you want to swap HCl for sulfuric acid?

Your FeSO4 + 2 NaCl ---> Na2SO4 + FeCl2 is interesting to say the least. I doubt that reaction would take place smoothly in practice, even though I believe the ∆H works out to be negative (I didn't double check it so don't take my word for it).




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[*] posted on 22-1-2012 at 17:10


Quote: Originally posted by entropy51  
If I remember correctly, the dry distillation of iron sulfate actually makes fuming sulfuric, or oleum.

But is there some reason that you're going through all these contortions to make H2SO4?

Is it academic interest? Or what? I hate to use the T word, but your posts kinds of have that aroma.

Some chemicals are just not worth making. I make hydrochloric acid just because the hardware stores no longer carry it and I kind of enjoy making chemicals.

But making sulfuric acid just seems to be so 17th century. If you don't know where to buy it, I can tell you.


I don't. Tried drain cleaner. No sulfuric acid drain cleaners. Tried asking low fume pool acid at Sigma Chemicals. They only come in 15L drums, and even if I want to buy it, I probably can't because the workers were already beginning to ask a frenzy of questions (like what's your pool's condition?, Do you have a sample of your pool water with you? etc. Buying a whole car battery, then draining the acids is too expensive and I can't find anywhere that sells car battery electrolytes.
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[*] posted on 22-1-2012 at 18:59


I can't believe Oz is so backward that they don't have sulfuric drain openers or battery acid. Somewhere.

I can believe that people don't want to sell them to 13 year olds, however.

It took me several years of chemistry discussions to earn the trust of the local pharmacists, but once I did, they would order anything I wanted.

Did you know that you can make HNO3 by dry distillation of a mixture of alum, CuSO4, and NaNO3? Of course you can't do this until you find a source of nitrates.

Less whining and more looking is my advice. Good luck.

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[*] posted on 22-1-2012 at 20:00


Quote: Originally posted by entropy51  
I can't believe Oz is so backward that they don't have sulfuric drain openers or battery acid. Somewhere.

I can believe that people don't want to sell them to 13 year olds, however.

It took me several years of chemistry discussions to earn the trust of the local pharmacists, but once I did, they would order anything I wanted.

Did you know that you can make HNO3 by dry distillation of a mixture of alum, CuSO4, and NaNO3? Of course you can't do this until you find a source of nitrates.

Less whining and more looking is my advice. Good luck.



Oh well, just asking. If you people won't help, then I'll look for them on my own like I've always did. Simple as that. And no, I'm not complaining.
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[*] posted on 22-1-2012 at 20:07


Also, if I tried to talk about those pool workers about how I am doing home chemistry and need sulfuric acid from the low- fume pool acids, then they will either:
.Sell nothing to me
.Sell things to me and act normal, then ring the police.
.Sell nothing to me AND kick me out of the shop.
.Sell nothing to me, then ring the police.
. Actually having a discussion, and selling me SOME of the things.
. Lecture about how these things are dangerous and telling me to go play/get a girlfriend (like you people have done)
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[*] posted on 22-1-2012 at 20:16


H2SO4 only loses to water on chemical use. So it cant be an easy task to completely hide it from you.

In my country, you (commom people, not firm/enterprise) could only buy max. 2lt of 98% H2SO4 per month from chemical supplier, and with your civil/adress information. Since I dont usually use large amounts of H2SO4 in short times, and dont make anything illegal with it, thats suits my experimental needs very well.

However, before 18 (actually before 16), I have boiled battery acid to obtain my own "98%" sulfuric (which was at most 85-95%). This is, IMHO, the only real way of "making" large amounts of concentrated H2SO4 by your own. The greatest drawback is that is an extremely dangerous operation, I and many members cant stress enough how dangerous is that.. Cold concentrated sulfuric can blind you instantly, so you can imagine what >300°C concentrated could do to your skin/face.. This without mention the SOx/H2SO4 nasty fumes.. Today Im much older and think that I dont need to expose myself with this kind of danger.

Stick with entropy51 idea. It dont worth the effort to make your own H2SO4, except only if you really cant find it (which is probably impossible.. Otherwise, you havent looked hard enough). Try to convince your parents or older friends to buy it for you (remember, if you have on hands a bottle of H2SO4 , treat it with greatest care.. It will not be nice if have a chance to hit you).

For me, the idea of FeSO4 pyrolisis being useful is just one: oleum. Oleum, unlike normal H2SO4 is very hard to come across and is a very valuable chemical, and anhydrous ferrous sulfate could be put in use to make it instead of plain sulfuric.. (even though, better options for SO3 generation are up to amateur reach and are already discussed in the forum), just bubbling SO3 in conc. H2SO4.

[Edited on 23-1-2012 by Aqua_Fortis_100%]




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weiming1998
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[*] posted on 22-1-2012 at 20:37


Quote: Originally posted by Aqua_Fortis_100%  
H2SO4 only loses to water on chemical use. So it cant be an easy task to completely hide it from you.

In my country, you (commom people, not firm/enterprise) could only buy max. 2lt of 98% H2SO4 per month from chemical supplier, and with your civil/adress information. Since I dont usually use large amounts of H2SO4 in short times, and dont make anything illegal with it, thats suits my experimental needs very well.

However, before 18 (actually before 16), I have boiled battery acid to obtain my own "98%" sulfuric (which was at most 85-95%). This is, IMHO, the only real way of "making" large amounts of concentrated H2SO4 by your own. The greatest drawback is that is an extremely dangerous operation, I and many members cant stress enough how dangerous is that.. Cold concentrated sulfuric can blind you instantly, so you can imagine what >300°C concentrated could do to your skin/face.. This without mention the SOx/H2SO4 nasty fumes.. Today Im much older and think that I dont need to expose myself with this kind of danger.

Stick with entropy51 idea. It dont worth the effort to make your own H2SO4, except only if you really cant find it (which is probably impossible.. Otherwise, you havent looked hard enough). Try to convince your parents or older friends to buy it for you (remember, if you have on hands a bottle of H2SO4 , treat it with greatest care.. It will not be nice if have a chance to hit you).

For me, the idea of FeSO4 pyrolisis being useful is just one: oleum. Oleum, unlike normal H2SO4 is very hard to come across and is a very valuable chemical, and anhydrous ferrous sulfate could be put in use to make it instead of plain sulfuric.. (even though, better options for SO3 generation are up to amateur reach and are already discussed in the forum), just bubbling SO3 in conc. H2SO4.

[Edited on 23-1-2012 by Aqua_Fortis_100%]


Actually, my parents are considering to buy things off eBay for me. So maybe in the not-so-distant future, I will get it. You are right about hot, concentrated sulfuric acid being dangerous, and I will take all the precautions necessary. SO3 is even more dangerous than simply H2SO4, and can dehydrate/carbonize your skin instantly. Also, would an international shipping of something hazardous, like sulfuric acid, be extremely expensive? Chemical suppliers don't cater to individuals here.
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[*] posted on 22-1-2012 at 21:22


Quote: Originally posted by weiming1998  
Oh well, just asking. If you people won't help, then I'll look for them on my own like I've always did. Simple as that. And no, I'm not complaining.
I don't know what you mean by "You people won't help." We can't buy chemicals for you. If you were in the US, we could refer you to all kinds of suppliers. But you aren't, and we can't. I have offered all the advice that I can. And you do sound like a lot of the kewls who visit here and bitch about not being able to obtain chemicals even though you are "entitled" to .

To tell the truth, the more you reveal about your self, the more I doubt that you could safely handle nitric acid.
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[*] posted on 22-1-2012 at 21:57


I really think I should stop talking to you now.

Also, sorry for sounding a bit "rude" or begging for help. I didn't mean to word it like that. Finally, no matter what you or anybody else thinks, I am not a k3wl that wants to blow things up with home-made explosives.

[Edited on 23-1-2012 by weiming1998]

[Edited on 23-1-2012 by weiming1998]
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[*] posted on 22-1-2012 at 22:43


Quote: Originally posted by weiming1998  
For method 3, wouldn't the H2SO4 formed immediately react with the Fe(OH)3?


First, here is another reference: "Hydrometallurgy in extraction processes", Volume 1, by C. K. Gupta, T. K. Mukherjee, pages 59 to 60:

"In some operations, the ferrous sulfate solution is left in ponds exposed to the air for several months so that oxidation and hydrolysis take place, thus regenerating the acid. The overall reaction is:

4 FeSO4 + O2 + 10H2O --> 4 Fe(OH)3 + 4 H2SO4 (22)

The oxidation of the ferrous sulfate is catalyzed by the autotrophic microorganism Thiobacullis ferrooxidans."

LINK:
http://books.google.com/books?id=F7p7W1rykpwC&pg=PA60&lpg=PA60&dq=4+FeSO4+%2B+O2+%2B+10+H2O+4+Fe(OH)3+%2B+4+H2SO4&source=bl&ots=fi SSl16x5f&sig=VsoTmxn7wDiLNncwWFwZPJB0SYU&hl=en&sa=X&ei=O_ocT-2DDefj0QHM1JixCw&sqi=2&ved=0CGkQ6AEwCQ#v=onepage&q=4%20FeSO4% 20%2B%20O2%20%2B%2010%20H2O%204%20Fe(OH)3%20%2B%204%20H2SO4&f=false

Note, the equation is nearly the same as the one I quoted. As I suspected, there is a time element and to my surprise, even a bacterial catalysis.

To answer your good question, I suspect that the jelly like Fe(OH)3 is physically separated from the H2SO4, although there could be, per your point, some Iron sulfate at the boundary itself, depending on the acid strength of the solution.
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[*] posted on 22-1-2012 at 23:02


Now I get why a freshly-made solution of FeSO4 is green but turns yellow over time. It's the gradual formation of Fe(OH)3!

Also, this reaction FeSO4 + 2 NaCl ---> Na2SO4 + FeCl2 works because FeCl2 is much more soluble in cold water than Na2SO4. So, in theory, this reaction (FeSO4 + 2 HCl ---> FeCl2 + H2SO4) wouldn't occur, because:
1, H2SO4 is a stronger acid
2, FeCl2 cannot precipate without his happening: FeCl2+H2SO4===>2HCl+FeSO4, driving the reaction back where it started.
3, The volatility of HCl is significantly higher than H2SO4
If this reaction really occurs, then it would be in an equilibrium, and a certain condition is required to prevent FeCl2 from reacting with the H2SO4.
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[*] posted on 23-1-2012 at 00:00


Quote: Originally posted by White Yeti  
Quote: Originally posted by AJKOER  
Method 2. By reacting with HCl (or other strong acid):

FeSO4 + 2 HCl ---> FeCl2 + H2SO4



*sigh*
If only it were that easy. Do you have an enthalpy of reaction value for that reaction?

Have no fear, I'll calculate it for you: -46kJ/mol (give or take).

On paper it seems like it would work, but in practice, why would you want to swap HCl for sulfuric acid?

Your FeSO4 + 2 NaCl ---> Na2SO4 + FeCl2 is interesting to say the least. I doubt that reaction would take place smoothly in practice, even though I believe the ∆H works out to be negative (I didn't double check it so don't take my word for it).


OK, here is a YouTube on FeSO4 + 2 HCl:

http://www.youtube.com/watch?v=MDm37HPc1X8

Yes, I agree with you on the FeSO4 + 2 NaCl reaction, and the comments at:

http://www.mychemistrytutor.com/questions/3197

However, this spectator reaction is presented in the context of:

FeSO4 + 2 NaOCl ---> Fe(OCl)2 + Na2SO4

to possibly explain:

FeCl2 + Fe(OCl)2 ---> 2 Cl2 (g) + 2 FeO

which actual makes sense in water as:

2 Cl2 + 2 H2O <---> 2 HCl + 2 HOCl

and upon adding in the FeO and grouping terms:

2 HCl + FeO ---> FeCl2 + H2O

2 HOCl + FeO ---> Fe(OCl)2 + H2O

(this last equation being the classic preparation of a hypochlorite by direct reaction with HOCl although even the temporary existence of Iron hypochlorite is debatable, but would be congruent with the known observed stability issues associated with Aluminum hypochlorite and also Zinc hypochlorite)

It would even appear that the following speculated equation is reversible:

FeCl2 + Fe(OCl)2 + 2H2O <---> 2 Cl2 + 2 FeO + 2 H2O

in a closed system.

Note, the evolution of Chlorine is definitely observed when FeSO4 is added to Bleach, and I am just speculating on possible paths. The fact that the literature appears to be uncomfortable with speculating on reaction products and paths here (and similarly in the case of the reaction of Fe in dilute HOCl forming Cl2 and FeCl3) leaves it to the practicing chemist to take his best guess (please correct me if I have mis-spoke).
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weiming1998
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[*] posted on 23-1-2012 at 01:02


By watching the video you've posted, I think I know what's happening. What's happening is that the Fe+,SO4-2, Cl- and H+ are all floating in the solution. Because they are all free ions, some H2SO4 or FeCl will be formed. But when you boil down the solution, the equilibrium shifts back to HCl and FeSO4 because HCl is more volatile. It won't even work if you try to dehydrate it at room temperature, because the bonds connecting Fe to SO4 is stronger than the bonds connecting it to Cl-, and the ion rejoins with the SO4-2. However, if you place some CaCO3 in the solution, some CaSO4 precipate will form, evidence that sulfuric acid exists in the solution.

The reason why FeSO4 is oxidized and hydrolyzed to Fe(OH)3 and H2SO4 is also because of this. As FeSO4 is a very weak base and also insoluble in water, It will precipate, driving the reaction forward. This can only form small amounts of H2SO4, as small amounts doesn't react with the Fe(OH)3/reacts so slowly that the oxidation and hydrolysis of FeSO4 is faster. However, once the concentration of H2SO4 gets to a certain point, the speed of acid-base neutralization and the hydrolysis of FeSO4 are the same, enabling no more H2So4 to form. If too much H2SO4 is in the solution, it will revert back to FeSO4 and H2O until it reaches this point, where both reaction stops.
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[*] posted on 23-1-2012 at 07:39


Quote: Originally posted by entropy51  
Is it academic interest? Or what? I hate to use the T word, but your posts kinds of have that aroma.



You don't say... ;)

'AJKOER's world of chemistry' revolves aroung hypochlorite and the mysterious 'NH4OH'.




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[*] posted on 23-1-2012 at 15:29


blogfast25:

FYI, some of my recent threads and responses had included with specific reference to the halogen family only, the following compounds:

Cl2, SOCl2, HCl + H2O2, HOCl, Cl2O, ClO2, NaCl, NaClO, CaCl2, Ca(OCl)2, Al(ClO)3, NaClO2, NaClO3, Mg(ClO)2, Mg(ClO3)2, Zn(ClO)2, KCl, KClO3, Zn(ClO3)2, FeCl2, FeCl3, Fe(OCl)2, HClO3 and even HClO4 (a world patent describing preparation of Perchloric acid via electrolysis starting with chloride fee HOCl).

With respect to Iodine:

I2, HI, CaI2, KI, HIO, HIO3, Ca(IO3)2 and KIO3.

so one would think you would have second thoughts on your comment "AJKOER's world of chemistry' revolves aroung hypochlorite and the mysterious 'NH4OH' " (surprise me!).

With respect to your academic/theoretical issue with this thread, with all the equations and postulates I have put forth, one would think you would offer at least one counter point or reference an error (surprise me!).


[Edited on 24-1-2012 by AJKOER]
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[*] posted on 23-1-2012 at 20:08


Quote: Originally posted by weiming1998  

The reason why FeSO4 is oxidized and hydrolyzed to Fe(OH)3 and H2SO4 is also because of this. As FeSO4 is a very weak base and also insoluble in water, It will precipate, driving the reaction forward. This can only form small amounts of H2SO4, as small amounts doesn't react with the Fe(OH)3/reacts so slowly that the oxidation and hydrolysis of FeSO4 is faster. However, once the concentration of H2SO4 gets to a certain point, the speed of acid-base neutralization and the hydrolysis of FeSO4 are the same, enabling no more H2So4 to form. If too much H2SO4 is in the solution, it will revert back to FeSO4 and H2O until it reaches this point, where both reaction stops.


Actually, perhaps not as those little microbes essentially may be driving the process and actually creating a fairly acidic product. To quote:

"Thiobacillus ferrooxidans is the most common type of bacteria in mine waste piles. This organism is acidophilic (acid loving), and increases the rate of pyrite oxidation in mine tailings piles and coal deposits. It oxidies iron and inorganic sulfur compounds. The oxidation process can be harmful, as it produces sulfuric acid, which is a major pollutant."

Also:

"This genus is thermophilic, preferring temperatures of 45-50 degrees Celsius. In addition, this is an acidophilic genus, preferring a pH of 1.5 to 2.5. A few species, however, only grow in a neutral pH."

And:

"Thiobacillus are strictly aerobic bacteria. All species are respiratory organisms.
Thiobacillus are obligate autotrophic organisms, meaning they require inorganic molecules as an electron donor and inorganic carbon (such as carbon dioxide) as a source. They obtain nutrients by oxidizing iron and sulfur with O2."

The way I interpret the last sentence is that the bacteria are using the oxygen, by assimilation and employing it as a tool, to create nutrients and H2SO4. Otherwise, O2 would not be reacting with the FeSO4.

Source: MicrobeWiki, the student-edited microbiology resource.

LINK:
http://microbewiki.kenyon.edu/index.php/Thiobacillus


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[*] posted on 24-1-2012 at 03:17


Hang on, this article says that this bacteria oxidizes Fe2+ to Fe3+ using oxygen and H+ ions! H+ ions? It would be consuming the H2SO4 formed by the oxidation of FeSO4. The formula would be this: 4FeSO4+2H2SO4+O2===>2Fe2(SO4)3+2H2O! It consumes the H2SO4 formed, not create it! So after a long time, you would end up with a solution of mainly Fe2(SO4)3, with traces of H2SO4 and Fe(OH)3. The solution would likely to end up less acidic than before.
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[*] posted on 24-1-2012 at 04:59


There are bacteria that consume sulfide and produce sulfate, the blue skin on stagnant water which you will find in healthy bogs is the residue of such bacteria. Incredibly low pH levels can be utilized by these little guys. The masses of iron oxide/hydroxide they leave behind is known as 'bog iron'.

By some mechanism, ether very low pH or other, I have seen a lot of evidence that these colonies are able to literally dissolve rocks and sand.

got fools gold? got to much time on your hands? got a nearby bog?
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[*] posted on 24-1-2012 at 05:37


Quote: Originally posted by Neil  
There are bacteria that consume sulfide and produce sulfate, the blue skin on stagnant water which you will find in healthy bogs is the residue of such bacteria. Incredibly low pH levels can be utilized by these little guys. The masses of iron oxide/hydroxide they leave behind is known as 'bog iron'.

By some mechanism, ether very low pH or other, I have seen a lot of evidence that these colonies are able to literally dissolve rocks and sand.

got fools gold? got to much time on your hands? got a nearby bog?


If my formulas are correct, only FeS2 would be oxidized by the bacteria by:
2FeS2+2H2O+7O2==bacteria===> 2FeSO4+ 2H2SO4. FeSO4 would further be oxidized/hydrolyzed into Fe(OH)3 and more H2SO4. But this time, the concentration of H2SO4 would increase dramatically because the bacteria produces H2SO4, which can, in theory, increase indefinitely until it gets too acidic and bacteria dies. But FeS, which is more available to me and easily made by iron powder+sulfur won't work. It would be: FeS+2O2===>FeSO4 . Anyone got an idea of the artificial synthesis of fool's gold?
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[*] posted on 24-1-2012 at 06:00



Actually, I think both Neil and Weiming1998 are correct. I have seen the reaction written to "completion". Here is an example, based on a quick search:

12FeSO4 +3 O2 +6 H2O = 4Fe2 (SO4) 3 +4 Fe (OH) 3

LINK (hit translate unless you do speak Chinese):
http://wenwen.soso.com/z/q168566732.htm

But, the partial reaction (to H2SO4) does occur as I have seen the reaction (see above citation in "Hydrometallurgy in extraction processes" by C. K. Gupta, T. K. Mukherjee, page 60), and also H2SO4 production is clearly cited in the article as an environmental hazard.

Remember the reaction time is in months!

[Edited on 24-1-2012 by AJKOER]
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Neil
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[*] posted on 24-1-2012 at 06:08


From wiki


"Pyrite has been used since classical times to manufacture copperas, or iron sulfate. Iron pyrite was heaped up and allowed to weather as described above (an early form of heap leaching). The acidic runoff from the heap was then boiled with iron to produce iron sulfate. In the 15th century, such leaching began to replace the burning of sulfur as a source of Sulfuric Acid. By the 19th century, it had become the dominant method."


For our next trick we will knock the edges off a square and attach it to a cart.


Maybe I'm daft but I thought that heating Fe + 2S gave FeS2



The bacteria that use sulphate as an energy source leave behind the iron hydroxides, they live on the wastes of bacteria that feed actual sulfides.

Edit: http://technology.infomine.com/enviromine/ard/Microorganisms...

[Edited on 24-1-2012 by Neil]
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[*] posted on 24-1-2012 at 06:16


Quote: Originally posted by weiming1998  
Also, if I tried to talk about those pool workers about how I am doing home chemistry and need sulfuric acid from the low- fume pool acids, then they will either:
.Sell nothing to me
.Sell things to me and act normal, then ring the police.
.Sell nothing to me AND kick me out of the shop.
.Sell nothing to me, then ring the police.
. Actually having a discussion, and selling me SOME of the things.
. Lecture about how these things are dangerous and telling me to go play/get a girlfriend (like you people have done)




"Hello kind sirs, I like working with copper and making copper inventions/jewellery/art and I really need some sulphuric acid for my pickling bath."

Also check around for jewellery supply shops and welding shops. Look for pickling solution.
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