MeshPL
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Iron catalyzed sulphuric acid production
Hello everybody!
Reading wikipedia i got an idea: how to produce sulphuric acid at home cheaply and easily. Not sure, if it will work though.
According to: https://en.wikipedia.org/wiki/Sulfuric_acid, the following reactions can occur (and do in rocks):
2FeS2 + 7O2 + 2H2O --> 2FeSO4 + 2H2SO4
2FeSO4 + H2SO4 + 1/2O2 --> H2O + Fe2(SO4)3
FeS2 + 7Fe2(SO4)3 + 8H2O --> 15FeSO4 + 8H2SO4
Presence of iron sulphates as a catalyst increases the rate of reaction, allowing the second and third reaction to work. Pre-adding some kind of
ferric salt as a catalyst would be a good idea.
If I would bubble enough air through FeS2 slurry, and stop before all FeS2 dissolves (I don't want the second reaction to waste my H2SO4), I should be
able to disstil H2SO4 out of the solution.
I also got an idea to replace FeS2 with elemental sulphur. According to wikipedia "List of standard Electrode Potentials", the following reactions
would potentialy occur. Note the need of iron catalyst this time:
S + 2H2O + 2Fe2(SO4)3--> SO2 + 4FeSO4 + 2H2SO4
SO2 + 2H2O + Fe2(SO4)3 -->2FeSO4 + 2H2SO4
2FeSO4 + H2SO4 + 1/2O2 --> H2O + Fe2(SO4)3
Would that be at least possible?
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byko3y
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It's nice to see someone trying to invent the same shit people have been doing for 400 years.
https://books.google.com/books?id=O4rzzkUQyzIC&pg=PA8&am...
Fe2S + O2 + H2O -> FeSO4 -> Fe2(SO4)3
Fe2(SO4)3 -> Fe2O3 + SO3
Maybe some day the idea of oxidizing the SO2 with nitrate salt will draw your attention.
[Edited on 19-7-2015 by byko3y]
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MeshPL
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For some reason I cannot view, that book...
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Fe2S + O2 + H2O -> FeSO4 -> Fe2(SO4)3
Fe2(SO4)3 -> Fe2O3 + SO3
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The second reaction seems to require high temperature. I'm well aware of the process for SO3 production requiring roasting of iron sulphates, but I
was curious, if you could do that without roasting, just distilling H2SO4 at thge right time out of pyrite slurry, to maximise the yield.
Also, is oxidation of elemental sulphur by iron (III) possible?
[Edited on 19-7-2015 by MeshPL]
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byko3y
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I'm pretty sure that many of you reactions require high temperatures. However I might be wrong and you will invent a new way of producing sulfuric
acid.
Some time ago I posted an idea about doing almost the same using copper as a renewable source of sulfate. http://www.sciencemadness.org/talk/viewthread.php?tid=62575&... It has one beautiful property - CuS is completely insoluble in sulfuric acid, thus
making it possible to collect precipitate and obtain a pure H2SO4. Drawbacks: toxicity and odor of H2S; some oxidation reactions might be tricky.
Also, you might already know that copper is used by enzymes for aerobic oxidation.
UPD: I forgot to mention that the iron catalyst is actually one of the options as a contact catalyst. AFAIK, it requires high temperatures, which
leads to decomposition of formed SO3 into SO2, thus leading to lower yields. Palladium has highest reactivity, but unfortunutely it also reacts with
impurities, that's why SO2 coming into catalyst bed should be exceptionally pure. Palladium-based contact process was used in the industry (somewhere
at 1900-1920) until vanadium catalyst was discovered, which is less vulnerable for poisoning.
[Edited on 19-7-2015 by byko3y]
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unionised
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Congratulations on posting a link to a Ukrainian web site.
This thread needs to remember that geology has plenty of time on its hands.
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aga
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Interesting point.
Perhaps reaction formulae should come with a clue as to the time taken in various conditions, a bit like catalysts are noted.
star
H ==> Pb (2.3 billion years +/-486.25 million)
hot
[Edited on 19-7-2015 by aga]
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MeshPL
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| Quote: |
star
H ==> Pb (2.3 billion years +/-486.25 million)
hot
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No, lead requires supernovae to form. As far as I remember nothing stable heavier than iron can be formed in regular stars in reasonable quantities.
So:
Supernovae
Anything ==> Pb (~few seconds)
F****** hot, high pressure
Note: the reaction is dangerous and may result in splashing of reaction mixture everywhere. The yields are low, and there are many by-products
including gold, platinum, iridium etc. which will need to be removed.
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I'm pretty sure that many of you reactions require high temperatures.
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Maybe... electrode potentials, measured for 20C, indicate, they MAY occur.
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This thread needs to remember that geology has plenty of time on its hands.
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In case of mines, acid mine drainage can occur within a few years... also geology cannot bubble fresh air through slurry of pulverized matterial.
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Boffis
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No, nature works the other way round and pours the water through a porous aerated honeycomb full of sulphides! Very effective!
The process of pyrite oxidation is complex and often bacterially assisted but the key process in the formation of free sulphuric acid is the
precipitation of the basic potassium ferric hydroxide sulphate minerals such as jarosite KFe3(SO4)2(OH)6 at about pH 2-3 from solutions containing
normal ferric sulphate rough according to:
K<sub>2</sub>SO<sub>4</sub> + 3Fe<sub>2</sub>(SO<sub>4</sub> <sub>3</sub> + 12H<sub>2</sub>O →
2KFe<sub>3</sub>(SO<sub>4</sub><sub>2</sub>(OH)<sub>6</sub> + 6H<sub>2</sub>SO<sub>4</sub>
This equation is only illustrative of the formation of one mineral and there are a vast number of natural iron sulphates.
 Curpocopiapite forming on openpit walls in Kazakhstan
 The iron precipitates but the copper stays in solution!!
 Copper sulphate encrusted blocks eastern Kazakhstan
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AJKOER
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Per Boffis comment "bacterially assisted", Wikipedia's ( link: https://en.m.wikipedia.org/wiki/Acidithiobacillus_ferrooxida...) says, to quote:
" Acidithiobacillus ferrooxidans (syn. Thiobacillus ferrooxidans) lives in pyrite deposits, metabolizing iron and sulfur and producing sulfuric acid."
[Edited on 20-7-2015 by AJKOER]
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