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Author: Subject: Pyrolysis of Metal Sulfates
chemistat
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[*] posted on 27-7-2020 at 09:18
Pyrolysis of Metal Sulfates


I am (like so many others) trying to find a convenient route towards sulfuric acid. I have read that copper sulfate can be pyrolitically decomposed to release SOx gasses. I began to wonder if this could be done with other sulfates and after reading into it, all I could find was a reference in an MSDS for MgSO4 that it "decomposes releasing sulfur dioxide and trioxide at 'very high temperatures'". Does anyone else know anything about what temperatures are involved in this process and whether it works with any other sulfates?

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macckone
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[*] posted on 27-7-2020 at 13:52


Epsom salt decomoses at about 1080C in normal atmosphere.
https://pubs.acs.org/doi/pdfplus/10.1021/ie102554f

All sulfates will decompose, the question is at what temperature?

Sulfur trioxide decomposes starting at 600C (or there abouts).

Iron and copper sulfate are traditional sulfates for making 'vitriol' aka sulfuric acid by decomposition even though the sulfates decompose above the sulfur trioxide temperature, sulfur dioxide is easily oxidized in hot water.

Sodium pyrosulfate decomposes about 460C by 600C it is fully decomposed.
You have competing reactions and sodium bisulfate is readily available.
Calcium sulfate decomposes over 1200C.

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[*] posted on 27-7-2020 at 14:58


I found some very cheap Ferrous Sulfate Heptahydrate and so it too have been thinking about this. Ferric Sulfate decomposes at 750 K according to wikipedia. Well below softening point of pyrex glass but decomp may not be complete at that point. Equation for ferric and the traditional (and presumably less efficient) ferrous sulfate should be:

Fe2(SO4)3 = Fe2O3 + SO3
2 FeSO4 = Fe22O3 + SO2 + SO3

Since Ferric Sulfate is almost insoluble it may(?) be very easily prepared by keeping Ferrous Sulfate solution in contact with Sulfur Dioxide and air and just collecting the Ferric Sulfate as it forms. Another way is to react Ferrous Sulfate with Sodium Persulfate and filtering off the product if it does not form a double salt. Could be a way to get more out of your persulfate. If pyrosulfate decomposes at such a high temperature it seems like not much is sacrificed in temperature at all. It may also be much less hydroscopic than Ferrous Sulfate judging by its solubility which is another bonus.
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draculic acid69
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[*] posted on 27-7-2020 at 16:37


I read somewhere that ammonium persulfate decomposes about 180'c evolving so3
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Eddie Current
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[*] posted on 27-7-2020 at 16:56


SO2 & SO3 @ 120C

http://journals.sagepub.com/doi/pdf/10.1080/1091581015263071...

More info:

https://www.peroxychem.com/media/241528/pxc089_persulfates_b...



[Edited on 28-7-2020 by Eddie Current]
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[*] posted on 28-7-2020 at 05:39


The Orginal Vitriol stoves used Fe(SO4)3. They made it when they rostet the FeSO4 in flat pans with stiring.

2 Parts Calciumsulfat and 1 part Ironoxide decompuses to SO3. When you ad
Calciumflourid and Calciumchlorid the tmperatur is much lower but it corrodes your reaction fessel away.

http://dingler.culture.hu-berlin.de/article/pj255/mi255mi02_...

There is an interesting reaktion with Sodiumhydrogensulfate and Magnesiumsulfate it forms an double salt and releases SO3 at araound 300C.

http://dingler.culture.hu-berlin.de/article/pj230/mi230mi05_...

all links only in german
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macckone
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[*] posted on 28-7-2020 at 10:34


Google translate - first link:

For the extraction of sulfuric acid from gyps.

If, according to Scheurer-Kestner ( Comptes rendus, 1884, vol. 99, p. 876), you heat 2 parts of calcium sulfate and 1 part of iron oxide to a bright red heat, all the sulfur is expelled. First sulfuric anhydride escapes, later sulfuric acid and oxygen. The addition of chlorine and fluorine calcium significantly reduced the decomposition of the sulfuric anhydride, but the crucibles were not sufficiently resistant to such a mixture. If the sulfates of other divalent metals are used instead of the calcium sulfate, the action is similar, for example when using lead sulfate, which seems to react at a lower temperature than the calcium sulfate, as well as magnesium sulfate.

second link:


For the production of sulfuric acid.

EA Parnell in Swansea (DRP No. 1351 of September 8, 1877) suggests mixing zinc blending with a concentrated zinc vitriol solution in the ratio that 1 mol. ZnS 3 mol. ZnS 4 O 4 comes to dry the mass in to heat in a closed furnace and to process the developed sulfurous acid in lead chambers for sulfuric acid.

JAW Wolters in Kalk (DRP No. 3110 of March 5, 1878) made the proposal to carefully melt anhydrous acidic sulfuric acid sodium, then mix it with the appropriate amount of anhydrous magnesium sulfate and heat it a little more, after which sulfuric anhydride distilled over. The remaining double compound of sulfuric acid sodium magnesium is broken down in a known manner by water and can be used again and again to represent the anhydride. It is emphasized as a particular advantage that the distillation of the sulfuric anhydride takes place at a relatively low temperature, so that the apparatus used is little used.
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[*] posted on 28-7-2020 at 10:44


Problem with this idea is that it's generally not effective or safe to lead SO3 gas into water with which it reacts so exothermically the product is inevitably a gas or "fine mist" as often put. Pyrolysis of alkali pyrosulfates is otherwise quite easy and would be a good choice for this process.

You can of course perform the traditional process -- lead SO3 into conc. H2SO4 giving oleum, then combine with water (always add acid to water) and repeat -- but this requires some conc. H2SO4 to start with.

In order to get this, one particularly nice way is to use the disproportionation of KHSO4 in water/ethanol mixtures from which it precipitates K2SO4. Sodium or ammonium bisulfates may also work. The resulting dilute sulfuric acid can be concentrated by heating.




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 28-7-2020 at 13:31


True, but i think the biggest issue is that the reason sulfuric acid is so useful is largely that it is so cheap. If you have to produce it by pyrolysing metal sulfates it loses much of its utility as a reagent though it is still useful as a catalyst. This may be a good way to produce Sulfur Trioxide or concentrate acid though.
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[*] posted on 28-7-2020 at 15:28


Informations about thermal properities of inorganic sulfates can be found (for example), here:
"Thermal Decomposition of Ionic Solids" (1999)
or
"High Temperature Properties and Decomposition of Inorganic Salts. Part 1. Sulfates" (1966)
Available from internet, for free :D




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[*] posted on 5-8-2020 at 18:05


This 2009 article,"Acid production by FeSO4·nH2O dissolution and implications for terrestrial and martian aquatic systems" may be of interest for terrestrial humans searching locally for an acid salt mixture as a substitute for H2SO4. To quote from the cited source:

"Combined experimental, modeling, and analytical results indicate that the rapid acidification of dilute waters in contact with nominally Fe2+-sulfate minerals (FeSO4nH2O) is caused by Fe3+ hydrolysis, which occurs when low levels (<1 mol%) of a contaminant Fe3+-sulfate phase are dissolved along with the FeSO4nH2O. This rapid acidification has previously been attributed to hydrolysis by Fe2+. However, dissolution experiments performed using ZnSO4nH2O, in which the Zn2+ cation has a higher hydrolysis constant (log K = −8.96) than Fe2+ (log K = −9.5), failed to produce significant changes in solution pH. We present the results of geochemical modeling simulations confirming that FeSO4nH2O dissolution alone cannot explain the experimentally observed change in pH from 5.65 to 3.50. Nor can the experimental observations be explained by oxidation of Fe2+ to Fe3+ in solution. Instead, our experimental results can be best explained by modeling the incorporation of <1 mol% Fe3+ contamination from any number of Fe3+ or mixed valence Fe-sulfate phases..."

Link: https://pubs.geoscienceworld.org/msa/ammin/article-abstract/...

So, perhaps an aqueous mix of FeSO4·nH2O and Fe3+ would work as a substitute for dilute H2SO4 in select applications.

Further, I speculate that pumping air into the aqueous FeSO4 plus an added acid source (like NaHSO4, where the purpose is to avoid the creation of a basic ferric sulfate) may result in an acid mixture with pH < 4 per above. That is, supplying a source of Fe3+, where the expected electrochemical action of oxygen on ferrous is given by:

4 Fe2+ + O2 + 2 H+ -> 4 Fe3+ + 2 OH-

Source: See discussion and links at https://www.sciencemadness.org/whisper/viewthread.php?tid=15...

Or, as sometimes presented, albeit less correctly from an operational perspective, in my opinion:

4 Fe2+ + O2 + 4 H+ -> 4 Fe3+ + 2 H2O

where adding acidic NaHSO4 reduces the likelihood of an insoluble basic ferric sulfate (as can occur in practice when employing the electrochemical reaction cited above). I suspect adding the NaHSO4 is required to achieve the pH effect per the statement above: "Nor can the experimental observations be explained by oxidation of Fe2+ to Fe3+ in solution", where the authors may not have a complete understanding in the implicit electrochemical reaction.

However, if the above embodiment is not successful, the author does further mention the employment of anhydrous Fe2(SO4)3 (surface chemistry?). So this may entail, per the electrochemical route specified above, obtaining a dry ferric salt, but the presence of Na2SO4 (OH- acting on NaHSO4) may preclude this. So, forget the acidification step above (or perhaps add MgSO4, albeit only slightly acidic, see https://chemistry.stackexchange.com/questions/49825/is-magne... ), and just try to obtain dry crystals of Fe2(SO4)3 and/or its basic salt.

Proposed chemistry with MgSO4 variant (which may create an interesting white precipitate of Mg(OH)2) with no added acid, just water with the purpose of obtaining eventually a dry ferric soluble sulfate, after filtering out, or waiting for the Mg(OH)2 to settle out, and decant:

4 FeSO4 (aq) + O2 (d) + 2 MgSO4 (aq) -> 2 Fe2(SO4)3 (aq) + 2 Mg(OH)2 (s)

where the MgSO4 should not be in stoichiometric excess.

A seemingly proposed weird and inventive experiment based on somewhat vague (or is that unexplained) reported chemistry accounting for a pH drop.

[Edited on 6-8-2020 by AJKOER]
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[*] posted on 6-8-2020 at 14:18


Quote: Originally posted by AJKOER  
...rubbish...

I am sorry, forum Einstein, but one more time you have proved you are an idiot.




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[*] posted on 8-8-2020 at 04:07


KmnO4:

Thanks for the feedback. I have since updated my mechanic on the formation of rust thread (at https://www.sciencemadness.org/whisper/viewthread.php?tid=15... ) with an advanced source (‘Thermodynamics of Oxidation of Iron and Carbon Steels in Water', a technical appendix with a mere 79 pages and 20 cited references, freely available as a pdf at https://link.springer.com/content/pdf/bbm%3A978-90-481-3477-... , which I would recommend if you need some 'put me to sleep material').

However, as this source apparently has the tumerity to seemingly support my supposition on reaction mechanics, you must be right (pure rubbish!).

By the way, I do find the following proposed reaction quite interesting:

4 Fe2+ (aq) + O2 (d) + 2 MgSO4 (aq) -> 4 Fe3+ (aq) + 2 Mg(OH)2 (s) + 2 SO4(2-)

and I plan on trying it with a more interesting and workable variation starting with elemental iron:

2 Fe (s) + O2 (g) + 2 MgSO4 (aq) --> 2 FeSO4 (aq) + 2 Mg(OH)2 (s)

that is, will iron actually rust in air in the presence of slightly acidic MgSO4 with the creation of FeSO4 (and even perhaps ferric) and an obvious precipitate of Mg(OH)2?

Cool if it does, and I will report back with pictures!

Note: Anyone performing this reaction, please avoid sunlight as any formed Mg(OH)2/MgO is likley photocatalytic potentially adding further chemistry.

[Edit] As a matter of couresty, as I have already started experimenting based on this electrochemical based reaction equation, I will post results/pictures in a new thread.

[Edited on 8-8-2020 by AJKOER]
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[*] posted on 15-8-2020 at 01:36


In the old days, i mean really old, they made sulfuric acid from Ferrous Sulfate by heating it until it decomposed.
I have no details though and there probably is many better alternatives.

It would be nice if there were a relatively easy way of making sulfuric acid as the politicians are on the path of banning every compound that can be little dangerous making hobby experimenters life really hard.

[Edited on 2020-8-15 by Mateo_swe]
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[*] posted on 15-8-2020 at 13:51


Overheated Steam on Magnesiumsulfate but i think you only get deluted acid.
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macckone
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[*] posted on 17-8-2020 at 19:27


Quote: Originally posted by Alkoholvergiftung  
Overheated Steam on Magnesiumsulfate but i think you only get deluted acid.

Yes, that was covered, the heat is 1080C, maybe slightly lower with the steam vs direct heat.
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