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Author: Subject: Small scale production of H2SO4 in the amateur lab
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[*] posted on 15-8-2020 at 13:21
Small scale production of H2SO4 in the amateur lab


Sulfuric acid is hard to get your hands on in certain countries and there also is upcoming regulation in the EU to limit its availability for the general public. For home chemistry purposes, this chemical, however, is so basic and of such great importance, that many things will become impossible, or at least very hard, if this chemical is not available anymore. For this reason we start a sticky thread on small scale production of H2SO4. There already is a thread on the lead chamber process, but that is just one specific method of making sulfuric acid, which may not be suitable for everyone.

In this thread any serious method can be discussed. Even better would be descriptions of practical attempts, even if they failed. Learning from failures can lead to future attempts with more success.

In order to limit the number of sticky threads, the lead chamber thread will not be sticky anymore, but it can be found directly by means of the above link.




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[*] posted on 15-8-2020 at 18:34


One of my first experiments was sulphuric acid by electrolysis of copper sulphate solution....I used carbon rods from batteries which discoloured it but seemed to filter out ok. Im unsure how scaleable it is and I also am unsure what concentration can be achieved before boiling it down to concentrate.
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[*] posted on 15-8-2020 at 21:18


Because of hazard shipping issues, it is practically impossible to find or purchase sulfuric acid in my state. In the past I have used a sulfur dioxide generator with 15 percent hydrogen peroxide, although this process is a lot of work for very little acid.
One method that I do not see talked about very often is the reaction of oxalic acid and iron (iii) sulfate in aqueous solution. Both reagents are cheap, if you can find oxalic acid, and should theoretically produce iron free acid due to the extreme insolubility of iron oxalate. The solution can then be boiled down or distilled to produce concentrated acid.
This method should be much easier then the SO2 and should produce more acid at once due to the convenience of not needing a gas generator.
I personally have not tried this method yet, but I have oxalic acid coming in the mail during the next few weeks and will share some observations.




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[*] posted on 15-8-2020 at 21:28


Quote: Originally posted by nzlostpass  
One of my first experiments was sulphuric acid by electrolysis of copper sulphate solution....I used carbon rods from batteries which discoloured it but seemed to filter out ok. Im unsure how scaleable it is and I also am unsure what concentration can be achieved before boiling it down to concentrate.


I did something similar. I did refine the process to something quite manageable. I will post it here later: including some photos.
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[*] posted on 16-8-2020 at 02:19


I'm seriously considering buying a 25kg bag of sulfur. I can get it in 1/2kg quantities from a home gardening supplier, but at a obscene price. Even including shipping the bag will be cheaper than 2kg from home garden supplier.
I just need to convince a farming friend that he won't get on any terrorist watch list by ordering it. Shouldn't be a problem, as he's producing ecological food and sulfur is approved for this. It actually looks like a very versatile compound that can be used both for soil improvement and as a fungicide.

The process will most likely be the contact process. Vanadium is the typical catalyst, but IIRC platinum could also be used. It's said to be sensitive to arsenic, but fertilizer grade sulfur should be pretty low in that so it might be worth a look. Perhaps a broken automotive catalyst could work?




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[*] posted on 16-8-2020 at 17:01


NurdRage have few videos on how to make sulfuric acid:

Copper(II) chloride method

Electrobromine process

Electrolysis of copper(II) sulfate

SO2/H2O2 method




If you are interested in aqueous inorganic chemistry look at https://colourchem.wordpress.com/main-page/

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[*] posted on 17-8-2020 at 06:05


A suggested route to very dilute H2SO4 based on a Wikipedia commentary on aluminum sulfate (https://en.wikipedia.org/wiki/Aluminium_sulfate ), to quote:

"When dissolved in a large amount of neutral or slightly alkaline water, aluminum sulfate produces a gelatinous precipitate of aluminium hydroxide, Al(OH)3. In dyeing and printing cloth, the gelatinous precipitate helps the dye adhere to the clothing fibers by rendering the pigment insoluble...Aluminium sulfate is sometimes used to reduce the pH of garden soil, as it hydrolyzes to form the aluminium hydroxide precipitate and a dilute sulfuric acid solution."

This method requires at least 3 steps, first obtaining or preparing aluminum sulfate, hydrolyzing it followed by a concentration step.

As to the first step, I was thinking of microwave heating of home available annealed Al foil in aqueous NaHCO3, as a pre-activation treatment. Rinse off the treated Al metal and place in a concentration solution of CuSO4. Copper should plate out leaving a solution of aluminum sulfate.

An interesting alternate path requiring some investigation would be first prepare Al(OH)3 and add the freshly prepared salt to albeit only slightly acidic aqueous MgSO4. May not react, or react slowly, or even may selectively react depending on the Al(OH)3 prep path (where high surface area, poor crystallinity of the Al(OH)3, small particile size could be promoting factors, see, forexample, relatedly https://www.researchgate.net/publication/257224933_Effect_of...), or possibly result in a Mg-Al salt (see https://www.sciencedirect.com/topics/chemistry/magnesium-hyd... ).

The last step is futher dilute the obtained aluminum sulfate and followed by boiling to concentrate.

[Edited on 17-8-2020 by AJKOER]
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[*] posted on 17-8-2020 at 06:31


Quote: Originally posted by AJKOER  
A suggested route to very dilute H2SO4 based on a Wikipedia commentary on aluminum sulfate (https://en.wikipedia.org/wiki/Aluminium_sulfate ), to quote:

"When dissolved in a large amount of neutral or slightly alkaline water, aluminum sulfate produces a gelatinous precipitate of aluminium hydroxide, Al(OH)3. In dyeing and printing cloth, the gelatinous precipitate helps the dye adhere to the clothing fibers by rendering the pigment insoluble...Aluminium sulfate is sometimes used to reduce the pH of garden soil, as it hydrolyzes to form the aluminium hydroxide precipitate and a dilute sulfuric acid solution."

This method requires at least 3 steps, first obtaining or preparing aluminum sulfate, hydrolyzing it followed by a concentration step.

As to the first step, I was thinking of microwave heating of home available annealed Al foil in aqueous NaHCO3, as a pre-activation treatment. Rinse off the treated Al metal and place in a concentration solution of CuSO4. Copper should plate out leaving a solution of aluminum sulfate.

An interesting alternate path requiring some investigation would be first prepare Al(OH)3 and add the freshly prepared salt to albeit only slightly acidic aqueous MgSO4. May not react, or react slowly, or even may selectively react depending on the Al(OH)3 prep path, or possibly result in a Mg-Al salt.

The last step is futher dilute the obtained aluminum sulfate and followed by boiling to concentrate.

[Edited on 17-8-2020 by AJKOER]


Hydrolysis of aluminium sulfate is more complicated than that.

[Al(H2O)6]3+ + H2O <--> [Al(H2O)5(OH)]2+ + H3O+

[Al(H2O)5(OH)]2+ + H2O <--> [Al(H2O)4(OH)2]+ + H3O+

[Al(H2O)4(OH)2]+ + H2O <--> [Al(H2O)3(OH)3] + H3O+

[Al(H2O)3(OH)3] complex can react with other molecules of [Al(H2O)3(OH)3] to form polymeric species.

All positively charged complexes are soluble in water.

The major complex is [Al(H2O)5(OH)]2+ followed by [Al(H2O)4(OH)2]+ and the minor complex is [Al(H2O)3(OH)3].

When I prepared saturated solution of potassium alum I never saw any formation of aluminium hydroxide (even after few months). Aluminium sulfate may behave differently, but I doubt that huge amounts of aluminium hydroxide are formed.

So this will be very inefficient way how to prepare sulfuric acid, because it will be very very dilute and highly contaminate with aluminium.

--------------------------------------------------------------------------------

I had a few ideas yesterday and I'll try it when I will have some time:

Oxidation of sulfur by hot conc. HNO3

Oxidation of sulfur by Fenton's reagent (which requires only catalytic amount of iron)

[Edited on 17-8-2020 by Bedlasky]




If you are interested in aqueous inorganic chemistry look at https://colourchem.wordpress.com/main-page/

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[*] posted on 17-8-2020 at 06:57


By far the simplest method I have heard of is to use the disproportionation of potassium bisulfate in ethanol:
http://www.sciencemadness.org/talk/viewthread.php?tid=79548

You can just boil down the filtrate to obtain reasonably concentrated sulfuric acid! It is still effective but less efficient and more "sticky" when using the sodium salt. The possibility of using NH5SO4 has also occurred to me but not been tested (I'll get to it someday...).

This method is also good for making a "seed" quantity of sulfuric acid that can then be treated with SO3 to give oleum and diluted etc. Using both methods together you can make lots of sulfuric acid from just bisulfate, ethanol and water -- but the intermediacy of SO3 is always a little scary.

[Edited on 17-8-2020 by clearly_not_atara]




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 17-8-2020 at 07:12


Quote: Originally posted by Fulmen  
I'm seriously considering buying a 25kg bag of sulfur. I can get it in 1/2kg quantities from a home gardening supplier, but at a obscene price. Even including shipping the bag will be cheaper than 2kg from home garden supplier.
I just need to convince a farming friend that he won't get on any terrorist watch list by ordering it. Shouldn't be a problem, as he's producing ecological food and sulfur is approved for this. It actually looks like a very versatile compound that can be used both for soil improvement and as a fungicide.

The process will most likely be the contact process. Vanadium is the typical catalyst, but IIRC platinum could also be used. It's said to be sensitive to arsenic, but fertilizer grade sulfur should be pretty low in that so it might be worth a look. Perhaps a broken automotive catalyst could work?


Do you have any Tractor Supply stores, Rural King or similar stores near you - or even any feed &/or seed stores? All of them should have "sulfur Flour" in 50lb/25kg bags for very good prices. I think maybe even pottery supply shops might have it as well, but I'm not sure, it may have been only available at a few places I looked.
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[*] posted on 17-8-2020 at 09:25


Nah. I ain't living in the US, and there simply isn't any shops selling anything useful anywhere close.



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[*] posted on 17-8-2020 at 12:32


If you can access epsom salt and plaster then you can also electrolyze magnesium sulfate with a plaster separator.
Sodium sulfate can also be electrolized in a similar chamber.
The electrolytic methods are less than ideal and produce weak acid with impurities for a lot of work.
The copper sulfate method is the easiest.

You can also superheat gypsum to high temperatures (1080C) to release sulfur dioxide.
Burning sulfur is going to be easier though possible not cheaper.
Burning DMSO is also a possibility as DMSO is available for use on horses.

Sulfur dioxide can be catalyzed to sulfur trioxide on a vanadium oxide catalyst or via a chamber method.

Both have been done by amateur chemists.
The chamber method is best done on a small scale using 5 gallon water bottles.
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[*] posted on 17-8-2020 at 13:48


I would still go for the bisulfate/pyrosulfate route as suggested by Atara.



Attachment: SO3_and_oleum.pdf (198kB)
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[*] posted on 18-8-2020 at 09:18


It sure would be nice if there was a pretty easy process for sulfuric acid preparation.
We might be able to buy it now but there is already politicians deciding that the common citizen should not have access to any chemical that is dangerous or can be used to make anything dangerous.
The ban is already in motion so any DIY way of preparing chemicals is good to have knowledge about.

I have tried the oxalic acid + Ferrous Sulfate method and it works.
FeSO4 + H2C2O4 --Heat --> FeC2O4 (s) + H2SO4
There is a thread about this method around here somewhere if i remember correctly.
If there is an easy way of making cation and anion membranes (maybe there is), there is a good way of making sulfuric acid and sodium hydroxide from sodium sulfate in a 3 compartment cell.
In the old days they made sulfuric acid by heating Ferrous Sulfate until it decomposed.
It will need a high temp oven to experiment with this method.
Since Ferrous Sulfate is cheap in 25kg bags (32 Euros on ebay + shipping), methods using this is nice.
I add some pdfs about preparation of sulfuric acid.

Attachment: chemicals_sulfuric_acid_preparation_from_oxalic_acid_thompson1849.pdf (209kB)
This file has been downloaded 65 times

Attachment: h2so4_vogel_eng_H2SO4_from_Oxalic_acid_and_Iron_sulphate.pdf (255kB)
This file has been downloaded 63 times

Attachment: US5928488 - Electrolytic Sodium sulfate salt splitter comprising a polymeric ion conductor.pdf (981kB)
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[*] posted on 18-8-2020 at 15:08


For preparation of copper sulfate
Start with ammonium hydroxide and magnesium sulfate
To form ammonium sulfate excess ammonia then add to copper oxide prepared by heating copper till it blackens and flakes off when it cools strong heating to drive off the ammonia from the complex to form copper sulfate




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[*] posted on 19-8-2020 at 13:37


Clearly Not Atara:

If I am reading a referenced source on salting-out correctly, NH4+ > K > Na....

So, heating (but below 350 C, which is a cited decomposition temperature forming sulfur oxides) ammonium sulfate (from, for example, the action of NH3 on aqueous Epsom's Salt, MgSO4) to create ammonium bisulfate (see http://www.sciencemadness.org/talk/viewthread.php?tid=198 ) and adding a select alcohol, may work as well.

There may also be other paths starting with ammonium sulfate, and I am looking at one now.

[Edited on 19-8-2020 by AJKOER]

[Edited on 20-8-2020 by AJKOER]
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[*] posted on 19-8-2020 at 14:22


I believe you can produce oleum by simply heating sodium bisulfate until it decomposes to sodium pyrosulfate and then heating the sodium pyrosulfate until it decomposes. It looks easy, but the high temperatures required cannot be safely achieved in ordinary borosilicate glass, and of course anything involving oleum is extremely hazardous. Len1 discussed the procedure in some detail in his book, Small-Scale Syntheses of Laboratory Reagents.

I could possibly try it out next weekend, but I'm not sure what to do with the oleum afterwards, and I'd have to do it outside, preferably in a location that won't terrorize the local homeless population... anyway, I ordered some sodium bisulfate.

oleum_production.png - 178kB
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[*] posted on 22-8-2020 at 03:32


Do you think Sodium metabisulfite (Na2S2O5) could be used as a replacement for the Sodium bisulfate in the process?
It is found in brew shops for ok prices.

[Edited on 2020-8-22 by Mateo_swe]
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[*] posted on 22-8-2020 at 05:29


JJay, have a look at the pdf I attached above. Here the decomposition is done in borosilicate and persulfate is cheap and available as etchant for circuit boards.
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[*] posted on 22-8-2020 at 06:03


Quote: Originally posted by Tsjerk  
JJay, have a look at the pdf I attached above. Here the decomposition is done in borosilicate and persulfate is cheap and available as etchant for circuit boards.


I saw it. The major drawback I see is that it requires anhydrous sulfuric acid as one of the starting materials. Also, sodium persulfate is much more expensive than sodium bisulfate. I considered giving it a try (and why not just pyrolyze the cheaper sodium bisulfate instead of sodium persulfate?), but I think the strongest sulfuric acid I have is 93%, so I'd like to know what the effect on yields is if there is a little bit of water present.
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[*] posted on 22-8-2020 at 06:05


Quote: Originally posted by Mateo_swe  
Do you think Sodium metabisulfite (Na2S2O5) could be used as a replacement for the Sodium bisulfate in the process?
It is found in brew shops for ok prices.

[Edited on 2020-8-22 by Mateo_swe]


It could be oxidized to sodium bisulfate, but sodium bisulfate is usually cheaper and easier to find.
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[*] posted on 23-8-2020 at 08:53


Bedlasky:

Thanks for your comment, as I agree, the hydrolysis of Aluminum sulfate with the 'acid' water to create a separable amount of very dilute H2SO4 is unlikely.

However, perhaps working with the slightly more powerful CO2 (aq) composition, leading to an insoluble salt product, may be worth exploring as a path to some very dilute H2SO4 itself, based on a Science Direct commentary. More precisely, the cited reaction, occurring in natural waters, of the bicarbonate ion acting on alum resulting in a precipitate of Al(OH)3 per the cited reaction (from a Science Direct compilation available at: https://www.sciencedirect.com/topics/chemistry/aluminium-sul...):

Al2(SO4)3.xH20 + 6 HCO3- (aq) <--> 2 Al(OH)3 (s) ....

but in a new suggested embodiment, exploring the use of slightly acidic carbonic acid. Note: the reaction is reversible, but using a reaction chamber, like a test tube, where the surface area contact of the created Al(OH)3 precipitate is reduced, may be beneficial and avoid stirring.

Did find a supporting reference (https://thefishsite.com/articles/managing-high-ph-in-freshwa... ) relating to care for treating pH problems in fresh water ponds, with naturally contained dissolved CO2 from air, to quote:

"Overtreatment with alum can cause a dramatic decrease in pH, possibly to levels more dangerous than the original high pH problem...A safer, longer lasting way to reduce high pH is to add carbon dioxide, which acts as an acid in water...The careful use of alum is probably the safest and most dependable emergency treatment."

Also, a more dated reference (at https://www.jstor.org/stable/pdf/41224990.pdf) claims that optimal efficiency relating to coagulation (Al(OH)3 formation) occurs not in alkaline pH, but at around pH 5.5 (which is also cited in this 1989 thesis at https://scholars.unh.edu/cgi/viewcontent.cgi?article=2599&am... ). Further, the addition of acid assists (like H2SO4 or other), and that CO2 is not particularly good. On the latter point to quote another source (https://drinking-water.extension.org/drinking-water-treatmen...):

"Citric acid and alum can be used instead, although they are more expensive."

The chemistry associated with the use of agents other than CO2 in transition metal and oxygen rich natural waters, however, may be related to the presence of active redox reactions and also, per the cited thesis (Pages 39-40), aluminum polymer formation per the presence of fulvic acid and the electrostatic properties of colloidals resulting in particle aggregation. But more insightful per a 2019 work (https://www.researchgate.net/profile/Wenzheng_Yu/publication...), to quote:

"When monomeric Al (alum) was added together with kaolin at pH 7, some [Al(OH2)(6-n)(OH)n](3-n)+ absorbed on the surface of kaolin, causing the surface of kaolin to be fully covered with alum precipitates/nanoparticles (Duan and Gregory, 2003). Then the flocculation process started."

So, hydrated Al(OH)3 can be induced (with select acids, kaolin based clay,...) to introduce deposits of varying charged surface complexes to assist in particle aggregation.

Other redox reactions in the presence of Fe ions are also possible. Further, with charged colloidal particles, solvated electron creation is possible, which may result, in the presence of O2 again, the creation of a superoxide radical anion:

O2 (d) + e- (aq) = .O2- (or .HO2 at pH < 4.88)

These radical species have been reported to interact with Al3+, as I have noted previously on SM with references discussed at https://www.sciencemadness.org/whisper/viewthread.php?tid=96... , and further promote redox chemistry.

On the role of acids and Al speciation, note the following source https://www.researchgate.net/figure/Monomeric-and-polymeric-... , to quote:

"Numerous studies have reported that there are various Al forms in soils in monomer, polymer or solid phase, and that their concentration depends on the degree and duration of hydrolysis of the Al compounds (Delhaize and Ryan, 1995) (Figure 1). Rout et al . (2001) found a significant correlation between low pH and high concentrations of phytotoxic Al species, which is related to the reduction of exchangeable bases in the soil solution (Mora et al ., 2006). Soil acidification is associated with inappropriate agricultural practices (Rengel, 1996), heavy winter precipitation that causes the loss of bases (Na + , K + , Ca 2+ , Mg 2+ ) due to leaching (Mora et al ., 2006), use of ammoniacal fertilizer (urea) and nutrient uptake by plants (Mora et al . , 2006)."

Still, assuming dilute H2SO4 is after all present, can the weak acid be effectively concentrated here?

The suggested theoretical investigation reminds me of an Oxalate acid path, but with a much weaker starting acid.

[Edited on 24-8-2020 by AJKOER]

[Edited on 24-8-2020 by AJKOER]
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[*] posted on 24-8-2020 at 07:16


Thermal decomposition of gypsum to sulfur dioxide for use conversion to sulfuric acid by chamber or contact process.
See attached file.

Attachment: Decomposition of Calcium Sulfate.pdf (3.3MB)
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[*] posted on 24-8-2020 at 13:19


Müller-Kühne process for Sulfuric acid and cement. works with an rotation stove and an mix of gypsum clay and sand the temperatur is around 700C.
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[*] posted on 24-8-2020 at 15:46


Alkoholvergiftung,
Müller-Kühne process requires a final temperature of 1150-1250C to release most of the SO2 and carbon is also utilized.
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