Sciencemadness Discussion Board - Small scale production of H2SO4 in the amateur lab

Small scale production of H2SO4 in the amateur lab

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woelen - 15-8-2020 at 13:21

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.

nzlostpass - 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.

Abromination - 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.

j_sum1 - 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.

Fulmen - 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?

Bedlasky - 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

AJKOER - 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]

Bedlasky - 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]

clearly_not_atara - 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]

RogueRose - 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.

Fulmen - 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.

macckone - 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.

Tsjerk - 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)
This file has been downloaded 1082 times

Mateo_swe - 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 967 times

Attachment: h2so4_vogel_eng_H2SO4_from_Oxalic_acid_and_Iron_sulphate.pdf (255kB)
This file has been downloaded 943 times

Attachment: US5928488 - Electrolytic Sodium sulfate salt splitter comprising a polymeric ion conductor.pdf (981kB)
This file has been downloaded 876 times

symboom - 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

AJKOER - 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]

JJay - 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.

Mateo_swe - 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]

Tsjerk - 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.

JJay - 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.

JJay - 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.

AJKOER - 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]

macckone - 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)
This file has been downloaded 1029 times

Alkoholvergiftung - 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.

macckone - 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.

Deleted post

Sulaiman - 25-8-2020 at 07:12

I deleted a method that was inappropriate for this thread
Sorry.

[Edited on 25-8-2020 by Sulaiman]

Fulmen - 25-8-2020 at 07:22

That sounds exceedingly complicated and dangerous compared to simple electrolysis of CuSO4.

JJay - 27-8-2020 at 22:26

I'm not 100% sure if I'm going to be able to find and unpack everything I need to do a trial run of oleum generation this weekend... my chemicals and equipment are packed in boxes and totes. The chemicals are well organized, but the equipment is not....

So anyway, I was putting together a list of what I would need earlier, and I drew a blank at gloves... what kind of gloves do people use for handling oleum? I would expect it to destroy almost any gloves immediately, perhaps setting them on fire.

Butyl rubber gloves are recommended for sulfuric acid according to: https://www.augusta.edu/services/ehs/chemsafe/PDF%20files/gl...

I've seen nitrile gloves recommended for sulfuric acid by other sources. I typically just wear latex gloves for sulfuric acid, but it does slowly attack them, so I try not to let any chemicals get on them, and I change gloves often.

Of course, oleum is in a whole different echelon than regular sulfuric acid. I wouldn't consider wearing latex, butyl, or nitrile gloves to handle oleum.

Viton gloves would be ideal, but they are extremely expensive, and I don't have time for shipping....

plastics - 28-8-2020 at 04:03

Shoot me down if you like but don’t bother wearing gloves

I didn’t when I made SO3/oleum from pyrolysis of sodium bisulphate in some quantity

Nothing encourages careful handling better than exposed flesh

As you say the stuff destroys most glove materials

Keep a bucket of water and ice handy - It is quite unexciting pouring oleum onto ice - I know I’ve done it

If the stuff was toxic rather than corrosive - yes I would wear gloves

Belowzero - 28-8-2020 at 06:25

Thanks Woelen for creating this thread, as you said in the U2U it would be awesome if this can lead to a workable and practical method.

I collected previous threads on the topic (and related threads such as SO3 production)

This is the first part.
The most important and extensive threads:

Axehandle posted this thread (2004)
H2SO4 by the Lead Chamber Process - success
https://www.sciencemadness.org/whisper/viewthread.php?tid=28...

garage chemist posted this in Prepublication (2008):
Sulfur trioxide from sodium hydrogen sulfate
https://www.sciencemadness.org/whisper/viewthread.php?tid=10...
And
Sulfur trioxide Oleum: the ferric sulfate method
https://www.sciencemadness.org/whisper/viewthread.php?tid=10...

len1 posted this in prepublication (2008):
Subject: Sulphur Trioxide and Oleum Using a Box Oven - Illustarted Practical Guide
http://www.sciencemadness.org/talk/viewthread.php?tid=10332

Microtek posted this thread (2017):
Lab scale contact process for H2SO4 (report)
https://www.sciencemadness.org/whisper/viewthread.php?tid=72...

Other useful and related threads:

Subject: oleum & SO3
http://www.sciencemadness.org/talk/viewthread.php?tid=727

Subject: SO3 from CaSO4 is it possible ?
http://www.sciencemadness.org/talk/viewthread.php?tid=15285

Subject: SO3 from pyrosulfate and H2SO4
http://www.sciencemadness.org/talk/viewthread.php?tid=63569

Subject: Sulfur Trioxide by m-Phosphoric Acid route (success)
http://www.sciencemadness.org/talk/viewthread.php?tid=65358

[Edited on 29-8-2020 by Belowzero]

macckone - 28-8-2020 at 09:28

This thread covers oxalate methods previously mentioned.
www.sciencemadness.org/talk/viewthread.php?tid=18963

They are much simpler than other methods if you have access to oxalic acid.

The easiest ways to get oxalic acid are as wood bleach (likely to be impacted by REACH), nitric acid oxidation of sucrose (birkeland-eyde reactor for nitric acid), oxidation of ethylene glycol (pick your oxidizer), and rhubarb leaves

CharlieA - 28-8-2020 at 16:50

Quote: Originally posted by plastics

Shoot me down if you like but don’t bother wearing gloves...

Nothing encourages careful handling better than exposed flesh...

... If the stuff was toxic rather than corrosive - yes I would wear gloves

Right on about exposed flesh!! I sometimes think that one can get careless/over-confidant with too much personal protective equipment (PPE). Personally, when handling corrosive material, I would wear proper safety goggles and/or a face shield before I would worry about gloves.
And after seeing too many videos with people in short pants and short sleeves, that is certainly "exposed flesh". A lab coat is cheap PPE.

Alkoholvergiftung - 28-8-2020 at 23:50

From Gyps and Lead and HCL.

http://dingler.culture.hu-berlin.de/article/pj139/ar139070

The invention of removing the sulfuric acid from the gypsum consists in using an intermediate body for this purpose, the sulfuric acid lead oxide, which gives off its sulfuric acid against hydrochloric acid with great ease, thereby transforming itself into chlorine lead, which then again with the help of finely ground gypsium is converted into sulfuric acid lead oxide.

If you pour over 1 1/4 equivalent of sulfuric acid lead oxide with 1 equivalent of hydrochloric acid at 21 ° Baumé, and heat both with stirring to about 50 to 60 ° Reaumur, the liquid is almost completely transformed into sulfuric acid after a short time. If you use e.g. 6 equivalents of sulfuric acid lead oxide and if 5 equivalents of hydrochloric acid at 21 ° Baumé are added, an acid is obtained in which the ratio of sulfuric acid to hydrochloric acid is 5: 1.

The more concentrated the hydrochloric acid with which one works, the more favorable this relationship is.

The exchange of hydrochloric acid for sulfuric acid takes place in quite different proportions, depending on how the hydrochloric acid is allowed to act in different strengths on the sulfuric acid lead oxide. Since the chlorine lead takes up a much larger space than the sulfuric acid lead oxide; as this salt generally retains a large part of the sulfuric acid obtained like a sponge, this fact provides a means of separating large masses of the sulfuric acid obtained from the lead chloride without filtration.

My method of making these facts useful in practice consists in setting up a wooden vessel lined with lead, depending on the scale in which the work is to be carried out, and which can be heated by means of steam through lead pipes. In this vessel, regardless of the shape, I put a certain amount | 284 | sulfuric acid lead oxide, which is then poured over with a mixture of sulfuric acid and hydrochloric acid or sulfuric acid alone from 18 ° Baumé until the salt is saturated or soaked with it. This weak acid is then always recovered in the course of the work, as will be shown below. It has no other purpose than to soak the lead salt with a liquid which does not change it and enables me to draw off the same quantity of acid which I give up. With diligent stirring and warming, as much hydrochloric acid of 21 ° Baumé is added as calculated for the first operation. After a while the liquid is left to stand still, to settle, and then just as much weak sulfuric acid is drawn off as hydrochloric acid was added to the measure. Furthermore, about half of the hydrochloric acid previously used is poured on for the second time, proceeding as the first, whereby all the sulfuric acid lead oxide is then converted into chlorine lead. After this acid has also been drawn off from the residue and combined with the first obtained, it is evaporated in vaulted lead pans up to 60 ° Baumé; but the escaping hydrochloric acid is collected by means of known apparatus.

By means of these operations, as a rule 3/4 of the sulfuric acid contained in the lead oxide sulfuric acid is immediately obtained for evaporation. If the correct quantity of water is now added to the chlorine lead, the quantity of weak acid which had to be artificially produced during the first operation is restored. When the lead salt settles, this weak acid is pumped onto a higher standing vessel and stored for further use. Further washing once or twice removes as much acid as possible from the lead chlorine; But so that this weak acid is not lost, it is suggested instead of ordinary water in the preparation of Glauber's salt and is saturated again with hydrochloric acid gas.

After the chlorine lead has been sufficiently washed out, the calculated quantity of Gyps is added in a very finely ground state with such a quantity of water that no stronger lye than 2 1/2 to 3 ° Baumé can arise. If the water is heated up to 60 ° Reaumur, the lead chlorine is converted back into lead oxide sulfuric acid in a short time. If one sees that the calcium chloride solution was not stronger than indicated above, and that it is not acidic, then, if the decomposition of gypsum has been conducted correctly, there must be no lead in solution. If this were the case, however, it is knocked out with a little lime. The newly recovered sulfuric acid lead oxide is ready for new decomposition after being washed out several times.

| 285 |
If this sulfuric acid, obtained from hydrochloric acid and gypsum, is to be brought on the market, it must be concentrated and cooked white by one of the known methods up to 66 ° Baumé, or treated according to my invention No. II.

II.
At that time, for the concentration of sulfuric acid, lead pans were first used, in which the acid was evaporated up to 60 ° Baumé; from there it came either in glass or platinum vessels, and then added up to 66 ° Baumé

Fulmen - 29-8-2020 at 02:55

DUDE! That's a great find, if I read it correctly the reaction only consumes gypsum and hydrochloric acid.

macckone - 29-8-2020 at 08:18

Fulmen, don't get to excited until you get it working.
These older journals are not always reliable.
The method seems plausible though.
The question remains if you can get it to work.

Fulmen - 29-8-2020 at 09:16

True, true. It does look kinda messy and crude, and can only produce dilute acid. But it has potential as it uses only cheap OTC ingredients.

I wonder if I should start by making a lead beaker that could handle a bit of acid.

Fulmen - 29-8-2020 at 12:50

So the basic reaction is something like this:

1. PbCl2 + CaSO4 => PbSO4 + CaCl2
2. PbSO4 + 2HCl => PbCl2 + H2SO4

But the devil is usually in the details, and I think we need a better translation.

macckone - 30-8-2020 at 10:09

Fulmen,
Yes and neither one of those reactions looks like it would proceed very well.

S.C. Wack - 30-8-2020 at 11:40

Lunge of course mentions this.

Attachment: gb18542161.pdf (3.1MB)
This file has been downloaded 828 times

Fulmen - 31-8-2020 at 08:32

Macckone: It's the second one that has me puzzled. The first one should go to completion if given enough time, but the second is somewhat counter to common knowledge. However, the text does stress the need for strong HCl, so it might be some special conditions at play.

macckone - 31-8-2020 at 08:50

S. C. Wack,
But does it work?
We all know there are patents and working patents.

The first step PbCl2 + CaSO4 -> PbSO4 + CaCl2
The reactants are practically insoluble.
One of the products is very soluble so in theory could be washed away to drive the reaction but that is going to take a lot of water.
The second step PbSO4 + 2HCl -> PbCl2 + H2SO4 would follow the same procedure, repeated washing with HCl to give a mixed product.

This is very cheap but it was never used in practice.
The only substantial cost is water and boiling it off.
For home use this is probably not prohibitive.
Commercially this is not practical of course.

macckone - 31-8-2020 at 08:55

One thought, because of the nature of the reaction,
this should be possible with a sohxlet extractor.
I don't own one so I can't try it but that would give the driving force.

macckone - 31-8-2020 at 09:10

Fulmen,
If you use a soxhlet, the HCl will boil as an ~20% constant temperature.
This would allow repeated extraction of the PbSO4 with 20% HCl
It will give a mixed product but the H2SO4 would be sequestered as it would not boil. My guess is that the equilibrium is somewhere around 1 to 2% H2SO4.

S.C. Wack - 31-8-2020 at 14:01

Quote: Originally posted by macckone

S. C. Wack,
But does it work?

I don't really care since it seems impractical and I can still get battery acid by the gallon at the hardware store...a better translation was asked for and provided...

The hint was, if you want to know old patented methods, read the old literature google and scihub have, and the highest authority is Lunge, although in retrospect his accuracy is not the best. I may have mentioned in the SO3 thread that Wolters is one of very few of the many etc.-method patents that Lunge sticks his neck out for. The early and late editions are literally voluminous, and this lead process is given short shrift within short shrift. Obviously if the process was ever used, he never heard of it.

macckone - 31-8-2020 at 14:56

S C Wack,
I think the soxhlet method would work.
The europeans are limiting sulfuric acid, hence this thread.

S.C. Wack - 31-8-2020 at 16:52

Like I'm not aware of this? Or like that makes this practical? Not long ago the Europeans had limited sulfuric acid and a love for working with lead. They did use lots of lead, not for this. I would assess the conversions separately and suspect one or both do not go well.

Abromination - 31-8-2020 at 19:22

Just a note on the oxalic acid process, it seems that iron (II) sulfate is not super ideal. The iron oxalate produced is a pain to filter, and some of the particles are fine enough to persist through multiple filtrations. It seems to mostly settle out as the acid is concentrated, but it still doesnt seem particularly great for acid production.
Magnesium sulfate is likely a better option, although magnesium oxalate has a slight solubility in water, especially when hot.

egret - 4-9-2020 at 14:11

Sulfuric acid by reaction of Sulfamic acid with nitric acid (wikipedia Sulfamic acid )

HNO3 + H3NSO3 → H2SO4 + N2O + H2O

The Nitric acid can be generated insitu from calcium nitrate

Ca(NO3)2 + 2 H3NSO3 → H2SO4 + CaSO4 + 2 N2O + 2 H2O

The CaSO4 is filtered out and almost pure solution of sulfuric acid is obtained.
Sulfamic acid is OTC cleaning agent, calcium nitrate is commonly available fertilizer.
Did anybody try?

clearly_not_atara - 4-9-2020 at 19:33

If I were to guess, this is made possible because the ionic reactions are like this:

PbCl2 (Ksp 1.7e-5) + CaSO4 (Ksp 4.9e-5) >> PbSO4 (Ksp 2.1e-8) + Ca2+ (aq) + 2 Cl- (aq)

PbSO4 + HCl + Cl- >> PbCl2 + HSO4- (aq)

That is, the sulfate ions are protonated in the acidic hydrochloric acid solution, but not in the neutral solution. The success of the calcium-to-lead metathesis thus depends on a neutral pH, while the lead sulfate-to-chloride metathesis happens at acidic pH, where the sulfate ion is protonated to bisulfate (pKa ~2).

Tsjerk - 5-9-2020 at 06:35

Adding methanol to a potassium hydrogen sulfate solution works and is nearly quantitative.

http://www.sciencemadness.org/talk/viewthread.php?tid=79548&...

macckone - 13-9-2020 at 10:03

Quote: Originally posted by clearly_not_atara

If I were to guess, this is made possible because the ionic reactions are like this:

PbCl2 (Ksp 1.7e-5) + CaSO4 (Ksp 4.9e-5) >> PbSO4 (Ksp 2.1e-8) + Ca2+ (aq) + 2 Cl- (aq)

PbSO4 + HCl + Cl- >> PbCl2 + HSO4- (aq)

That is, the sulfate ions are protonated in the acidic hydrochloric acid solution, but not in the neutral solution. The success of the calcium-to-lead metathesis thus depends on a neutral pH, while the lead sulfate-to-chloride metathesis happens at acidic pH, where the sulfate ion is protonated to bisulfate (pKa ~2).

This changes with ion concentrations and removal as well.
Both reactions are equilibrium, so if you can remove the calcium chloride, the reaction proceeds reasonably well.
If you remove the sulfuric acid the second reaction proceeds as well.
Hence my thought on the soxhlet extractor where you have continuous if slow removal of reactants from one side of the equation.

AJKOER - 24-9-2020 at 08:02

Came across an interesting electrochemical based reaction to quote:

"2 FeCl2 + 2 H2O + 1/2 O2 --> Fe2O3 + 4 HCl (1)

Hydrochloric acid is regenerated, and commodity oxides of iron are obtained as a result of the reaction [8, 9]."

Source: http://rudmet.net/media/articles/Article_CIS_vol.15_18_pp.28...

Turning to sulfate chemistry, with respect to ferric sulfate by Atomistry at http://iron.atomistry.com/ferric_sulphate.html to quote:

"Upon dilution, ferric sulphate solutions readily undergo hydrolysis, precipitates being obtained which, however, have no well-defined composition.
A study of the electric conductivities of aqueous solutions of the salt indicates that the hydrolysis proceeds in two stages, embodying (1) a rapid change unaccompanied by precipitation, and (2) a slower change, progressing at a measurable rate, and accompanied by the production of a so-called basic salt. Colloidal ferric hydroxide does not appear to be formed during hydrolysis, the salt thus differing from ferric chloride and nitrate...In dilute solution ferric sulphate is reduced by metallic iron to ferrous sulphate."

More recent research on the action of O2 on iron salt solutions, for example, this fully free available 2008 reference: "Air Oxidation of Ferrous Iron in Water" by Ahmet Alıcılar, et al, at http://www.jieas.com/fvolumes/vol081-5/3-5-11.pdf.

“Abstract: Air oxidation of ferrous iron in water was studied...Thereafter, the experiment was successively repeated by blowing air to the solution without and with inert packing. Lastly, the catalytic effect of ferric hydroxide was investigated. While the maximum yield of 86 % is catalytically achieved by blowing air at a neutral medium, the oxidation was almost completed in an alkaline solution even at stationary oxidation was almost completed in an alkaline solution even at stationary atmosphere. The reaction was first order with respect to Fe2+. “

And further:

“Oxidation of iron is achieved by addition of chemical oxidants. However, it can be easily and low cost carried out by contact with air (Wong, 1984). During oxidation of Fe2+ salt aqueous solutions, poor soluble compounds including Fe3+ oxides are formed (Domingo et al., 1994). The composition of precipitate formed depends on numerous parameters such as temperature, pH, concentration, feed rate and anion nature (Das & Anand, 1995; Tolchev et al., 2002)....The oxidation kinetics of Fe(II)(aq) species has been previously reviewed by many workers (Wehrli, 1990; Zhang et al., 1992). The stoichiometry for the overall oxidation of Fe2+ ions by O2 is given by Eq. (1) (Burke & Banwart, 2002).

O2(aq) + 4Fe2+ + 6H2O ↔ 4FeOOH(s) + 8H+ (1) “

So, similar to FeCl2, the slow air pump oxidation of a RT dilute aqueous FeSO4 and metal Iron mix to the point of possible dilute H2SO4 creation may be a subject of my future investigations. Substituting (NH4)2SO4(aq)/Fe for FeSO4(aq)/Fe, may also be interest.

[Edited on 24-9-2020 by AJKOER]

Σldritch - 30-9-2020 at 00:33

I just had a crazy idea of how to make sulfuric acid. As we all know sulfuric acid can be made by adding oxalic acid to Ferrous Sulfate. Ferrous Sulfate being essentially free oxalic acid is what makes this an expensive process. But what if you made the oxalic acid in situ? If you put out a bowl of Ferrous Sulfate, sucrose and a catalytic amount of nitric acid in the dark one might expect the following reactions to take place:

C₁₂H₂₂O₁₁ + 12 HNO₃ = 6 C₂O₄H₂ + 11 H₂O + 12 NO
NO + O₂ + H₂O = HNO₃ (catalyzed by iron in solution)
C₂O₄H₂ + FeSO₄ + 2 H₂O = FeC₂O₄·2H₂O + H₂SO₄

With the overall reaction being:
C₁₂H₂₂O₁₁ + 9 O₂ + 6 FeSO₄*5H₂O = 6 FeC₂O₄·2H₂O + 6 H₂SO₄ + 23 H₂O

The main problem may be overoxidation of the sugar and the slow reaction between sugar and nitric acid. The former is catalyzed by light/iron couple which should be easy to avoid unless iron by itself can do it too. Oxalate may not be rremoved fast enough due to solubilizing effects from the sugar. The latter problem could probably be solved by adding more nitric acid. Nitrogen oxides may be lost but conditions should be very favorable for reoxidation so it may be manegable.

Just like with the oxalic acid process a nice bonus is the Ferrous Oxalate dihydrate byproduct which may be useful to make iron or iron oxide powder though it may be diffucult to contain it 100g+ scale.

Bezaleel - 5-11-2020 at 03:38

Quote: Originally posted by Tsjerk

Adding methanol to a potassium hydrogen sulfate solution works and is nearly quantitative.

http://www.sciencemadness.org/talk/viewthread.php?tid=79548&...

That is of course a great find. But will KHSO4 remain available OTC when H2SO4 won't?

njl - 5-11-2020 at 04:32

Probably, considering it has commercial use as a fertilizer, pool chemical, food additive etc. Those tend to be difficult to regulate.

clearly_not_atara - 20-1-2021 at 16:15

Hypothetically speaking:

Suppose you wanted to use what seems to be the consensus option for conc. H2SO4: you start with 15%-ish H2SO4 and boil it down to around (80%)?, then you feed SO3 through it (via Na2S2O7) until you have roughly 50% oleum, then re-mix with 15% H2SO4 and repeat, etc.

What % do you need to concentrate to in order for it to be safe to pass SO3? You can't pass SO3 into water -- it boils -- but do you need to boil down to 50%? 80%? 90%?

Fyndium - 21-1-2021 at 11:10

We'll look after 5 years, when EU has noticed that any chemicals left from other methods of banning remain at large, and decides to ban them.

The bisulfate method sounds by far the best method to date. The reagents are bearably priced and available, and the process is high yielding and part of the reactants can be reused. For concentrating the acid, efficient methods should be developed, because in general it is PITA.

draculic acid69 - 25-1-2021 at 04:03

I recently seen a video of Epsom salt In a terracotta pot Inside a plastic tub being electrolyzed produced h2so4.cant remember the channel but it was an Australian guy

[Edited on 25-1-2021 by draculic acid69]

Piroz - 25-1-2021 at 15:52

It looks like ferric oxide is a great catalyst for oxidation of sulfur dioxide to sulfur trioxide:

Quote:

Ferric oxide possesses the power of catalytically promoting the combination of sulphur dioxide and oxygen at red heat. The action is perceptible at temperatures just above 400° C., attaining a maximum at 625° C. when 70 per cent, of the sulphur dioxide is converted into trioxide. The origin of the ferric oxide is of considerable importance, that prepared from the hydroxide being particularly active. Admixture of copper oxide increases the efficiency, as does also the presence of arsenic at temperatures above 700° C.

source: http://iron.atomistry.com/ferric_oxide.html
I'm not sure about activity of Fe2O3 from rust (I got large amount of that compound from iron scrap by oxidation under water, then roasting it on glowing charcoal) I'll try to use it as catalyst soon.
There is another reaction that I'm focusing on:

Quote:

Sulphur dioxide may be passed into a solution of a ferric salt for a similar purpose, or it may be generated in the solution by addition of an alkali sulphite and a little dilute mineral acid. Thus, ferric sulphate is reduced in accordance with the equation

Fe2(SO4)3 + SO2 + 2H2O = 2FeSO4 + 2H2SO4.

Source: http://iron.atomistry.com/iron_salts.html

The last reaction may be used for recovering sulfur dioxide that's no oxidised on the catalyst.

[Edited on 25-1-2021 by Piroz]

pip - 31-1-2021 at 17:11

Call me both a risk taker and lazy, but I’d be offering to pickup used car batteries and distilling before going through that much work for that little yield.

A quick and relatively cheap way to make sulfuric acid in EU

P-α3 - 25-9-2021 at 13:20

The Best way is to mix NaHSO4 with HCl in stochiometric amounts, maybe excess HCl if you want to be on the safer side. Then boil it, let it cool down as much as possible to crystalize the NaCl. Filter out the NaCl. Then Distill the acid. First to come out is the HCl then the Sulfuric acid. Both of the chemicals are cheap to buy in the EU atleast and are a great substitute against other ways of obtaining sulfuric acid. I calculated that with my efficiencies of making I can get almost as low in price to commercially available sulfuric acid in the US. So I think its pretty good apart from the fact that you will always have the paranoia that the round bottom or the flat bottom will crack when you try distilling the acid.

Tsjerk - 25-9-2021 at 13:57

Or you crash out Na2SO4 with methanol/ethanol from a solution of NaHSO4.

Edit: KHSO4 works better.

[Edited on 25-9-2021 by Tsjerk]

woelen - 26-9-2021 at 09:27

But how do these methods help in getting H2SO4? In the EU we still can buy H2SO4 (15%).

Is boiling down 15% H2SO4 more difficult than making the acid from HCl and NaHSO4 or making it from ethanol and KHSO4? Boiling down is a pain in the ass, but I'm afraid that making pure _concentrated_ acid from HCl and NaHSO4 is even more a pain in the ass.

If you really want concentrated acid (e.g. 96%), from chemicals which can be obtained legally in the EU, and without boiling down so much that you reach 96%, then I see only one method:
- boil down the 15% acid until it reaches 85% or so. At that concentration, further boiling will produce intense fumes and you will lose a lot of acid if you want to get it at 95%. This is not practical.
- add chlorosulfonic acid (slowly and carefully, while stirring) to the 85+ % acid. This will increase the concentration to 95% with production of a lot of HCl (which can be absorbed in e.g. methanol to create an interesting side-product).
This method is not for the faint of heart though. Chlorosulfonic acid is EXTREMELY corrosive and the reaction with 85% acid is very exothermic. It must be done very slowly with good stirring. I, however, think that this reaction is less dangerous than distilling or boiling appr. 95% acid.

The above method is nice for making a small amount of concentrated H2SO4, e.g. 50 ml. I would not use it for makin g liters of conc. H2SO4. But making liters of conc. H2SO4 is not wise anyway. Just make what you need for experiments you do now and maybe a few experiments to come. For many more experiments boiled down acid to e.g. 40% is suitable and that is not an issue at all, giving only water vapor and no need to heat to insanely high temperatures like 300 C.

[Edited on 26-9-21 by woelen]

teodor - 6-10-2021 at 14:08

100 ml of 85% H2SO4 contains 26.7g H2O. Let say, we try to get 100% H2SO4 (to keep the math a bit simpler). The reaction of this water with chlorosulfonic acid will require 173g of the acid. It will give 145g H2SO4, so the result will be 296g H2SO4 total. So, we will get 160 ml of 100% H2SO4. So, ~170 g HClSO4 -> 160 ml 100% H2SO4. And for 95%, roughly 85g HClSO4 -> 120 ml (and + 100ml 85% = a lot of 15 % H2SO4). This is not the most economical method.
On the other hand, if you can buy somewhere SiCl4 I would give a try to another method.

[Edited on 6-10-2021 by teodor]

teodor - 10-10-2021 at 09:24

I tried to prepare ~100% H2SO4 by mixing 2.4 ml HClSO3 and 6.5 ml 96% H2SO4. Eventually, it doesn't result in HCl evolution except for several small bubbles. Probably the preparation requires heating. I used the mix to prepare 6M H2SO4 solution in glacial acetic acid. When I've added the acetic acid the solution became warm & highly saturated with HCl and was boiling probably because of HCl evolution. After some time the pressure stopped building. Now it is a fuming liquid with the mixed smell of AcOH and HCl.
So, I think the mixing of chlorosulfonic acid and sulfuric acid doesn't cause a fast reaction and easy HCl elimination if the concentration of H2SO4 is already high.

woelen - 10-10-2021 at 09:46

That's a quite surprising result. If you look at the reaction between water and chlorosulfonic acid, then I would expect that mixing chlorosulfonic acid with something, which contains a little water quickly destroys that little water. Getting rid of dissolved HCl is not an issue, simply heat and you drive it off, but if the destruction of water is not fast and complete with chlorosulfonic acid, then that indeed is a problem.

I have done the prep of 100% H2SO4 from 20% oleum and 96% H2SO4 and that works really nice (although mixing of these still produces noticeable heat) and results in a slightly fuming acid. So, I expected that the reaction with chlorosulfonic acid also would be fast and complete, but apparently the HCl, bound to the SO3, spoils the reaction to quite some extent.

teodor - 10-10-2021 at 23:13

I think the reason is that water is not quite "free" in 96-98% H2SO4. For example, according to https://doi.org/10.1021/ja01630a063 in H2SO4 with a concentration of 90%+ the water exists only in the fully ionized form (H3O+). And probably H3O+ and HClSO3 can coexist. (I will elaborate more on this in my thread about the reaction of chlorosulfonic acid with metals. I have mixtures of 96% H2SO4 and HClSO3 in dichloromethane which I keep in closed bottles for several days - there is no pressure).

woelen - 11-10-2021 at 01:06

If this hypothesis is true, then making pure H2SO4 from HSO3Cl and 90% H2SO4 could be done, if also some SO3 is added:

In 90% H2SO4 nearly all water is present as H3O(+):

H2SO4 + H2O <--->>> HSO4(-) + H3O(+)

SO3 adds to sulfate (and bisulfate): SO3 + HSO4(-) ---> HS2O7(-).
This compound strongly attracts hydrogen ion: HS2O7(-) + H3O(+) --> H2S2O7 + H2O

Then the free H2O can react with HSO3Cl to form H2SO4 + HCl.

This may be interesting from a theoretical point of view, I do not see this as a practical method for preparing H2SO4. Getting SO3 (or strong oleum) is much harder than getting HSO3Cl. At least, oleum is MUCH more expensive, if you can get it at all. I once read the term 'poor man's oleum' for chlorosulfonic acid, as it can be used as a substitute in some special cases, but it definitely is not a snap-in replacement for true oleum. Some people over here succesfully made small quantities of oleum, but making and isolating it safely is HARD. I do have a small quantity of oleum, but I only use that for really special experiments (e.g. making chloryl compounds), because of its very high price and the good chance that I'll never get an opportunity to get my hands on it again.

teodor - 12-10-2021 at 02:55

Last time I bought 30% oleum (you probably can guess where) at the same price per kg as chlorosulfonic acid. But the point is if you can buy oleum you probably can buy 98% H2SO4 also. So, this method is definitely not "scalable".

As per equilibrium, I can cite the book "Sulfur in organic and inorganic chemistry, Volume 2" (A.Senning):
"Sulfuric acid is slightly self-dissociated into sulfur trioxide and water
(1) H2SO4 <-> H2O + SO3
... Water is nearly completely ionized as a base
(2) H2O + H2SO4 <-> H3O(+) + HSO4(-)
Sulfur trioxide is completely converted to disulfuric acid, H2S2O7. This acid is ionized as a moderately weak acid
(3) H2S2O7 + H2SO4 <-> H3SO4(+) + HS2O7(-)
Thus, since the ions H3SO4(+) and HSO4(-) are in equilibrium as a consequence of the autoprotolysis reaction, it follows that the ions H3O+ and HS2O7- must also be in equilibrium
(4) 2H2SO4 <-> H3O(+) + HS2O7(-)
... The complete self-dissociation reaction in the sulfuric acid solvent system can be described then by the above equations". (Then the 4 equilibrium constants are given for 25C).
This matter is a bit hard for me, but probably I would understand that with a help of some visual or mathematical model. Because all 4 equations are in a single system, I am unable to model the situation you have described without the understanding of this 4 equations system first.

As a practical method for a basic laboratory (no distillation), I probably will do some experiments following Tsjerk suggestion of using KHSO4/NaHSO4 and alcohol.

Because getting and isolating SO3 by all known methods is hard in an average home lab I have a plan to continue my experiments about "oxidation of SO2 in pyrosulfuryl chloride" but of course it is only the field of thoughts and experiments yet.

Another field of possible experiments could be separating the mixture of H2SO4 and H3PO4 with a help of some organic solvent. If it is ever possible, it will allow concentrate H2SO4 with P2O5 which is cheaper and more available than chlorosulfonic acid or oleum. But I doubt based on the same considerations about water ionization.

As a practical approach, I see another way. We probably can use 85% H2SO4 + P2O5 in many reactions which require 95%+ H2SO4. Because we know how to use 36% HCl and can avoid using 100% HCl in most cases by some tricks. Now probably is the time to invent more tricks for using 85% H2SO4 instead of concentrated acid.

zerodan - 27-10-2021 at 13:11

Quote: Originally posted by woelen

If this hypothesis is true, then making pure H2SO4 from HSO3Cl and 90% H2SO4 could be done, if also some SO3 is added:

In 90% H2SO4 nearly all water is present as H3O(+):

H2SO4 + H2O <--->>> HSO4(-) + H3O(+)

SO3 adds to sulfate (and bisulfate): SO3 + HSO4(-) ---> HS2O7(-).
This compound strongly attracts hydrogen ion: HS2O7(-) + H3O(+) --> H2S2O7 + H2O

Then the free H2O can react with HSO3Cl to form H2SO4 + HCl.

This may be interesting from a theoretical point of view, I do not see this as a practical method for preparing H2SO4. Getting SO3 (or strong oleum) is much harder than getting HSO3Cl. At least, oleum is MUCH more expensive, if you can get it at all. I once read the term 'poor man's oleum' for chlorosulfonic acid, as it can be used as a substitute in some special cases, but it definitely is not a snap-in replacement for true oleum. Some people over here succesfully made small quantities of oleum, but making and isolating it safely is HARD. I do have a small quantity of oleum, but I only use that for really special experiments (e.g. making chloryl compounds), because of its very high price and the good chance that I'll never get an opportunity to get my hands on it again.

I'm not sure how's that related to the original goal of this thread.
Unless bottles of chlorosulfuric acid are sold at grocery stores and I missed a memo, what's the point if the starting reagent is even harder to get than the final product.
I wouldn't even know where to source that apart from sigma or shady ebay sellers (which would probably sell you h2so4 regardless of the ban).

Linus1208 - 27-10-2021 at 15:34

What about thermal decomposition of NaHSO4?

NaHSO4 decomposes at 280°C to form H2O and at 480°C to form SO3, so the temperatures are achievable. If you bubble the gases in a small amount of water (there's already the stoichimetric amount present in the reaction), you should be able to get quite a high concentration of sulfuric acid.
You also don't need to seperate byproducts.
I don't know how efficient this process would be, but low effiency should be that of a problem, as technical NaHSO4 is available in kg quantities in the hardware store as pH- granules for swimming pools (at least in Germany, I guess it is similar in other (EU) countries).

Jome - 27-10-2021 at 23:45

How about SO2Cl2 to get rid of those last pesky few % of water?

woelen - 28-10-2021 at 00:18

SO2Cl2 reacts with water very slowly. Besides that, it does not dissolve in H2SO4, nor in water, so you only have a reaction at the contact leyer between the two solutions. SO2Cl2 also is not OTC, but making it seems doable for the more advanced amateur.

@zerodan: In my previous post I already wrote myself that the described method is interesting from an academic point of view, but should not be considered a practical method for making H2SO4 in larger quantities.

Where I live, however, chlorosulfonic acid, can be obtained legally, while getting H2SO4 at more than 15% is illegal. The world is a strange place indeed

but reality for me is that chlorosulfonic acid can be ontained more easily than conc. H2SO4. A few weeks ago I actually purchased some HSO3Cl from a respected supplier (not some shady eBay seller) who sells to private individuals.

Linus1208 - 28-10-2021 at 02:10

Oh, and I thought being able to buy kilos of red phosphorus and a few 100g of KMnO4 but heavily restricted KNO3 in Germany was weird

Keras - 28-10-2021 at 03:15

Also, I just ordered p-toluenesulphonic acid. It is still legal to buy it in any quantity. It is quite handy when you need sulphuric acid for catalysis. I wonder if one can use it to make ether out of ethanol the classic way. I will try that when I put my hands on it.

teodor - 29-10-2021 at 01:19

I also have an idea to compare p-toluenesulphonic acid vs H2SO4/HCl for ester synthesis. I plan 3 experiments of making sec-butyl acetate using these 3 acids.
As for di-ethyl ether, I suppose KHSO4 should also work.

SWIM - 29-10-2021 at 11:36

Ozone does have some solubility in concentrated sulfuric acid; so maybe you could oxidize sulfur to SO3 in sulfuric acid with ozone gas.

If viable, it could be useful for bringing the concentration up (boil off water to 80% or so, then treat with sulfur/ozone), or even for making fuming acid.

I've read a few things that make it look like sulfur goes straight to SO3 in ozone reactions and not through an intermediate SO2 stage, but I'm not at all sure about this.

edit: the oxygen would need to be awfully dry.

[Edited on 29-10-2021 by SWIM]

teodor - 30-10-2021 at 04:41

I believe O3 can also contaminate the product with persulfuric acid.

Could PbO2 be used for sulfur oxidation somehow?

SWIM - 30-10-2021 at 09:24

I hadn't heard that Ozone forms persulfuric acid.
I thought you needed hydrogen peroxide.

I don't know much about lead dioxide oxidations, sorry.

AJKOER - 5-12-2021 at 12:50

Concept pending my acquisition of TiO2, namely, a photolysis experiment involving a suspension of TiO2 acting on hot aqueous ammonium sulfate.

Starting reactions:

(NH4)2SO4 = 2 NH4+ + SO4(2-)

NH4+ = H+ + NH3

With TiO2 as a photocatalyst, expect UV light generation of e- and h+ (an electron hole capable of converting OH- to .OH).

Possible reactions forming associated products with various yields:

e- + H+ = •H

e- + •NH2 --> NH2-

e- + •OH --> OH-

e- + •SO4- --> SO4(2-)

•H + •H --> H2

•H + NH3 --> •NH2 + H2

•H + SO4(2-) --> •HSO4- --> H+ + •SO4-

•H + •SO4– --> HSO4-

•H + •NH2 --> NH3

h+ + OH- (from water) --> •OH

h+ + SO4(2-) --> •SO4-

h+ + NH2- --> •NH2

•OH + NH3 --> H2O + •NH2

•OH + SO4(2-) --> OH- + •SO4–

•OH + •SO4– --> HSO4- + 1/2 O2 (or possibly HOSO4-)

•OH + •NH2 --> NH2OH

•NH2 + •NH2 --> N2H4

•SO4– + •SO4– --> S2O8(2-) (See https://www.sciencedirect.com/topics/chemistry/sulfate-radic... )

•SO4– + OH- --> SO4(2-) + •OH (See https://www.sciencedirect.com/topics/chemistry/sulfate-radic... )

•SO4– + H2O --> SO4(2-) + •OH + H+ (See https://www.sciencedirect.com/topics/chemistry/sulfate-radic... )

S2O8(2-) + hv --> •SO4– + •SO4– (See https://www.sciencedirect.com/topics/chemistry/sulfate-radic... )

HSO5- + hv --> •OH + •SO4– (See https://www.sciencedirect.com/topics/chemistry/sulfate-radic... )

So, a heated concentrated solution of ammonium sulfate undergoing TiO2 UV photolysis may form several transient radical species while liberating from solution NH3, H2 and even some toxic N2H4 (so best performed with ventilation) leaving behind H2O and H2SO4 (and more like H2S2O8 in small amounts).

Note, per a sciencedirect reference, TiO2 does not dissolve in dilute warm H2SO4 (see https://www.researchgate.net/post/In-which-solvent-can-TiO2-... ) so the photocatalyst should continue to function here (pending my experimental verification). I am also considering the substituting a photo catalytic dye for TiO2 with the understanding of a less pure product.

Also, ammonium sulfate can be sourced from the action of aqueous or gaseous ammonia on aqueous Epsom Salt (a highly pure MgSO4 hydrate as people enjoy a good healthy mineral bath).

A Little Off Topic: Upon presentation of the cited reactions above, I noticed that h+ + •SO4– reaction is missing! That is, what is/can be
the action of an electron hole on the sulfate radical anion?

Per this related article "Sulfate Radical Anions (SO4•-) as Donor of Atomic Oxygen in Anionic Transannular, Self-Terminating, Oxidative Radical Cyclizations" at https://pubs.acs.org/doi/abs/10.1021/ol006527y suggests a speculative answer to be:

h+ + •SO4– --?--> SO3 + O

where the above speculated reaction (under appropriate conditions) very interestingly involves both the formation of both SO3 and atomic oxygen. And, since the addition of water to SO3 is a path to H2SO4, this discussion is not particularly off topic.

[Edited on 5-12-2021 by AJKOER]

AJKOER - 6-12-2021 at 07:12

I actually just seemingly found a supporting verification to my suggested use of a dye here on the ammonium sulfate path to an acid (see https://www.dharmatrading.com/home/did-you-know-how-to-take-...). To quote:

"Ammonium Sulfate is a leveling agent, which means it slows the absorption of the dye into the fiber. It also causes the dye bath to become acidic very gradually, so the dye fixes over time rather than all at once."

where the slow nature of the acidification could be due to normal light exposure (this is not a photolysis experiment per se, but an exercise in the application of a dye) and the presence of fabric may scavenge radicals that could have been attacking the (NH4)2SO4.

[Edited on 7-12-2021 by AJKOER]

teodor - 7-12-2021 at 06:20

AJKOER, I was wondering where you are. It's nice to see you again.

According to SM wiki:

"Sulfamic acid melts at 205 °C before decomposing at higher temperatures to water, sulfur trioxide, sulfur dioxide, and nitrogen.

H3NSO3 → H2O + SO3 + SO2 + N2".

In EU probably everybody can buy 1 kg H3NSO3 per 6 EUR and for bigger quantities, it is like 1.6 EUR/kg. The thermal decomposition is probably at a temperature close to H2SO4 boiling point but Wikipedia says it starts at 205C. If so it is a bit easier than concentrating H2SO4 by distillation.

But there are also other possible decomposition reactions, also catalysts ...

AJKOER - 27-12-2021 at 10:37

Thanks Teodor.

I guess you detected the warmth in my threads.

Now, renting in a more tropical setting likely inspired by a neighbor braking her ankle on ice on her way to work!

As such, more limited on reagents but finding friendly galvano-assisted reactions a path to new things.

Also, some Patents on the horizon in 3 diverse areas (computational theory relating to small sample settings, new paragon in casino gaming and, of also, a new/safer/green bleaching).

Tsjerk - 8-1-2022 at 09:04

Maybe it is not all that practical to get sulfuric acid up to 80%, but from that point on: Can sulfuric acid be concentrated with a Dean Stark trap with hexane?

No matter how slow, the only thing used is energy to boil the hexane, which boils quite low.

macckone - 8-1-2022 at 09:10

Quote: Originally posted by Tsjerk

Maybe it is not all that practical to get sulfuric acid up to 80%, but from that point on: Can sulfuric acid be concentrated with a Dean Stark trap with hexane?

No matter how slow, the only thing used is energy to boil the hexane, which boils quite low.

you can concentrate sulfuric acid with nothing more than a beaker, a watch glass, a hot plate and a tray filled with sand and baking soda to catch splatter. I find fiberglass mesh works better. yes it is slow, yes it is tedious but it works, yes you get sulfuric acid escaping, but it works otherwise you couldn't concentrate sulfuric acid by distillation.

clearly_not_atara - 23-1-2022 at 11:33

Quote: Originally posted by Tsjerk

Maybe it is not all that practical to get sulfuric acid up to 80%, but from that point on: Can sulfuric acid be concentrated with a Dean Stark trap with hexane?

What I do know is that you can use dilute (50%, which is easy) sulfuric acid to make PTSA with excess toluene and a Dean-Stark. That leads to the following sequence:

- disproportionate KHSO4 or maybe NaHSO4 with ethanol

- heat H2SO4/EtOH/H2O mixture to ~180 C to drive off residual ethanol as ethylene/ether (needs good ventilation)

- Dean-Stark dilute H2SO4 with toluene to obtain paratoluenesulfonic acid

So it might actually be easier to make solid tosylic acid than concentrated sulfuric acid. I also know that PTSA will precipitate from sufficiently acidic aqueous solutions but I'm not sure about the practical use of this.

Crazy_Chemist - 9-3-2022 at 22:21

I had time to buy a liter of battery acid at 37.5% before it was banned. I have not used it yet, but I have seen that it is easy to concentrate it by boiling off the water at about 100 ° C.

Colleen Ortiz - 17-3-2022 at 04:33

Hello,
The contact method produces sulfuric acid from sulfur, oxygen, and water. Sulfur is burnt in the first phase to create sulfur dioxide.
S (s) + O2 (g) → SO2 (g)
In the presence of a vanadium(V) oxide catalyst, this is then oxidized to sulfur trioxide.
2 SO2 + O2 (g) → 2 SO3 (g) (in presence of V2O5)
Finally, the sulfur trioxide is processed with water to generate 98-99 percent sulfuric acid (typically as 97-98 percent H2SO4 with 2-3 percent water).
SO3 (g) + H2O ( l) → H2SO4 (l)
Because of the extremely exothermic nature of the reaction, directly dissolving SO3 in water is impracticable. Instead of a liquid, mists develop. Alternatively, the SO3 can be absorbed into H2SO4 to make oleum (H2S2O7), which can then be diluted to produce sulfuric acid.
H2SO4( l) + SO3 → H2S2O7(l)
When oleum reacts with water, it produces concentrated H2SO4.
H2S2O7(l) + H2O(l) → 2 H2SO4(l)

Texium - 17-3-2022 at 10:05

Quote: Originally posted by Colleen Ortiz

Hello,
The contact method produces sulfuric acid from sulfur, oxygen, and water. Sulfur is burnt in the first phase to create sulfur dioxide.
S (s) + O2 (g) → SO2 (g)
In the presence of a vanadium(V) oxide catalyst, this is then oxidized to sulfur trioxide.
2 SO2 + O2 (g) → 2 SO3 (g) (in presence of V2O5)
Finally, the sulfur trioxide is processed with water to generate 98-99 percent sulfuric acid (typically as 97-98 percent H2SO4 with 2-3 percent water).
SO3 (g) + H2O ( l) → H2SO4 (l)
Because of the extremely exothermic nature of the reaction, directly dissolving SO3 in water is impracticable. Instead of a liquid, mists develop. Alternatively, the SO3 can be absorbed into H2SO4 to make oleum (H2S2O7), which can then be diluted to produce sulfuric acid.
H2SO4( l) + SO3 → H2S2O7(l)
When oleum reacts with water, it produces concentrated H2SO4.
H2S2O7(l) + H2O(l) → 2 H2SO4(l)

Thank you for reciting a textbook once again... this thread is supposed to be a discussion of practical methods though. Everyone knows how the contact process works in theory. Showing a working contact process system built by an amateur would be another story. If you really aren't a bot, could you please explain why all of your posts sound like one?

Jinc8 - 16-4-2022 at 04:49

Would it be possible to create H2SO4 with Citric acid/citrate salts and Sulfate salts (Ferrous in my case)? I have a bunch of citric acid and I'd like to use that instead of the usual Oxalic acid, if possible.

I'm not sure how citric acid itself would react with FeSO4, as the former seems to react differently depending on the ph of solution (for example with NaHCO3 which produces monosodium Citrate, while NaOH produces Trisodium Citrate)

Or I could just buy 1kg of Oxalic acid for like 6€ and make it that way.

[Edited on 16-4-2022 by Jinc8]

clearly_not_atara - 17-4-2022 at 08:18

Oxalic acid is 80 times stronger than citric acid, so the product concentration will be 80 times better with oxalic acid assuming the same precipitation characteristics. But citrate salts also don't precipitate as easily as oxalate salts. In short, no, you can't.

Keras - 17-4-2022 at 23:00

Quote: Originally posted by clearly_not_atara

So it might actually be easier to make solid tosylic acid than concentrated sulfuric acid. I also know that PTSA will precipitate from sufficiently acidic aqueous solutions but I'm not sure about the practical use of this.

I’m not sure it’s worth the effort, given that PTSA is freely available.

Jinc8 - 18-4-2022 at 01:32

Quote: Originally posted by clearly_not_atara

Oxalic acid is 80 times stronger than citric acid, so the product concentration will be 80 times better with oxalic acid assuming the same precipitation characteristics. But citrate salts also don't precipitate as easily as oxalate salts. In short, no, you can't.

Yeah, I tried doing it yesterday in a small test tube and nothing really happened

Thanks for the explanation

H2SO4 from CaSO4 (gypsum), PbCl2 (or Lead) and HCl

RU_KLO - 18-3-2023 at 07:26

Today I tested the possibility to make H2SO4 from PbCl2 + CaSO4.

you will find information in some previous post (starting here: https://www.sciencemadness.org/whisper/viewthread.php?tid=15...)

1. PbCl2 + CaSO4 => PbSO4 + CaCl2
2. PbSO4 + 2HCl => PbCl2 + H2SO4

Main problems:

1) working with lead salts
2) low solubility of reagents
a) being CaSO4 allmost insoluble
b) PbCl2 In water: 0.673 g/100 mL water at 0 °C; 0.99 g/100 mL water at 20 °C; 3.34 g/100 mL water at 100 °C, solubility product Ksp = 1.7×10−5 at 20 °C.)

from literature found:

PbCl2 is more soluble in hot water.
CaSO4 decrease its solubility with increasing temperature (being sweet spot at aprox 40°) in (https://www.researchgate.net/figure/Gypsum-solubility-in-H-2...)
CaSO4 decrease its solubility as more CaCl2 gets into solution (https://pubs.acs.org/doi/10.1021/je050217e)

For doing this experiment a temperature of 70° (+/-10°) was chosen. (patent states 54°-84° - https://www.sciencemadness.org/whisper/viewthread.php?tid=15... )

PbCl2 was made previously (months ago from battery lead + HCl) it is not anhydrous, and was in a "clumpled" form. Did not powderize, but dissolution in H2O was helped with a stirrod smashing the clumps.

Aproximately 1.88gr PbCl2 was dissolved in 100ml H2O at 70° (+/-10°). It fully dissolved.
Aprox. 0.47gr CaSO4 was dissolved in 100ml H2O at 40° in another beaker. It did not fully dissolved (as when standing for 5 minutes, some CaSO4 precipitate was left.
Whats left made the second CaSO4 solution.

CaSO4 sol. was poured through a filter (cotton plug) in the PbCl2 solution while hot with stirring. This was done to avoid confusion of newly precipitate (PbSO4) with non dissolved CaSO4.

A white precipitate was formed. (PbSO4), it was decanted and the clear liquor was transfered to another beaker.

The liquor was tested for Pb+ ions in a test Tube with HCl (20%). A white precipitate indicates that PbCl2 was not fully consumed.
So a second solution of CaSO4 was made with 50ml H2O and the not dissolved CaSO4 from the first CaSO4 sol. Heated to aprox 40°. It also did not fully disolved.
It was added via the same filter.
A little more PbSO4 was formed.

Both PbSO4 where put together and 100ml H2O was added, decanted, water removed. this process was done twice to wash the PbSO4.aprox 10ml of water was left with the white precipitate. pH was slighly acidic (~ 6 - yellow in a universal pH paper). It was not weighted.

Then 100 ml 20% HCl was added and heated to 90°. (at aprox 50° the the solution became clear)
After 5 minutes a sample was taken - aprox 10 ml - and tested with a saturated CaCl2 solution. (No precipitate was found - or there was no H2SO4 or it was highly diluted)

As I do not own a fume hood, waited till next morning for concentrating outside.
The next morning a prepipitate was found (PbCl2) in the beaker. It was very "diamond" shine. Solution was poured in a new beaker, trying to avoid PbCl2 to come over.

Solution volume was reduced from 100ml to aprox 10ml. More PbCl2 crashed out. The solution was clear yellow. It was transfered to a new test tube (trying to avoid PbCl2 to come over- which was very small).

10ml saturated CaCl2 was added and a white precipitate was formed, covering (after 10 minutes standing) 1/3 of a test tube (20ml solution).
If Im not wrong, this confirmates H2SO4 was produced.

The recovered PbCl2 was meassured, giving 0.79gr. (but take this with a Ton of salt, because it was weighted wet)

So the procedure works, being the main problem CaSO4 solubility.

I will try this procedure again, but instead of making a CaSO4 solution, I will pour it directly into the hot PbCl2 solution. The knack of this procedure is to get the most PbSO4 salt.

Rainwater - 18-3-2023 at 09:44

Quote: Originally posted by RU_KLO

For doing this experiment a temperature of 70° (+/-10°) was chosen. (patent states 54°-84° - https://www.sciencemadness.org/whisper/viewthread.php?tid=15... )

Code:


PbSO4 + 2HCl = PbCl2 + H2SO4
ΔH -68.8686  kJ/mol 
ΔS -229.2414  J mol/K 
temp 298.15  Kelvin  25  Celsius  
ΔG = ΔH - TΔS 
ΔG -0.52  kJ/mol 
ΔG=0   T= 300.42 K 27.27 C

Thermodynamics indicates higher temperature will cause the reverse reaction to occur.
The reaction might progress better under less heat.
This will be up to kinetics.

But getting PbSO4 to disolive is difficult.
One thing that might be worth trying.

The solubility of PbSO4 is greater in sulfuric acid than in water.
https://nvlpubs.nist.gov/nistpubs/jres/22/jresv22n1p55_A1b.p...
Other than using a "policeman" to wipe the acid from the beaker. The paper suggest a 9~11% H2SO4 solution will provide the highest solubility. But your local law enforcement might assist you

[Edited on 18-3-2023 by Rainwater]

RU_KLO - 18-3-2023 at 13:36

the 70° was for first part (PbCl2 + CaSO4). This was a choice between better solubility of PbCl2 and not to hot (away from 40°) - best solubility temperature for CaSO4)

the second part (PbSO4 + HCl) started at room temperature then was heated till 90° and noted that at 50° aprox the solution cleared.

Probably as you say when arround 30° H2SO4 was produced which increased the solubility of PbSO4.

But in the end for concentrating, 108°-110° was achieved.

Is there another way to check if H2SO4 is produced in a solution containg HCl? my idea was to use CaCl2 which will precipitate Ca2SO4.

Rainwater - 18-3-2023 at 14:17

Quote: Originally posted by RU_KLO

Is there another way to check if H2SO4 is produced in a solution containg HCl?

Bring the solution to 120c. This will remove all HCl and some water
http://www.chm.bris.ac.uk/motm/h2so4/h2so4h.htm

Quote:

Firstly, add dilute HNO3. This is to prevent precipitation of other insoluble barium compounds such as BaCO3 or BaSO3. Secondly, add Ba(NO3)2 (aq). If sulfuric acid or a sulfate is present a white precipitate will be immediately observed.
Ba2+ (aq) + SO42- (aq) BaSO4 (s)

blogfast25 - 18-3-2023 at 14:17

Quote: Originally posted by RU_KLO

Is there another way to check if H2SO4 is produced in a solution containg HCl? my idea was to use CaCl2 which will precipitate Ca2SO4.

Ba(NO₃)₂: because BaSO₄ is extremely insoluble. It's a well-known test for sulfates.

[Edited on 18-3-2023 by blogfast25]

RU_KLO - 19-3-2023 at 05:02

Ok, Thanks, but no Barium salt.

If you have to buy only one barium salt, that allows making others (like Ba(NO3)2)
Barium in my country is not OTC (the only Barium salt that can be purchased in a pharmacy is Barium Sulphate (for contrast agent in XRay).
Or maybe Baryte mineral.

Which one should I try to buy? or just try to buy Ba(NO3)2.

Rainwater - 19-3-2023 at 05:44

A less sensitive test
2NaCl + H2SO4 → 2HCl + Na2SO4
HCl + NH3 = NH4Cl

Again br8ng the solution to 120c to remove HCl.
Then close by place a container of ammonia, add salt(NaCl) to the H2SO4 solution and look for white smoke(NH4Cl).

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