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Author: Subject: Small scale production of H2SO4 in the amateur lab
teodor
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[*] posted on 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.
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SWIM
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[*] posted on 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]




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teodor
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[*] posted on 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?
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[*] posted on 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.











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AJKOER
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[*] posted on 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]
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AJKOER
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[*] posted on 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]
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teodor
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[*] posted on 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 ...
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[*] posted on 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).
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[*] posted on 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.
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[*] posted on 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.
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clearly_not_atara
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[*] posted on 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.




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 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.



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[*] posted on 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)
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[*] posted on 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?



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[*] posted on 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]
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clearly_not_atara
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[*] posted on 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.



[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 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.
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[*] posted on 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
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