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teodor
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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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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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I believe O3 can also contaminate the product with persulfuric acid.
Could PbO2 be used for sulfur oxidation somehow?
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SWIM
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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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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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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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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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AJKOER
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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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Tsjerk
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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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macckone
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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.
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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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| 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?
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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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Crazy_Chemist
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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.
Amateur chemist, just for fun!
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Colleen Ortiz
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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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Texium
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| 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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Jinc8
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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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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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Keras
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| 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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Jinc8
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| 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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RU_KLO
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H2SO4 from CaSO4 (gypsum), PbCl2 (or Lead) and HCl
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.
Go SAFE, because stupidity and bad Luck exist.
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Rainwater
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| 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
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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]
"You can't do that" - challenge accepted
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RU_KLO
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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.
Go SAFE, because stupidity and bad Luck exist.
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Rainwater
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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)
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"You can't do that" - challenge accepted
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blogfast25
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| 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.
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Ba(NO3)2: because BaSO4 is extremely insoluble. It's a well-known test for sulfates.
[Edited on 18-3-2023 by blogfast25]
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RU_KLO
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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.
Go SAFE, because stupidity and bad Luck exist.
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Rainwater
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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).
"You can't do that" - challenge accepted
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