| Pages:
1
..
3
4
5 |
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
The latter. Or BaCl2.
Or: react Desert Rose with Al powder:
BaSO4 + 8/3 Al ==> BaS + 4/3 Al2O3
Then leach out the Ba2+ with strong HCl.
[Edited on 19-3-2023 by blogfast25]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by Rainwater  | 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). |
That's not SA specific. Assuming the test works it will work for almost any acid. But it's not a bad idea...
[Edited on 19-3-2023 by blogfast25]
|
|
|
Twospoons
International Hazard
Posts: 1349
Registered: 26-7-2004
Location: Middle Earth
Member Is Offline
Mood: A trace of hope...
|
|
I stumbled across this patent https://patents.google.com/patent/US20040234441A1/en describing the conversion of ammonium sulfate to ammonia and sulfuric acid. I know patents
can be a somewhat dubious source of information but if it works it would be a reasonably cheap and simple method. The temperatures are reasonable and
it doesn't involve SO3. It does probably require accurate temperature control though.
| Quote: |
The present invention concerns a method for manufacturing of ammonia gas (NH3) and sulphuric acid (H2SO4) characterized wherein ammonium sulphate
((NH4)2SO4) is mixed with concentrated sulphuric acid, and by heating the mixture to a temperature above 235° C., whereby ammonium sulphate is
melted, and thereafter by heating the mixture to above 280° C. but below the boiling point of concentrated sulphuric acid (290° C.-320′), by which
ammonium sulphate is decomposed into ammonia gas and sulphuric acid liquid simultanously, and where the produced sulphuric acid during the reaction
time is mixed ideally with the concentrated sulphuric acid from the initial basis mixture.
[0011]
By this method it is achieved to decompose ammonium sulphate to a fully extent into ammonia gas and sulphuric acid liquid, where the ammonia gas
during processing is evaporated and separated out from a rest solution consisting of sulphuric acid.
|
Helicopter: "helico" -> spiral, "pter" -> with wings
|
|
|
Texium
|
Thread Moved
5-12-2023 at 14:18 |
clearly_not_atara
International Hazard
Posts: 2819
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
^In brief, I think this patent is wrong, but based on what I've read before about ammonium sulfate decomposition, similar reactions do exist. The
thermal decomposition sequence is, roughly:
ammonium sulfate + heat >> ammonium bisulfate + ammonia (g)
ammonium bisulfate + more heat >> ammonium pyrosulfate + water (g)
ammonium pyrosulfate + lots of heat >> SO2 + H2O + N2 (undesirable)
So if you add ammonium sulfate to sulfuric acid, you will primarily get ammonium bisulfate, which undergoes the thermal decompositions described
above. But you can react ammonium bisulfate with sodium or potassium salts to get the mixed sulfate, which then releases ammonia and water to give
sodium or potassium pyrosulfate, eventually SO3. But this involves generating SO3 at a very high temperature, which is what most people would strongly
prefer not to do.
However, you could use ammonium bisulfate itself as a reagent, e.g. rxn with bromide salts to release HBr. In this rxn, the bisulfate is about a
million times more acidic than the ammonium, so HBr should be evolved more easily than NH4Br. Ammonium sulfate may be cheaper or easier to get than
bisulfates, so this could be convenient sometimes.
|
|
|
LookingForAnswers
Harmless
Posts: 19
Registered: 25-3-2024
Location: United Kingdom
Member Is Offline
Mood: No Mood
|
|
Which method is better to make a catalyst bed? nh4vo3 or salts of vanadium organic acids? I bought v2o5 right away, but I read somewhere that
impurities in the commercial v2o5 reduce the performance of the final product. Is it true ?
Chemistry Pope
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Carbothermic Reduction of Calcium Sulfate
Hello,
I always wondered if it was possible to make something useful out of Calcium Sulfate. I heard about the process which was used in the eastern German
state to make Sulfuric Acid some time ago. When I read about this again in this thread, I decided to give it a try.
The goal is to reduce CaSO4 with Carbon to SO2 in a first step and capture it in solution. In a possible next step this can then hopefully be oxidized
to Sulfuric Acid. Maybe it is even possible to reduce the SO2 to elemental Sulfur.
Anyway, here the setup so far:
The firing chamber consists of an old soup can into which a flowerpot is inserted. On the side of the can is an opening for a pipe through which a
vacuum is applied. This tube leads to 4 wash-bottles. After that comes the vacuum pump.
Here the wash-bottles:
Gypsum pellets with carbon mixed in - interestingly they look like no carbon was added. I assume the gypsum simply covers all the carbon particles.
The pellets were made by spraying the mix with water in a protein powder drum and then rolling the drum.
At the end additional charcoal was placed on top of the pellets to get the reaction started (picture 2).
I was hoping that the reaction would be done in about an hour, but 5 hrs later I can still see glowing coal inside.
Once it is finished, does anyone have a good idea, how the SO2 can be oxidized in solution? I saw someone posted the methods from NurdRage here. Are
there any other options which are easy?
I thought maybe using some Potassium Chlorate. but then I would lose 1/3 of the acid as the Bisulfate...
Greetings!
|
|
|
Alkoholvergiftung
Hazard to Others
Posts: 197
Registered: 12-7-2018
Member Is Offline
|
|
They add Fe2O3 for Cement so Fe2O3 is an Catalyst for the SO2 to SO3.
You can make the same reaction with MgSO4 2 parts and 1 part Fe2O3 with much lower temperatures.
The reduction of SO2 to sulfur was done with H2S gas they bubblet it through H2SO3 solution.
[Edited on 4-3-2025 by Alkoholvergiftung]
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Results: Complete Failure
Thank you, Alkoholvergiftung!
I have heard about the Claus-process. Did not realize, that this reaction can also be done in aqueous solution - that is exciting :-)
Anyway, just wanted to share the results from yesterday:
I carefully sniffed the contents from the wash bottles and did not detect the distinct smell of SO2 in any of them. They smelled deliciously like
'Liquid Smoke Barbecue' sauce, lol.
Then I thought that maybe I had Calcium Sulfide in the burning chamber and took a pellet from there and tossed it in some HCl to see if I can smell
rotten eggs, but again no. So, no reaction took place.
Here some ideas for future improvement:
1. The pellets were not properly dried before the procedure - just air dried. They sure had a lot of hydration and extra water. I hoped that all the
additional coal would drive out the H2O, but it seems temperatures just did not get high enough.
2. I mixed 2 portions of coal into the gypsum - one finely powdered as the reducing agent and a second bigger portion as little chunks. This was meant
to be additional burning coal. However, the chunks were completely covered with gypsum powder and did probably not catch fire. That means next time I
will only mix the reduction coal with the gypsum and add the burning coal later.
3. The seal from pipe to can was not very good and thus there was not a lot of suction coming through the burning chamber. Maybe a copper pipe with
proper fitting would work better here.
4. As Alkoholvergiftung mentioned, a catalyst could be added to help the reaction and also convert the SO2 to SO3.
These are the ideas. I will see, if I can repeat the experiment sometime this week.
Concerning Magnesium Sulfate: 'draculic acid69' mentioned the electrolysis of MgSO4. This should be easier than carbothermic reduction since MgSO4 is
water soluble. I tried this small scale some years ago and had some success. At that time, I did not follow that path further, as I could still buy
H2SO4 at Walmart (which I still can). I will give that one another try and see if I can reproduce my findings from back then.
The important part is to have a very big surface area cathode because of all the Mg(OH)2 plating out.
|
|
|
metalresearcher
National Hazard
Posts: 790
Registered: 7-9-2010
Member Is Offline
Mood: Reactive
|
|
What about this ?
A few weeks ago I tried heating K2S2O8 (available at online shops) in a test tube (another topic here) and presumably some SO3 was released.
When capturing this into water, will that work ?
https://www.metallab.net/jwplayer/video.php?f=/forums/K2S2O8...
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
I have not tried that method, but 'garage chemist' started a threat about it:
https://www.sciencemadness.org/talk/viewthread.php?tid=5495
I am not sure though, if K2S2O8 is regulated over in Europe.
However, it can be made by electrolysis with the right anode.
|
|
|
MrDoctor
Harmless
Posts: 48
Registered: 5-7-2022
Member Is Offline
|
|
i had this idea a while ago, ill drop it here and see what people think.
I found that you can buy these cartridge heaters that can glow red hot very cheaply in the profile of a pencil, and it gave me the idea for a SO3
generator that kind of looks like a ketene lamp, in the middle of a flask you would have this heater, inside a ceramic catalyst containing some
combination of vanadium, iron, copper and molybdenum off the top of my head, and it would catalytically oxidize the refluxing contents of that flask.
I need to experiment though with these metal oxides because while i dont think i could, it would be extremely convenient if it turned out i could
combine some or all all on a single catalyst structure which could just hover above fuming sulfur and completely oxidize it to SO3 without being
poisoned in a way that cant be resolved with a heating cycle, even if this comes at the cost of efficiency. But i think it will moreso end up being
that two catalysts will be required individually.
For the time being i need to fill in some knowledge gaps like what SO3 and oleum will do at what temperatures, to elemental sulfur, as well as, if SO3
will even absorb well into hot sulfuric acid
One particular idea i had was to have a flask in which sulfuric acid and sulfur were refluxed causing the acid to form SO2, though it also produces
water and aerosolized sulfur. higher up in the flask would be a heated catalyst, i still need to figure out an arrangement which can run indefinitely
that wont be damaged by sulfur, sulfur oxides, or acid, but there shouldnt be too much sulfur present, also while i cant confirm it, i think SO3 under
these conditions just also reacts with sulfur directly as well so it might work out that as long as the catalyst is too hot for condensates and the
atmosphere is lean, that conditions should remain ideal.
As So3 forms, it should absorb into the sulfuric acid regenerating the acid and increasing its volume alongside the controlled introduction of water.
the flask would be fitted with a reflux condenser, although some considerations need to be taken so that little to no sulfur compounds leave, but
oxygen depleted air does, and i think that a reflux condenser might be able to do this at a sufficiently low flow rate. without having to go all the
way down to temps that would condense the SO2, i think there should be a way to tune the reaction such as that both S and SO2 levels remain
perpetually so low that the losses arent substantial, as well as cooling it somewhat so its density at the low air flow rate discourages escape.
For that part i have no idea how that will play out, but if its unworkable theres many other ways to do this. If however, sulfuric acid can be formed
inside a flask by the controlled addition of air, water and sulfur in a reflux, that would be dope.
otherwise its always possible to use a few stages instead to keep things clean, healthy and more efficient.
Also, if one was to use quartz glassware, it would not be difficult at all to make the catalysts using sections of steel tubing with the inside and
outside coated heavily in catalyst substrate, and then heat them inductively using a simple ZVS driver. while they rest on some low density solid such
as perlite or are strung up somehow. Soon im going to try experimenting with my ground glass quartz test tubes and other glassware, inductively
heating 2 catalyst stages to see if its possible to continuously just get SO3 output until the sulfur runs out, i think with a careful balance it
could be done, although induction heating suffers from the fact that its difficult to measure how hot you are heating things, you have to use a
literal thermocouple (a piece of metal conducting heat away from the EMF area) or some contactless IR method.
[Edited on 5-3-2025 by MrDoctor]
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Hello MrDoctor,
It sounds like you are trying to solve many problems at once.
I think your thoughts are valid, but as an amateur chemist you can solve only so many problems...
Might be best to break the problem down into smaller tasks and then start experimenting, but I think that is the route you wanted to take anyway.
Do you have a drawing of how you picture your apparatus?
Best Greetings!
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Magnesium Sulfate Electrolysis
As promised in a previous post, here a setup for MgSO4 electrolysis:
1240g Epsom salt were dissolved in ~2.5 liters of water.
This mixture was then filled in both the cathode and anode compartment.
The hope is that [SO4 2-] ions migrate from the cathode compartment into the anode compartment and the [Mg 2+] ions do the opposite.
Cell is running since yesterday and current has stabilized at:
1.9A @ 17.5V
I was hoping for a much smaller voltage with the big metal bucked as cathode.
When I tried this some time ago, I used a small stainless-steel pot and was running at 1A @ 12V, which seems like the same magnitude - so this is not
better.
I will let this run as long as it works and see where it brings us.
Here a close-up of when the cell was filled:
[Edited on 5-3-2025 by Mister Double U]
|
|
|
teodor
National Hazard
Posts: 985
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
There are a lot of ways, but what should be the solvent?
1. You can get SO3 by sulfamic acid decomposition and then react it with water to get H2SO4.
2. You can get sulfamic acid crystalls in sulfuric acid solution by reacting SO2 (you have it) with hydroxylammonium sulfate water solution (you don't
have it yet), but it requires 4 atm pressure or suitable catalist, probably you can do some research:
(NH3OH)2SO4 + 2SO2 -> NH2SO3H + H2SO4
The bonus side is that you can concentrate H2SO4 you get as the result of the reaction up to 82-85% and then react it with SO3 from the first
procedure to get concentrated H2SO4 or oleum.
3. You can get hydroxylammonium sulfate by Raschig process, reacting SO2
(you have it) with NH4NO2 (you don't have it yet).
4. You can get NH4NO2 by reacting ammonium carbonate (you can buy it) with nitrogen oxides (you cannot buy it).
5. You can get nitrogen oxides by catalitic decomposition of ammonium persulfate (you can buy or make it by electrolisys).
6. If you ever reach this point stop here and just use the lead chamber process. Try to optimize it to use the ammonium persulfate directly without
separating nitric oxides.
|
|
|
MrDoctor
Harmless
Posts: 48
Registered: 5-7-2022
Member Is Offline
|
|
| Quote: Originally posted by Mister Double U | Hello MrDoctor,
It sounds like you are trying to solve many problems at once.
I think your thoughts are valid, but as an amateur chemist you can solve only so many problems...
Might be best to break the problem down into smaller tasks and then start experimenting, but I think that is the route you wanted to take anyway.
Do you have a drawing of how you picture your apparatus?
Best Greetings! |
Its more like, i have a very generalized approach id like to take, and based on how that goes, everything else could change quite a bit and the
possibilities branch out, since one of the only constraints here is, making this work in regular boro glassware, or with very limited amounts of
ground glass quartz, if any.
I dont have any drawings but ill post some pictures once i get a working combustion catalyst working to solve the problem of oxidizing sulfur without
fire or sulfuric acid since both of those introduce challenges that gets in the way of automation.
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Update On: Carbothermic Reduction of Calcium Sulfate - Partial Success
Yesterday I did another run with the following changes:
1. Pipe was sealed to the can with plaster - same which is used as Calcium Sulfate for the reaction.
This adhered very well and was still tight after the run completed.
Picture shows the content of burning chamber before ignition and addition of top coal layer.
2. The weight ratio of coal to CaSO4 was changed to 2:1 - huge excess of coal. This was done so the mixture itself could burn. Both reagents were
mixed as powders and then sprinkled with a little water to create a coarser powder which would allow gas flow between the particles.
A top layer of coal was added for ignition.
3. The first wash-bottle which did not have a dispersion frit was left out of this setup.
Results:
Again, there was no noticeable smell of SO2 in the wash-bottles. However, when HCl was poured on the ashes from the burning chamber, there was a lot
of bubbling and smell of H2S occurred. So, some reaction took place :-). I wonder, if I had let this run longer if the formed CaS would have liberated
SO2 on oxidation, which I could have recovered from the wash bottles (I stopped the experiment, because I needed to go to bed).
Here a picture from the acidified ashes:
Side note: The raw CaSO4 was also acidified for comparison to the ashes and it showed also some bubbling. However, there was no H2S smell present at
all.
[Edited on 6-3-2025 by Mister Double U]
|
|
|
imidazole
Harmless
Posts: 22
Registered: 18-10-2014
Location: Massachusetts, USA
Member Is Offline
Mood: Radical
|
|
SO2 aq Oxidization to Sulfate salts -> Salt pyrolysis
I'll have to find my notes, but long ago I thought of a long, annoying cyclic method to produce Sulfuric acid from Sulfur/ SO2.
I think the initial idea was to bubble SO2 and air from a Sulfur burner into a solution of basic Iron/Copper mud/water with a bunch of silica gel
(might cause issues) to promote adsorption, with CuCl2 (like in the Copper Chloride process) as a catalyst to oxidize the SO2 into SO4 ions, and keep
the initial ph a little higher. the idea was that the low ph would keep the SO2 in solution as SO3 2- ions, where they would be oxidized to the
Sulfate salts. Ferric Sulfate heats through multiple stages, dehydrating and eventually releasing Sulfur Trioxide around 580C , which when bubbled
into water creates Sulfuric acid.
lots of work, but once you have such a thing built, you could make fairly large amounts of sulfuric acid, and you could concentrate it beyond what the
Copper Chloride process permits.
Overly ambitious addition:
This is where it's obvious I came up with this before hearing of "less is more"
Adding Halogens like Chlorine, Bromine, Iodine to this or the Copper Chloride process oxidizes things faster, and could be used to get valuable
mineral acids at the same time, so it could be more worth doing if you wanted to make HBr for instance.
Thinking off that, I thought about adding a chloride electrochemical cell to the mix, producing Cl2, which would become HCl, which, as the reaction
continues, would reach a saturation point and would then would release Hcl as a gas. The gas could be fed back into the electrochemical cell to react
with the NaOH produced from electrolysis, essentially becoming a loop where the end reaction is make hydrogen gas, and oxidize Sulfur Dioxide to
Sulfuric acid.
[Edited on 6-3-2025 by imidazole]
|
|
|
imidazole
Harmless
Posts: 22
Registered: 18-10-2014
Location: Massachusetts, USA
Member Is Offline
Mood: Radical
|
|
Possible carbothermic reduction boost?
I vaguely remember some long as heck method involving SiO2 to liberate something or other, I forget the exact reaction but it was something like C+Ca
Oxyanion +SiO2 -> element+Calcium Silicate
Only sort of related, but I thought it would be worth mentioning
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Thank you, imidazole.
It seems like using Copper as an oxidant in one way or another might be useful in this quest.
I guess it is fair to say that this is a 2-tiered problem:
1. Generation of significant quantities of SO2 from easily available materials.
2. Oxidation of SO2 to H2SO4 - preferably in aqueous solution, because then the SO2 can be captured prior to starting the second reaction.
|
|
|
Alkoholvergiftung
Hazard to Others
Posts: 197
Registered: 12-7-2018
Member Is Offline
|
|
https://dingler.bbaw.de/articles/ar113036.html#:~:text=Braun...
Interesting link in German of produktion of sulfuric acid without nitric acid only with an pumic stone catalyst impregnated with MnO2. There are
pictures (drawings) of the sulfuric acid plant. The Factory is from 1848 so it is simply made.
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Update On: Magnesium Sulfate Electrolysis
Brief update on the MgSO4 electrolysis.
By the end of yesterday the current had dropped to 1.5A @ 17.5V.
I added 50g of Na2SO4 in the hope that the NaOH generated at the cathode would make the Mg(OH)2 deposit more conductive, but this did not help much.
I then thought that the clay diaphragm might be the problem and held the anode directly in the cathode compartment and the current immediately jumped
to 10A - so, the clay pot is not very conductive with these parameters.
I remembered, that when I first did this experiment some years ago, the stainless-steel pot I used was shallower than the clay pot and that I had the
liquid level ~4cm higher in the anode compartment compared to the cathode compartment.
So, I removed 500ml from the cathode side and the level is now ~5cm lower than the anode side.
I also took the whole anode compartment out of the bucket an let it sit on its own so the diaphragm could 'soak' in the anolyte - which I believe is
more conductive because of the H2SO4 formed.
This morning I started the cell back up and it is now running at 2.1A @ 12V.
This is a significant improvement.
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Danke, Alkoholvergiftung!
Is an interesting read. Do you think that small quantity of Ammonia which is in the process gas is oxidized to Nitric Acid and then oxidizing the SO2
like in the Lead Chamber Process?
|
|
|
Alkoholvergiftung
Hazard to Others
Posts: 197
Registered: 12-7-2018
Member Is Offline
|
|
Yes I thought that too. That the Ammonia is oxidized to nitric acid. But the drawings are interesting. So you sew how an simple factory works. On an
other read they wrote with mich enough survace earea you dont need NOx because the SO2 can oxidize on the walls of the lead chamber. But it wasnt
recommend because of the longer time an much more lead impurities in the acid.
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Update 2 On: Magnesium Sulfate Electrolysis
Cell is running even better today. Had to dial down the Voltage:
3.28A @ 8.7V
The liquor in the cathode compartment is looking orange now. I assume [Fe3+] ions are leached out of the clay diaphragm by the H2SO4 which is
generated in the anode compartment. Difference between the liquid levels is still 5cm (anode - cathode). I did fill up some of the evaporated liquid
in the cathode compartment (outside) with the prior removed cell liquor.
[Edited on 7-3-2025 by Mister Double U]
|
|
|
Mister Double U
Harmless
Posts: 19
Registered: 21-2-2025
Member Is Offline
|
|
Update 3 On: Magnesium Sulfate Electrolysis - Successful
This morning, I stopped electrolysis and started processing the content of the cell. When I stopped the current flow the cell was running at 3.98A @
7.2V. Difference in liquid level was still 5cm (anode - cathode).
At the end of distillation 249g of an oily clear liquid were collected.
Density was 1.79g/cm^3 which corresponds to ~86% concentration.
This means that 214g of H2SO4 were formed.
Electrolysis ran for 5.5 days at an average current of 3A (estimated).
With a 2-electron reaction assumed the theoretical amount would have been 724g -> Current efficiency = 29%.
After distillation there was a lot of MgSO4 left in the flask, so it would have been wise to let the cell run a couple more days.
The H2SO4 formed did have some SO2 dissolved in it (smell). However, I am unable to quantify this amount.
Here some pictures:
Liquid in the anode compartment:
Mg(OH)2 crust in cathode compartment (with Fe contamination):
Anode Liquor (extracted through siphoning):
Flask after Distillation:
Receiving Flask:
Effect of Liquid on Paper Towel:
So, this process does work but is not very efficient.
+ Chemicals easily available (Epsom salt).
+ Materials easily available (assumed DC power source present). As anode even Lead could be used (assumption).
+ Unconsumed MgSO4 can be reused.
- Low current efficiency.
Y'all have fun!
[Edited on 9-3-2025 by Mister Double U]
|
|
|
| Pages:
1
..
3
4
5 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
