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Author: Subject: H2SO4 by the Lead Chamber Process - success
Lambda-Eyde
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[*] posted on 16-8-2010 at 13:45


Quote: Originally posted by un0me2  

Also works with I2 & Br2.


Are you sure that iodine would be a strong enough oxidant for this?
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[*] posted on 16-8-2010 at 19:47


Quote: Originally posted by un0me2  
SO2 is soluble as all fuck in water - the clathrate is insoluble till it melts (about 15-20'C IIRC) but collection of the clathrate should give a strong solution of SO2/H2O in about the right proportions. Actually, dissolving SO2 in water is endothermic, so be aware of that. I've seen STRONG solutions of SO2 in water, they stink like hell (yum SO2), but a strong solution is quite workable. Adding more SO2 as the oxidation proceeds should be feasible.


Around 20 deg., about 10g SO2 will solubilize in 100g H2O. Which isn't that much. If you were bubbling SO2 into just H2O to form H2SO3, most SO2 could be lost (depending on your ratios), or unless it was recovered (maybe by setting up a series, but it's too much work), so it might work better to just oxidize the SO2 directly using Cl2 and water. That's if you care about the sulfur loss. If not then it would be easier to gas the H2SO3 solution with Cl2 because you would be working on generating and maintaining gas flow for only one gas.

If you were to solubilize the SO2, you could do it like this: first solvate estimated SO2 in an large excess of cold water so you don't need to worry about solubility. Then take the solution and gas it with Cl2. The fumes that come over collect with water which would then contain HCl, some HClO and Cl2. Then boil the H2SO3 oxidized solution just to get out the water (probably not until white fumes form), some more HCl should come over then. Finally, boil the aq. HCl containing solution to purify it, around 20% concentration will be reached, when you go beyond this, HCl strength will actually decrease on boiling.

Alcohol and ether solubilize at least over two times more SO2 than H2O, and I've handled those solutions. No fun, no fun at all. Up there with liquid ammonia.

Quote: Originally posted by Lambda-Eyde  
Are you sure that iodine would be a strong enough oxidant for this?


Br2 and I2 are one of the oxidants which are listed to oxidize SO2 in Gmelin's Handbuch (it was mentioned on page 5 of this thread).

[Edited on 17-8-2010 by Formatik]
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[*] posted on 16-8-2010 at 21:14


Do you think bubbling Cl2 + SO2 into water could get H2SO4 concentrated past azeotropic? I don't remember if there are any side reactions that could occur in a mixture of HCl, conc. H2SO4, Cl2, SO2 and H2O... I can't think of any off the top of my head, besides formation of sulfuryl chloride, but that should not occur without a catalyst right?

[Edited on 17-8-2010 by 497]
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[*] posted on 17-8-2010 at 02:53


What does this have to do with the lead chamber process?

I note that chlorine was well known in the past, and they chose to use the dearer nitrous fumes.

The solubility of SO2 in water is of course discussed in Mellor.

EDIT: It may well have been used at the later stages, and was patented earlier.

Early patents didn't seem to catch on. References from Lunge: http://books.google.com/books?id=RAhCAAAAIAAJ
1854: http://books.google.com/books?id=SgALAQAAIAAJ&pg=PA503
1863: http://books.google.com/books?id=-T4oAQAAIAAJ&pg=PA39
1904, German:
http://v3.espacenet.com/publicationDetails/originalDocument?...
http://v3.espacenet.com/publicationDetails/originalDocument?...

But I think the reference you want is:
http://dx.doi.org/10.1002/ange.19230365503

[Edited on 18-8-2010 by S.C. Wack]
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[*] posted on 19-8-2010 at 21:37


References in Gmelin verify the reaction goes as thought: when SO2 and Cl2 are led into water, this exotherms a bit and accumulates the H2SO4 as the HCl concentration decreases. Neumann described the reaction is going rapidly and almost completely (95-100% theoretical amounts were converted), the sulfuric and hydrochloric acids result immediately as fine droplets/fog, these are difficult to absorb and also pass over, as gases and water initially interact.

The patent mentioned of Stolle, leads same parts SO2 and Cl2 into water, eventually raising the temperature to 250 deg., yielding 90% H2SO4 and conc., free from Cl2 and SO2, aqueous HCl. Neumann's process is much more descriptive.

Neumann also described despite having used a Cl2-excess, a significant amount of SO2 got solubilized in H2SO4, since SO2 solubility increases with H2SO4 concentration. Though experiments also showed conc. H2SO4 which had Cl2 or SO2 solubilized in it, after blowing in air for 15 minutes, were almost completely removed.

Quote: Originally posted by S.C. Wack  
What does this have to do with the lead chamber process?


It seems this thread is the designated stickied sulfuric acid thread. I would retitle it as the sulfuric acid preparation thread, or remove the non-Chamber discussions and sticky those with said title instead. Good eye on that reference, I also found it through Gmelin.

Quote: Originally posted by 497  
Do you think bubbling Cl2 + SO2 into water could get H2SO4 concentrated past azeotropic? I don't remember if there are any side reactions that could occur in a mixture of HCl, conc. H2SO4, Cl2, SO2 and H2O... I can't think of any off the top of my head, besides formation of sulfuryl chloride, but that should not occur without a catalyst right?


I doubt it's of concern. Neumann described that after the reaction heat slows down, that the gases come out ununited. This heat is especially large when water is first consumed in the reaction. Their later experiments used additional heat (60-92 deg), to make the reaction go much faster.

Concerning the concentration of H2SO4 obtained by combination of SO2 and Cl2 with H2O, Neumann says it is that of the Chamber acid or Glover acid (66-88%). That's the raw figure then, it can be concentrated further by regular means. For practical purposes, instead of H2O, conc. HCl was recommended. Then when a specific gravity of 1.6 is reached, the hydrochloric acid content has been nearly completely removed.

Attachment: Gmelin Cl, 102.pdf (715kB)
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Attachment: DE353742.pdf (45kB)
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Attachment: Neumann, Z.ang.Ch.36,377.pdf (1.5MB)
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[Edited on 20-8-2010 by Formatik]
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[*] posted on 24-8-2010 at 00:29


Nice route to clean H2SO4 & clean HX acids but... Anyway, the only reason it was posted was to allow those who didn't realise the alternative existed. Personally I like the construction tips on the Lead Chamber-type processes, but yeah...



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[*] posted on 19-9-2010 at 22:41


in reference to the earlier posts about sulfur with potassium nitrate in a chamber with water...could this be done without the KNO3 if you substituted oxygen gas from say the mini welding kits at home depot? they sell small bottles of welding oxygen that might be useful in this process without forming the nitrogen compounds from the KNO3. any thoughts?
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[*] posted on 19-9-2010 at 23:09


Quote: Originally posted by Rogeryermaw  
..... without forming the nitrogen compounds from the KNO3. any thoughts?


Nitrogen compounds (NO and NO2) are an integral part of the process, check the chemistry!
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[*] posted on 28-9-2010 at 04:29


hello guys
it is very important to me

can i convert sulfuric acid to hydrogen sulfide????????







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[*] posted on 27-10-2010 at 16:01


I have a granulated sulfur product designed for garden use that is 90% pure elemental sulfur, I have ground this up to powder in the past and I assume I could make it into cakes, but is this too impure for the "lead chamber" reaction here, assuming of course, that the impurities are inert, which may not be true?
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[*] posted on 27-10-2010 at 16:11


The normal additives are clay, mostly as a handling and flammability reduction, and wetting agents, more commonly used with powdered forms. The clay will be no problem except that it makes burning more difficult. Organics such as wetting agents will result in some generation of H2S as the sulphur is heated, and in carbonaceous gun forming in hot liquid sulphur - not a problem when just combusting the S.

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[*] posted on 16-11-2010 at 22:59


I was very excited about This method for H2SO4 production when i first heard of it, but unless there is absolutely no other way for you to get H2SO4 i would not consider it as a practical method.
Using a 'backyard' style setup it is almost impossible to avoid contact with the noxious fumes emmited from the burning sulfur/KNO3 mix. Is not good for the lungs etc.
On top of this the procedure must be repeated >10 times to get a very dilute solution of acid. It may be practical on a huge scale, but anything under a 100 L volume container is a waste of time.
Living in an urban enviroment it is a great way to draw attention from the neighbours, considering burning sulfur is not exactly the most subtle of smells..
I have given to this method the best of my abilities and have come to the conclusion that effort >> reward.
I understand that acid is not available to everyone, but I will happily pay the 10 or so dollars for 1 L of battery acid and boil it down over making the stuff if the option is there.
Just my 2 Cents.
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[*] posted on 16-11-2010 at 23:28


Well guess what? those two cents didn't help me in the slightest in progressing this process. Thanks for your time but if your going to tell us to buy it your wasting our time.




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[*] posted on 17-11-2010 at 00:53


Thats my point exactly, this process is unfeasable for diy at home H2SO4 production.
Practical experimentation has proved this.
I have thought of ways to improve this process, but all require more effort than warranted for an otc product.
The 'Lead chamber Process' sounds cool, but it is just not an efficient way to make sulfuric acid.
Major improvements will have to be made before it becomes anything more than a last resort.
If i have offended anyone, too bad. This is honest feedback and I make no appologies for telling the truth.

[Edited on 17-11-2010 by BenZeen]
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[*] posted on 19-11-2010 at 04:32


So, tried this on a very very minor scale the other day, and had little to no success. I was using a 2L plastic bottle as the vessel, with a small beaker with 40g S and 80g (I think) of KNO3 sitting inside on top of 100ml of water. Obviously ridiculously small scale, but I thought it'd be fun to give it a shot. Problem was, I couldn't get the sulfur to burn properly: I could get a weak flame, but not enough for the chunks to catch, and certainly not enough to decompose the KNO3. I was only using a cigarette lighter to try and light it though, so I'm thinking of getting a small pen-sized propane torch to try again with a bit more heat.
Also, it turns out that blu-tac is not quite as good a sealant for containing SO2 as I'd thought it would be :)
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[*] posted on 28-11-2010 at 15:12


The cheapest source of a sulfur compound for me would be sodium persulfate. I plan to be making sulfuric acid by sulfur dioxide and chlorine. Making sulfur trioxide from the persulfate seems very troublesome and low yielding. Can anyone think of a route to sulfur dioxide from persulfate? If I only could get bisulfite, metabisulfite or sulfur but those aren't nearly as cheap or available.
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[*] posted on 23-12-2011 at 11:59


spam reported to woelen



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[*] posted on 16-2-2012 at 18:43


I just thought of a interesting (but requiring safety precautions and nevertheless potentially dangerous) way to prepare H2SO4 from household chemicals: Bleach (NaClO), Epsom salt (MgSO4.7H2O), ammonia water and Acetic acid. (Please comply with local laws)

Step 1. Prepare HOCl. React NaOCl and Acetic acid (HAc), or select another preparation method. Distill to collect HOCl (note, as a significant part of the HOCl and Cl2O come over first, distill half of the solution to concentrate the HOCl). Repeat the distillation to further concentrate. One could also use acetone to extract the HOCl (see Patent 3078180 for details).

NaClO + HAc --> NaAc + HOCl

Step 2. Prepare (NH4)2SO4. Add NH4OH to MgSO4 and filter out the white Mg(OH)2 precipitate. If not pure NH4OH, heat it and drive the NH3 gas into the aqueous MgSO4 solution (it is important to remove any organic based additives from the ammonia water):

MgSO4 + 2 NH4OH --> Mg(OH)2 (s) +(NH4)2SO4

Step 3. Cautiously add small quantities of HOCl to aqueous Ammonium sulfate. Employ an appropriate open vessel to limit the potential condensation of any explosive NCl3 vapors and also address the foaming reaction. Apply gentle heat and perform the reaction outdoors as the reaction is expected to also produce toxic gases (chloramines). Take safety precautions in the event of exploding droplets of NCl3 which have been accidentally allowed to collect. One would expect the following:

(NH4)2SO4 + 2 HOCl ---> 2 NH4ClO + H2SO4

but ammonium hypochlorite is so unstable one only observes an exothermic decomposition reaction finally resulting in NCl3 vapors and water. In particular, the reaction chain between NH4OH and HOCl is given as:

A. MonoChloramine
NH4OH + HOCl --> NH2Cl + 2 H2O

B. DiChloramine
NH2Cl + HOCl --> NHCl2 + H2O

C. TriChloramine (or Nitrogen Trichloride)
NHCl2+ HOCl --> NCl3 (g) + H2O

Note: If any Nitrogen trichloride remains in solution, a hydrolysis occurs:

NCl3 + 3 H2O → NH3 (g) + 3 HOCl

and if any ammonia dissolves, the reaction chains starts again. EDIT: For safety, I recommend targeting the required amount of HOCl needed to react with (NH4)2SO4 to produce only MonoChloramine, which is to be driven off with heat.

Upon completion, the remaining solution should be dilute H2SO4 (note, this is not a synthesis of NCl3 as we are limiting the amount of HOCl employed and avoiding vapor condensation). Why still a potentially dangerous synthesis, working with SO3 is not exactly safe either, and both routes should only be attempted by skilled chemists taking appropriate safety precautions.
-----------------------------------------------------------------------------
Alternate method for preparation: Just treat an aqueous solution of concentrated (NH4)2SO4 with Cl2. Avoid an excess of Chlorine to limit the possible formation of an explosive oily liquid. Upon heating, the final product should be H2SO4 and HCl. One could use Ag2O to remove the HCl (via AgCl) and further encourage the reaction move to the right, or follow the procedures in other synthesis (like Cl2 + H2SO3) that also result in H2SO4 and HCl formation. This variation could form more concentrated H2SO4. An optional advanced modification would be to first run the Cl2 over heated Na2CO3 (180 C) to change the chlorinating mix to Cl2O/Cl2 thereby reducing the HCl formation, and then cautiously react with aqueous Ammonium sulfate only (dry (NH4)2SO4 is reputedly explosive with any strong oxidizer including Cl2O and conc HOCl).


[Edited on 17-2-2012 by AJKOER]
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[*] posted on 16-2-2012 at 19:12


Quote: Originally posted by AJKOER  
I just thought of a interesting (but requiring safety precautions and nevertheless potentially dangerous) way to prepare H2SO4 from household chemicals: Bleach (NaClO), Epsom salt (MgSO4.7H2O), ammonia water and Acetic acid. (Please comply with local laws)


You've got a sense of humour, when should sulphuric acid be handled without care and precaution? When should bleach and ammonia not be handled without precaution?

The mere reiteration of safety shows you have no experience and there is no guarantee this method would ever work.




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[*] posted on 16-2-2012 at 21:06


White Yeti:

Here in the wonderful USA, one is not legally allowed to use one's household reagent's in any manner prohibited by the bottle's label (a Federal offense no less).

Excuse my legal prose, but in a land where your neighbor's want to sue you for no reason (yes, it has happened) and the Patriot Act has strip us of due process to "protect" us from whoever, one is advised to tread/speak like a lawyer 24/7, or else, I would be a dumb ass. Now, unfortunately, I just sound like a dumb ass.

FYI, I do believe that the chlorination synthesis is both doable and potentially powerful when properly executed. Reason, chlorine water is:

Cl2 + H2O <---> HOCl + H(+) + Cl(-)

and the process is not limited by the initial concentration of HOCl, as in the first synthesis, as the Chlorination is ongoing. Also, water is consumed as the reaction is forced to right.
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[*] posted on 17-2-2012 at 00:42


Quote: Originally posted by AJKOER  
I just thought of a interesting (but requiring safety precautions and nevertheless potentially dangerous) way to prepare H2SO4 from household chemicals: Bleach (NaClO), Epsom salt (MgSO4.7H2O), ammonia water and Acetic acid. (Please comply with local laws)

Step 1. Prepare HOCl. React NaOCl and Acetic acid (HAc), or select another preparation method. Distill to collect HOCl (note, as a significant part of the HOCl and Cl2O come over first, distill half of the solution to concentrate the HOCl). Repeat the distillation to further concentrate. One could also use acetone to extract the HOCl (see Patent 3078180 for details).

NaClO + HAc --> NaAc + HOCl

Step 2. Prepare (NH4)2SO4. Add NH4OH to MgSO4 and filter out the white Mg(OH)2 precipitate.

MgSO4 + 2 NH4OH --> Mg(OH)2 (s) +(NH4)2SO4

Step 3. React Ammonium sulfate and HOCl taking precautions. First, use a very flat vessel with low walls to limit potential condensation of explosive NCl3 vapors. Apply gentle heat and perform the reaction outdoors as the reaction is expected to also produce toxic gases (chloramines). Take safety precautions in the event of exploding droplets of NCl3 which have been accidentally allowed to collect. One would expect the following:

(NH4)2SO4 + 2 HOCl ---> 2 NH4ClO + H2SO4

but ammonium hypochlorite is so unstable one only observes an exothermic decomposition reaction finally resulting in NCl3 vapors and water. In particular, the reaction chain between NH4OH and HOCl is given as:

A. MonoChloramine
NH4OH + HOCl --> NH2Cl + 2 H2O

B. DiChloramine
NH2Cl + HOCl --> NHCl2 + H2O

C. TriChloramine (or Nitrogen Trichloride)
NHCl2+ HOCl --> NCl3 (g) + H2O

Note: If any Nitrogen trichloride remains in solution, a hydrolysis occurs:

NCl3 + 3 H2O → NH3 (g) + 3 HOCl

and if any ammonia dissolves, the reaction chains starts again.

The remaining solution is just dilute H2SO4 (note, this is not a synthesis of NCl3 as we are avoiding vapor condensation). Why a potentially dangerous synthesis, working with SO3 is not exactly safe either, and both routes should only be attempted by skilled chemists taking appropriate safety precautions.
-----------------------------------------------------------------------------
Alternate method for preparation: Just treat an aqueous solution of concentrated (NH4)2SO4 with Cl2. The final product will be H2SO4 and HCl. One could use Ag2O to remove the HCl (via AgCl) and further encourage the reaction move to the right. This variation could form more concentrated H2SO4.


[Edited on 17-2-2012 by AJKOER]


More HOCl chemistry?!?!?!?!??!?!?!?!??!

But this method do seem interesting though. I have one question, how to I keep the HOCl from spontaneous decomposing into HCl, which decomposes further into chlorine gas?
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[*] posted on 17-2-2012 at 00:52


Distilling hypochlorous acid? Are you joking? If it is, it isn't very funny. Also, using Ag salts for preparing sulfuric acid? That's not really a good joke either.



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[*] posted on 17-2-2012 at 02:03


AJKOER, you must be realistic. As Lambda-Eyde already states, HOCl is not something you want to distill. I once have read about distilling this, but this can only be carried out with great difficulty and lots of losses. HOCl is notoriously unstable and easily decomposes into a mix of HCl, O2, H2O and Cl2!

Making H2SO4 along the route you propose is the most difficult and most expensive route which I have ever seen :D .

But please, if you really think that this might give useful results, try it and report on it. The experiment may have high educational value and if you use small quantities the risks associated with the experiment are not that high.




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[*] posted on 17-2-2012 at 12:35


Quote: Originally posted by woelen  
AJKOER, you must be realistic. As Lambda-Eyde already states, HOCl is not something you want to distill. I once have read about distilling this, but this can only be carried out with great difficulty and lots of losses. HOCl is notoriously unstable and easily decomposes into a mix of HCl, O2, H2O and Cl2!

Unless I am missing something, the reference posted by S C Wack here says that HClO can be distilled without great difficulty.

The link to the reference is active if you go back to that post:

Quote: Originally posted by S.C. Wack  
It can be distilled. Boric acid is preferred.

Attachment: jcs_101_444_1912.pdf (699kB)


[Edited on 27-9-2009 by S.C. Wack]
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[*] posted on 17-2-2012 at 13:44


Quote: Originally posted by weiming1998  
More HOCl chemistry?!?!?!?!??!?!?!?!??!

But this method do seem interesting though. I have one question, how to I keep the HOCl from spontaneous decomposing into HCl, which decomposes further into chlorine gas?


On questions relating to the distillation and stability of HOCl, see Watt's Dictionary Chemistry, page 16:

"A dilute solution of HClO may be distilled with partial decomposition, the distillate is richer in HClO; Gay-Lussac found that, on distilling a dilute solution to one half, the distillate contained five-sixths of the total HClO"

http://books.google.com/books/reader?id=ijnPAAAAMAAJ&dq=...

Per another recent source (page 552):

"Relative Volatility.
Hypochlorous acid is more volatile than water and aqueous solutions can be distilled to yield solutions of higher concentration"

http://www.scribd.com/doc/30121142/Dichlorine-Monoxide-Hypoc...

With respect to stability, my recollection is that HOCl's stability is a function of medium (water versus polar solvents) and inversely related to age, concentration (dilute solution are more stable), temperature, pH and UV light exposure. Also, contamination with organic matter and certain heavy metals can greatly expedite decomposition.

More recent literature employs the use of polar solvents (see cited Patent reference) to extract HOCl and achieve higher concentrations than observed in aqueous environments.

Please note that I have edited the original synthesis to stress using the precise amount of HOCl (and no excess) to consume the (NH4)2SO4, and target the formation of only Monochloramine, which is to be driven off by heating. Best done by carefully adding and stirring small amounts of HOCl to concentrated aqueous (NH4)2SO4 (and not the reverse, which could result in the creation of some NCl3 as an explosive yellow oily liquid). Also, never react dry (NH4)2SO4 with any strong oxidizer (explosion hazard). Also, while the chlorination of (NH4)2SO4 is perhaps a better route, by not properly controlling the amount of Cl2 generated, one could easily and tragically form NCl3 as this is one of its cited modes of preparation.

To be honest I am not too enthusiastic to work with Chloroamines, and am first also investigating a possibly safer synthesis (well at least absence the Chloroamines) involving again aqueous (NH4)2SO4, but using H2O2 (of various strengths) in the presence of mild heat and a catalyst (Pt among others).

As a sidebar, the concept is based on the apparent observation that upon boiling NH4OH + H2O2 + catalyst, the creation of NH4NO2 is apparently accomplished (a given reference is completely in German, however). One can actually, however, perform a demo by adding a mineral acid like H2SO4, which results in the evolution of NO and NO2.

2 NH4NO2 + H2SO4 --> (NH4)2SO4 + 2 HNO2

2 HNO2 --> NO + NO2 + H2O

Since we are forming H2SO4 from (NH4)2SO4, this new synthesis then becomes a path to HNO2/HNO3 and not H2SO4. However, it is also possible that NH4NO2 is formed and mostly decomposes on boiling being highly unstable decomposing into N2 gas:

NH4NO2 (aq) ---Heat---> 2H2O + N2

leaving just H2SO4. Or, we could get a mixture of H2SO4 and HNO2 with NO and NO2 fumes, or neither (I am still working on it obviously).

Caution: Ammonium Nitrite is acutely toxic compond, and as a solid, the unstable NH4NO2 (a two hour half-life, decomposing between 60 to 70 C) is considered a high explosive with limited applications due to thermal and shock sensitivity.


[Edited on 17-2-2012 by AJKOER]
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