Sciencemadness Discussion Board - H2SO4 by the Lead Chamber Process - success

H2SO4 by the Lead Chamber Process - success

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Lambda-Eyde - 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?

Formatik - 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]

497 - 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]

S.C. Wack - 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]

Formatik - 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: Neumann, Z.ang.Ch.36,377.pdf (1.5MB)
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[Edited on 20-8-2010 by Formatik]

un0me2 - 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...

Rogeryermaw - 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?

Xenoid - 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!

aeacfm - 28-9-2010 at 04:29

hello guys
it is very important to me

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

FrickinA - 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?

not_important - 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.

BenZeen - 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.

Sedit - 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.

BenZeen - 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]

Jianaran - 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

bluetrain - 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.

White Yeti - 23-12-2011 at 11:59

spam reported to woelen

AJKOER - 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]

White Yeti - 16-2-2012 at 19:12

Quote: Originally posted by AJKOER

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.

AJKOER - 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.

weiming1998 - 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?

Lambda-Eyde - 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.

woelen - 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

.

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.

entropy51 - 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]

AJKOER - 17-2-2012 at 13:44

Quote: Originally posted by weiming1998

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]

AJKOER - 19-2-2012 at 19:16

A SAFE TEST RUN ON THE CHLORINATION OF (NH4)2SO4 TO MAKE DILUTE H2SO4 (WITH Na2SO4 IN THIS TEST)

I thought of a somewhat safe procedure for a small test run on the chlorination of Ammonium sulfate to make H2SO4. This procedure is only intended to gain insights (or suggest improvements) as to the original cited synthesis. Note, this particular preparation is intended to quickly and somewhat safely create dilute H2SO4 containing Na2SO4 in this instance. The process was:

1. Prepare (NH4)2SO4 per the reaction of NH4OH on a slight excess of Epom Salt (MgSO4.7H2O). Filter out the Mg(OH)2 to obtain a clear solution of (NH4)2SO4.

2. Add NaHSO4 to NaClO/NaCl (6% Bleach) to create only the required amount of Chlorine. As a departure from the original synthesis, one immediately add this solution mix to (NH4)2SO4 from Step 1, which now also includes a small amount of 3% H2O2 (note, precise quantities are listed for this test run below) in a large (3 liter) clear plastic bottle.

Some of the reaction equations:

NaHSO4 + NaOCl --> Na2SO4 + HOCl

NaHSO4 + NaCl --> Na2SO4 + HCl

Cl2 + H2O <---> HCl + HOCl

and, per Watt's for dilute Chlorine water solutions only:

HCl + HOCl + H2O2 --> 2 HOCl + H2O

where an excess of H2O2 is to be avoided (as it would further reduced the HOCl to HCl and O2). Reference: "Watts' dictionary of chemistry", Volume 2, page 16, link: http://books.google.com/books?id=ijnPAAAAMAAJ&q=HOCL#v=o...

The key speculated reaction in this synthesis is:

(NH4)2SO4 + 2 HOCl + 2 H2O --> 2NH2Cl + 4 H2O + H2SO4

After the reaction is completed, one can add more H2O2 to remove unwanted NH2Cl (this is a suggested improvement to the original synthesis in converting HCl to HOCl, and also possibly decomposing Chloroamines):

2 NH2Cl + 2 H2O2 --> N2 + Cl2 + 2 H2O

Note, no heat was applied to the solution to reduce the possibilty of the NCl3 formation.

--------------------------------

More precisely, here are the quantites employed:

MgSO4 11.1 grams 6.6 ml
NaOCl/NaCl 110 grams 100 ml
NH4OH 30.7 grams 33 ml
H2O2 2.76 grams 3 ml
NaHSO4 10.85 grams 3.96 ml
93.2% NaHSO4 11.64 grams 4 ml

However, preparation of the (NH4)2SO4 via MgSO4 (which also includes the addition of 10 ml H2O to MgSO4 and an about 3 ml of H2O2) upon filtering suffered a loss in the amount of Ammonium sulfate available (the solution went from to a total of 46 ml to 35 ml after filtering). I proceeded anyway with the synthesis. The reaction upon mixing the (NH4)2SO4 and Chlorine generating solution at first displayed white smoke upon pouring into the 3 liter vessel. Letting noxious gases escape for about a minute (do outdoors, these were Chlorine and I would guess Chloroamines and perhaps Nitrogen), I compressed the plastic container (to test for continuing gas expansion) and sealed it. A cloud was formed that dissipated after several minutes of shaking, but a greenish tinct was observed. After an hour, a more intense noticeable greenish color was observed, which even with shaking, remained. I decided at this point to prepare and add more (NH4)2SO4 to make up for the original shortfall. Immediately, upon adding 11 ml of the new (NH4)2SO4/H2O2 mix, all visual evidence of Cl2 disappeared, however I did not witness a significant gas generation as previously. The whole reaction I found surprisingly less violent than I anticipated (but, for a more concentrated preparation, this may not be the case). The final product was a clear solution with a strong Chlorine smell with seemingly greater viscosity (most likely testing acidic as well given the chlorination). I plan on testing it further starting by freezing the solution.

[Edited on 21-2-2012 by AJKOER]

Kola - 21-2-2012 at 11:29

I don't know much about lead Chamber process but I know H2SO3 can be safely prepared in a laboratory by dissolving SO2 in water. SO2 can be easily prepared by heating sulphur &Oxygen in the presence of a catalyst eg Pt. The H2SO3 can be oxidized with H2O2 to yield H2SO4...any problems with this Scheme?

White Yeti - 21-2-2012 at 12:28

Quote: Originally posted by Kola

I don't know much about lead Chamber process but I know H2SO3 can be safely prepared in a laboratory by dissolving SO2 in water. SO2 can be easily prepared by heating sulphur &Oxygen in the presence of a catalyst eg Pt. The H2SO3 can be oxidized with H2O2 to yield H2SO4...any problems with this Scheme?

Expensive, the sulphur is not too expensive, but the hydrogen peroxide is definitely expensive when working on small scales.

Kola - 21-2-2012 at 13:57

Quote: Originally posted by White Yeti

Expensive, the sulphur is not too expensive, but the hydrogen peroxide is definitely expensive when working on small scales.

expensive but less risky, besides I figure O2 can do the oxidation but the resulting solution will be dilute cos of SO3 evolution..and perhaps the reaction wont go to completion

White Yeti - 21-2-2012 at 15:07

Quote: Originally posted by Kola

expensive but less risky,

Hydrogen peroxide is dangerous mind you. The 3% is safe, but once you go over 30% (the lowest concentration for this to be useful), things get dangerous.

Of course, 18M sulphuric acid is not by any stretch of the imagination a compound to be handled without care.

Also, sulphur dioxide is dangerous. Do not treat it as an ordinary gas, it will choke you if you do not employ proper ventilation. If you have asthma, this gas will not have mercy, so be careful.

AJKOER - 22-2-2012 at 14:35

UPDATE ON THE TEST RUN (CHLORINATION OF AMMONIUM SULFATE)

First, I previously stated that the key speculated reaction in this synthesis of dilute H2SO4 was:

(NH4)2SO4 + 2 HOCl + 2 H2O --> 2NH2Cl + 4 H2O + H2SO4 [1]

I have since found a reference stating that stock solutions for the purpose of testing Chloramine levels are prepared by adding NaOCl (Bleach) to (NH4)2SO4. This supports the speculated key reaction if one replaces Na with H (but not as to speed as NaOCl is clearly more ionic than HOCl). See page 52, top left, at:

http://www.fwrj.com/techarticles/0608%20tech%205.pdf

However, I also found an interesting equation that suggests a portion of the H2SO4 could be consumed by liberated NH3:

"The overall chloramine decomposition reaction, based on acid-catalyzed disproportionation reaction of NH2Cl (Eq 2.43):

3 NH2Cl + H+ --> N2 + NH4+ + 3 H(+) + 3 Cl(-)" [2]

Source: "White's Handbook of Chlorination and Alternative Disinfectants" by Black & Veatch Corporation, page 150 (middle of page):

Link:
http://books.google.com/books?id=mGVbIoW2lNAC&pg=PA150&a...

If we rescale Equation [2] to correspond to Equation [1] and add H2SO4:

2 NH2Cl + 2/6 H2SO4 + 4/6 H2SO4 --> 2/3 N2 + 2/6 (NH4)2SO4 + 2 HCl + 4/6 H2SO4

which suggests that 1/3 of the expected H2SO4 yield is reduced by the presence of Monochloramine. However, the author also cites the following reaction (also on page 150):

"NHCl2 + NH2Cl --> N2 + 3 H(+) + 3 Cl(-)"

So, if we added more HOCl, then per the equation:

HOCl + NH2Cl ===> NHCl2 + H2O

(as a reference of the above and many other associated reactions, see: http://www.h2o4u.org/chloramination/chemistry.shtml )

then, some NHCl2 could be created, which correspondingly reduces the NH2Cl. This would increase the H2SO4 yield as upon adding more HOCl, we have

NH2Cl + (HOCl + NH2Cl) + 4 H2O + H2SO4

= NH2Cl + (NHCl2 + H2O) + 4 H2O + H2SO4

Or:

(NH2Cl + NHCl2) + 5 H2O + H2SO4 = N2 + 3 HCl + 5 H2O + H2SO4

So the new target reaction is:

(NH4)2SO4 + 3 HOCl + 2 H2O --> N2 + 3 HCl + 5 H2O + H2SO4

and in the direct application of HOCl synthesis (as opposed to chlorination) the best one could hope for is apparently dilute H2SO4 in HCl.

Also, the original targeting of the production of NH2Cl is not advised as the Monochloramine is more stable than NHCl2 which, in contrast, is much more volatile and easier removed. In fact, prolonged boiling is required for NH2Cl's decomposition combined with aeration only slowly decomposing it, and only very advanced/costly filtering being able to remove it from solution.

In genral, one should also be aware that the chemistry of Chloramine is complex with the species of chloramine rapidly and constantly shifting as a function of temperature, pH, turbulence, and Cl2/NH3 ratio. See: http://www.chloramine.org/literature_pdf/chloramine_facts_06...

------------------------------------
Update Frozen Solution

An interesting salt has separated (tiny long thin crystals which in solution shimmer and resembles cotton fibers). Some bubbling also in a perfectly clear solution with a strong chlorine-like smell.

AirCowPeaCock - 22-2-2012 at 17:45

These guys do NOT like Chloramines(;

AJKOER - 23-2-2012 at 19:31

Actually, I do not like Chloramines either. If in California or (or Massachusetts) and you see a green color in the local swimming pool, avoid getting the water (actually Chloramine treated) in your eyes, and certainly don't let a pet drink it (toxic to dogs and fish). Apparently, Campden pills which releases SO2 in water are able to remove Chloramines in about a minute.

------------------------------

I thought it was interesting to note if we take the new target equation scaled by two and net the H2O:

2 (NH4)2SO4 + 6 HOCl --> 2 N2 + 6 HCl + 6 H2O + 2 H2SO4

and noting that:

6 HOCl <--> 3 Cl2O + 3 H2O

we have:

2 (NH4)2SO4 + 3 Cl2O --> 2 N2 + 6 HCl + 3 H2O + 2 H2SO4

Note, the treatment of dry Ammonium Sulfate with any strong oxidizer (as DiChlorine Mono-oxide) is strongly not advised (at worst, explosive), and it is intended that a concentrated aqueous (NH4)2SO4 solution be treated.

Also, upon adding more Cl2O:

2 (NH4)2SO4 + 6 Cl2O --> 2 N2 + 6 HCl + 6 HOCl + 2 H2SO4

but:

6 HCl + 6 HOCl <==> 3 Cl2 + 3 H2O

So:

2 (NH4)2SO4 + 6 Cl2O --> 2 N2 + 3 Cl2 + 3 H2O + 2 H2SO4

So assuming this very vigorous reaction can even be performed

, a massive treatment of a concentrated (NH4)2SO4 solution with Cl2O gets one at best 40% H2SO4 (and a lot of Chlorine). Caution, DiChlorine Mono-oxide is many times more poisonous than Cl2 and has explosive properties as well.

Kola - 26-2-2012 at 15:26

Quote: Originally posted by AJKOER

Actually, I do not like Chloramines either. If in California or (or Massachusetts) and you see a green color in the local swimming pool, avoid getting the water (actually Chloramine treated) in your eyes, and certainly don't let a pet drink it (toxic to dogs and fish). Apparently, Campden pills which releases SO2 in water are able to remove Chloramines in about a minute.

------------------------------

I thought it was interesting to note if we take the new target equation scaled by two and net the H2O:

2 (NH4)2SO4 + 6 HOCl --> 2 N2 + 6 HCl + 6 H2O + 2 H2SO4

and noting that:

6 HOCl <--> 3 Cl2O + 3 H2O

we have:

2 (NH4)2SO4 + 3 Cl2O --> 2 N2 + 6 HCl + 3 H2O + 2 H2SO4

Note, the treatment of dry Ammonium Sulfate with any strong oxidizer (as DiChlorine Mono-oxide) is strongly not advised (at worst, explosive), and it is intended that a concentrated aqueous (NH4)2SO4 solution be treated.

Also, upon adding more Cl2O:

2 (NH4)2SO4 + 6 Cl2O --> 2 N2 + 6 HCl + 6 HOCl + 2 H2SO4

but:

6 HCl + 6 HOCl <==> 3 Cl2 + 3 H2O

So:

2 (NH4)2SO4 + 6 Cl2O --> 2 N2 + 3 Cl2 + 3 H2O + 2 H2SO4

So assuming this very vigorous reaction can even be performed

this scheme is too complex...i don't think its logical for any laboratory chemist. Maybe can there be a simpler pathway?
Axhandle's synthesis is much more realistic and it can be used on a large scale.
Have you experimented your theory? What were your results?

AJKOER - 29-2-2012 at 20:58

Kola: I agree with your comment. The most practical approach is also the simplest, mentioned in the original synthesis, just over chlorinate (NH4)2SO4 forming HCl and H2SO4. To be more precise, per the equation:

2 (NH4)2SO4 + 6 HOCl --> 2 N2 + 6 HCl + 6 H2O + 2 H2SO4

upon adding 6 HCl to each side:

2 (NH4)2SO4 + 6 HOCl + 6 HCl --> 2 N2 + 12 HCl + 6 H2O + 2 H2SO4

and as:

6 Cl2 + 6 H2O <---> 6 HOCl + 6 HCl

2 (NH4)2SO4 + 6 Cl2 + 6 H2O --> 2 N2 + 12 HCl + 6 H2O + 2 H2SO4

or:

2 (NH4)2SO4 + 6 Cl2 --> 2 N2 + 12 HCl + 2 H2SO4

or, more precisely working with aqueous Ammonium sulfate:

(NH4)2SO4 + x H2O + 3 Cl2 --> N2 + 6 HCl + H2SO4 + x H2O [CAUTION: No heat to avoid the creation of an oily yellow NCl3 explosive]

at which point, removal of the HCl (without heating) via dilute H2O2, aeration,.. becomes the concern, without significantly diluting the H2SO4 or excessive expense. Note, adding FeSO4 to bleach has been suggested previously on a Sciencemadness thread as a means, without acid, to generate Cl2.

Nevertheless, the synthesis is far from perfect with smelly toxic fumes and stills provides much more HCl (to be addressed) than H2SO4.

[Edited on 1-3-2012 by AJKOER]

Formatik - 29-2-2012 at 23:42

Quote: Originally posted by AJKOER

(NH4)2SO4 + x H2O + 3 Cl2 --> N2 + 6 HCl + H2SO4 + x H2O [CAUTION: No heat to avoid the creation of an oily yellow NCl3 explosive]

at which point, removal of the HCl (without heating) via dilute H2O2, aeration,.. becomes the concern, without significantly diluting the H2SO4 or excessive expense. Note, adding FeSO4 to bleach has been suggested previously on a Sciencemadness thread as a means, without acid, to generate Cl2.

It's full well possible that the nitrogen trichloride might not be skipped as an intermediate and accumulates regardless of overchlorination, in which case it would have to be decomposed by standing for 24 hours, maybe even less. It would make it an immense explosion hazard. Cold temperatures also favors nitrogen trichloride formation. And then if the reaction goes as thought you may be able to isolate some sulfuric acid.

For generation chlorine, oxidizing hydrochloric acid has been the preferable chemical route.

AJKOER - 1-3-2012 at 14:23

Formatik:

In my test run, I did avoid an excess of Chlorine (via stoichiometric calculations) and without heating, observed no yellow oily liquid. Given that the formation of the deadly NCl3 is highly endothermic, these measures I would think, are necessary first safety measures. Note, strong light may serve as a heat substitute (as well as a trigger), and should also be avoided.

Now, NCl3 reputedly has limited solubility in water, but with continuous solution turbulence (dangerous, I would otherwise believe) and aeration (commercially employed to remove NCl3 in so called Breakpoint Chlorination of water) may speed up the hydrolysis/removal of any Nitrogen trichloride that has formed.

Still, I think your comment is to some extent appropriate ("immense explosion hazard"), and hence my reluctance to make especially the chlorination path the recommended synthesis route.

As a source on dealing with unwanted NCl3 see: "White's Handbook of Chlorination and Alternative Disinfectants" by Black & Veatch Corporation, page 23, where the authors mention breakpoint chlorination followed by aeration, and an alternate method of pre-treating the chlorine gas with UV light in the spectrum 3600-4400 (my speculation on how this works is as NCl3 is know to explosively decompose with UV light that this treated Cl2 decomposes the NCl3 on formation). Link:

http://books.google.com/books?id=mGVbIoW2lNAC&pg=PA23&am...

See also, for example, Patent 4435291, "Breakpoint chlorination control system", which does mention the use of mechanical agitation.

http://www.patentgenius.com/patent/4435291.html

[Edited on 2-3-2012 by AJKOER]

Formatik - 2-3-2012 at 19:51

What I was thinking when saying nitrogen trichloride formation is favored in cold temperatures was in reference to hydrolysis of nitrogen trichloride which increases with temperature, whereas when it is kept cool then nitrogen trichloride better accumulates. I had phrased it badly, cold does not promote its formation but slows its hydrolysis. Ideal temperatures for forming nitrogen trichloride are cool to warm, but not freezing.

AJOKER, it's a good sign you observed no yellow oily liquid. However, nitrogen trichloride is an elusive compound not only because of decomposability, but because it needs to be let sit to see it (typically after 20 minutes), because it can be in finely divided form where especially also milky appearance of solutions makes it harder to see.

The reaction between ammonium sulfate solution and chlorine could be:

6 Cl2 + (NH4)2SO4 --> 2 NCl3 + H2SO4 + 6 HCl

Then NCl3 could possibly react with an excess ammonium sulfate per this equation:

(NH4)2SO4 + 2 NCl3 = 2 N2 + 6 HCl + H2SO4

Note: similarly that if NH4Cl is present in large excess the following reaction has been observed to occur slowly with ammonium chloride: NCl3 + NH4Cl = N2 + 4 HCl (Bray, Dowell, J. Am. Soc. 39 [1917] 905).

Thus we could eventually have something like the:

(NH4)2SO4 + 3 Cl2 = N2 + 6 HCl + H2SO4

Dilute H2SO4 decomposes NCl3 rapidly, but dilute HCl is said not to decompose NCl3 (according to Davy, Phil. Trans. 103 [1813] 1). However both acids when concentrated decompose the trichloride rapidly under N2 evolution according to the same reference (these might also act as trigger initiators to the pure compound when added to it). So there is the possibility that as nitrogen trichloride is formed, can be decomposed by dilute H2SO4 before it has the chance to accumulate (speculation on assuming it forms in an atmosphere corresponding to dilute sulfuric acid). Though Chapman and Vodden (J. chem. Soc. 95 [1909] 138) describe NCl3 as solubilizing more rapidly in aq. HCl than in water or aq. H2SO4, because HClO which is formed in equilibrium reaction of water hydrolysis, is destroyed by HCl but not water or sulfuric acid. The presence of both acids seems to have some promotional decomposing effect.

Note in some points it may look conflicting but it may really be more complicated with nitrogen trichloride. I've formed it in acidic ammonium sulfate solution and it did not decompose right away (probably too dilute of an acid concentration), in fact it could be isolated.

Addendum: we may be interested in destroying any NCl3, and there are substances which are reported to prevent the formation of NCl3 mentioned in Gilb. Ann. 47 (1814) 58 coal powder, CO2, sulfur powder, etc. But if the nitrogen trichloride is needed for the reaction to work it could be in that sense, counter-productive to destroy it.

[Edited on 3-3-2012 by Formatik]

AJKOER - 3-3-2012 at 07:28

Formatik:

Some interesting stuff (at least to me).

OK, I found a source with interesting comments on the decomposition of NCl3.

First, the decomposition reaction:

NCl3 + H2O → NHCl2 + HOCl

is first order in time, so your comment on letting NCl3 sit for hours is supported.

Next, the source cites work by Kumar that NCl3, in the presence of excess HOCl, will either decompose into HOCl or be subject to reduction/oxidation. Noting that Chlorine water is:

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

I would argue that upon adding a small amount of NH3 (via the hydrolysis of (NH4)2SO4), the reaction is first moved to the right forming more HOCl, as much more rapidly:

H(+) + Cl(-) + NH3 --> NH4Cl

which would thus add to the decomposition of NCl3 per the above author's assertion on the effect of increasing HOCl (see also H2O2 comment below).

In addition, the author's assertion on the ability of HOCl to decompose NCl3 adds support to my equation, to quote:

Quote: Originally posted by AJKOER

UPDATE ON THE TEST RUN (CHLORINATION OF AMMONIUM SULFATE)
............

So the new target reaction is:

(NH4)2SO4 + 3 HOCl + 2 H2O --> N2 + 3 HCl + 5 H2O + H2SO4

and in the direct application of HOCl synthesis (as opposed to chlorination) the best one could hope for is apparently dilute H2SO4 in HCl.

However, on old Sciencemadness thread (http://www.sciencemadness.org/talk/viewthread.php?tid=3347 ), the implication is clearly that one can directly formulate NCl3 with a sufficient amount of HOCl (so carefully monitor/restrict the quantity of HOCl to keep all your body parts!). To quote:

Quote: Originally posted by kazaa81

Hi,

in roguesci's forum I found a different synthesis of NCl3, with acetic acid, sodium hypochlorite and ammonium nitrate (probably any ammonium salt would be right). I've reformatted this and post it now:
==
Nitrogen chloride

3 NaOCl + 2 CH3COOH + NH4NO3 -----> NCl3 + NaNO3 + 2 CH3COONa + 3 H2O

although an excess of HOCl may disrupt the created Nitrogen trichloride. Note, Acetic acid plus NaOCl yields HOCl, so this is a reworking of the synthesis. Also, replacing a sulfate with a nitrate in a pure HOCl reaction could produce dilute HNO3.

---------------------------

Finally, your comment on accumulation of NCl3 is in agreement with the author's statement that NCl3 can be quite stable in the presence of excess Cl2. However, excess Cl2/H2O equals HCl/HOCl, and in the presence 3% H2O2, per my updated synthesis, this is mostly HOCl and may hinder NCl3 without even assuming that H2O2 directly attacks the Nitrogen trichloride.

Source: Per page 14 at link:

http://scholar.lib.vt.edu/theses/available/etd-042299-143911...

[Edited on 4-3-2012 by AJKOER]

[Edited on 4-3-2012 by AJKOER]

AJKOER - 4-3-2012 at 10:52

Apparently Chloramines can be decomposed by super chlorination, potassium peroxymonosulfate (2KHSO5·KHSO4·K2SO4), and ozone. The latter two processes support my contention that dilute H2O2 may be effective on NCl3. To quote:

"Chloramines can be removed from the water by the following three methods:

By adding a high dose of chlorine and raising the levels to 10 times the level of combined chlorines (5 to 10 ppm) for a minimum of 4 hours. This is called super chlorination. To remove chloramines, the ratio of chlorine to chloramines must be at least 7.6 to 1. If this ratio is not obtained more chloramines will be produced.
By adding a non-chlorine shock to the water. The most common chemical used for this is potassium peroxymonosulfate (MPS). This "shocking" requires the addition of 1 oz. per 625 gallons of water.
By adding ozone to the water. If an ozone generator is installed and wired so that it comes on each time the pump comes on, then oxidation of the ammonia and nitrogen compounds will take place on a continuous basis. This reduces, and can even eliminates the need for shocking. Each time ammonia and nitrogen enter the ozonated water, they are oxidized by the ozone."

Link:
http://www.rhtubs.com/chlorine.htm

SO3 generator.

Cemtek - 18-5-2012 at 17:32

I'm in need of a small SO3 generator to produce a known quatity on SO3 (between 1ppm & 100ppm) from SO2 & H2O for testing & calibrating a laser based SO3 analyzer.
Does any one know of a company that makes something like that?

elementcollector1 - 16-9-2012 at 09:19

I bought a bottle of Rooto's H2SO4 at McLendon's yesterday, and after measuring 100mL on a scale, it weighed 168.9 g. This translates to a density of 1.69, or 78%.
However, MSDS says 93.2%.
Which one is the correct value?
(Does my scale need calibration?)

ScienceSquirrel - 11-12-2012 at 03:53

Quote: Originally posted by elementcollector1

I bought a bottle of Rooto's H2SO4 at McLendon's yesterday, and after measuring 100mL on a scale, it weighed 168.9 g. This translates to a density of 1.69, or 78%.
However, MSDS says 93.2%.
Which one is the correct value?
(Does my scale need calibration?)

I would check the scale by weighing 100g and 200g of water. I have done this with my 0.1 and 0.01g scales using a 100ml pipette and they are spot on.
If you use a measuring cylinder it will be nowhere near as accurate but it should do.
It is drain opener and i doubt most end users care if it is ca 80% or 90%, it will still open drains so I would take the MSDS with a pinch of salt!

Poppy - 28-12-2012 at 17:49

Not that this equation should be proposed at all. After all, the way up here has shown uses of bleach, so let come the exotric equations!!!
I've been interested on making tripped equations out of standard electrode solutions, such an equation should provide means of a different contact for implanting the damn oxygen to sulfur dioxide, the problems found is, first, to surpass the thumb-rule 0,6V that should give aprecciable reaction rates:

SO4(2-) + 4H+ + 2e- --> SO2 + 2H2O +0.17 (reversing it)

SO2 + 2H2O --> SO4(2-) + 4H+ + 2e- -0.17
2Fe3+ + 2e- --> 2Fe2+ +0.77V

_____________________________________________
SO2 + 2H2O + 2Fe3+ --> SO4(2-) + 4H+ + 2Fe2+ +0.5V

furthermore quite acidic solutions are already needed to bear Fe3+, so the equilibrium would shift, which could be compensated with pressurized SO2 until the overvoltage is achieved.

Seriously, why is the direct reaction so slow?
K = 1.08
E° = 1,03.10^(-3)V

(From harsh calculations) I think the harsh Fe3+ enviroment would increase chances of success.

White Yeti - 29-12-2012 at 08:34

Using Fe⁺³ is really overkill. Why not use oxygen?
O2(g) + 2 H2O + 4 e− ----> 4 OH−(aq)

Poppy - 29-12-2012 at 10:41

Fe3+ enchances surface contact. A 1 molar solution is in a steric proportion of 1 to 55 iron/ H2O
oxygen bubbles would hardly beat this
Indeed, the Fe3+ feedback is provided by bubbling oxygen, but since Fe2+ oxydises faster in the presence of oxygen than does SO2, iron as a catalyst...

White Yeti - 29-12-2012 at 14:52

I see the reasoning, but your previous post made it seem as though you were using ironIII as the oxidizer, not as a catalyst, otherwise you would have included oxygen in the overall reaction as well.

Try it out and see how it goes.

Poppy - 29-12-2012 at 20:13

Quote: Originally posted by White Yeti

I see the reasoning, but your previous post made it seem as though you were using ironIII as the oxidizer, not as a catalyst, otherwise you would have included oxygen in the overall reaction as well.

Try it out and see how it goes.

Sure, then pick in the apparatus to be given a try to this.
Direct oxygen doesn't work because it would flush, I mean, just vent the SO2 away.
It must proceed alligator style, step by step (wait, did this make any sense?)

testimento - 13-7-2013 at 06:02

What are methods of producing sulfur dioxide the most straightforward way from OTC sources? OTC means for ex. sodium sulfate, potassium sulfate, etc.

One idea was to flush the salt with very hot air(1000-1500C) to cause it to pyrolyze into oxide and sulfur dioxide. Another idea was to roast it with carbon to form sulfide, react with acid to form H2S and burn it to form SO2.

Jay - 18-7-2013 at 12:27

The most straightforward way to get sulphuric acid is to BUY IT AS BATTERY ACID. The price is pretty good. Do you guys not have it available in your area?

Zyklon-A - 21-12-2013 at 16:47

I've looked through this thread, and I haven't had this question answered; How high % sulfuric acid can you get, if you ran the reaction multiple times with exess KNO3 and S, before you boil it?

DubaiAmateurRocketry - 22-12-2013 at 17:14

I find the SO3 process much more convienient/ promising for us amateurs.

Mesa - 22-12-2013 at 18:05

@Zyklonb: Your question was answered on the same page it was posted. It's also on the first 5 links seen when typing "Lead chamber process" into google and hitting enter.

@else: Would it be possible to selectively adsorb aforementioned chloramines on something akin to molecular sieves? I was looking into the preparation of surface modified cellulose type membranes which seems quite achievable for home chemists(starting as far back as unrefined wood pulp. Probably better off buying some cotton though.)

Zyklon-A - 23-12-2013 at 06:49

It doesn't say anything about, % on the fist page, on Wikipedia, it says, ''Later versions of chamber plants included a high temperature Glover tower to recover the nitrogen oxides from the chamber liquor, while concentrating the chamber acid to as much as 78% H2SO4'', but I just have a 5 gallon bucket, not a ''Glover tower''.

DubaiAmateurRocketry - 23-12-2013 at 14:43

Quote: Originally posted by Zyklonb

It doesn't say anything about, % on the fist page, on Wikipedia, it says, ''Later versions of chamber plants included a high temperature Glover tower to recover the nitrogen oxides from the chamber liquor, while concentrating the chamber acid to as much as 78% H2SO4'', but I just have a 5 gallon bucket, not a ''Glover tower''.

Read this.

http://www.sciencemadness.org/member_publications/SO3_and_ol...

I strongly suggest this method for sulfuric acid. It can problem high concentration oleum or sulfuric acid.

Zyklon-A - 23-12-2013 at 15:48

Thank you, I could not find that, earlier.

testimento - 10-2-2014 at 18:27

I have been thinking few different approaches of manufacturing of sulfuric acid. All of them uses calcium sulfate as a precursor. It is first reduced with carbon to calcium sulfide, which is heated with some calcium sulfate to yield sulfur dioxide and calcium oxide. Since sulfur dioxide is highly soluble in very cold water (up to 230g per liter), it could be stored there before use. Another possibility is to liquify it into a good freezer(-30C minimum) where a common steel BBQ gas bottle is placed with proper quality needle valve. The valve can be closed when the procedure is done and bottle brought to room temp, where it gains about 5-10 bars of pressure and pressure boiling point equilibrium keeps it stable. Common steel is happy with SO2, but it must be dried with CaCl2 trap just in case.

From there, SO2 can be used for few methods to cause sulfuric acid. The "easiest" method is to bubble air through the SO2-water and slowly heat it to cause SO2 vaporization and lead this composition through CaCl2 to dry it. After this it would be lead through, preferably, quartz tube, with either powdered or impregnated vanadium pentoxide catalyst, where it would turn into SO3. This would be bubbled through 70-90% sulfuric acid which is heated up to 70-90C, and the SO3 is turned into H2SO4 and eventually oleum. SO3 must not be condensed(note 46C BP), because it can form strange matter alpha/beta structures that may spontaneously decompose upon heating with such force it can shatter glass. The unreacted SO2 can be vented off, or preferably directed through another V2O5 tube and bubbled into another pot of sulfuric acid. It could also be bubbled through water storage trap or freezer cold trap to obtain total yield of 90-100% with no SO2 losses.

Another variation is to do the heating and air bubbling the same way, but lead the SO2 into nitric acid tank, where H2SO4 is formed and NO is released, which is lead into a middle tank similar to Ostwald process reaction tank(2 NO + O2 = 2 NO2), and the NO2 is partially condensed and lead into water where nitric acid is formed again. This structure could be made high so the heavier-than-air-NO will remain there and oxidize again to NO2, condense and form nitric acid. Through this effective yields of 100% from SO2 and 60-80% from HNO3 can be obtained. Downside is the required equipment, upside is that everything is mostly OTC. In most simple terms, this process requires only basic labware, DIY electric furnace, some SS and quartz tubing and some V2O5 and can easily fit on a decent tabletop.

[Edited on 11-2-2014 by testimento]

jock88 - 11-2-2014 at 09:54

DO be careful with SO3.
Eyes gone if you make ONE mistake.

testimento - 11-2-2014 at 10:39

A chemist who isn't using full face respiratory mask is a candidate for darwin awards.

Xenon1898 - 26-3-2014 at 19:41

Quote: Originally posted by Samkk

Sulfuric acid production is described via equipment process flowsheet to easy understand the mechanism of the process

Just a nit, but this process flowsheet doesn't have an exit path shown for the acid product.

macckone - 27-3-2014 at 08:04

Quote: Originally posted by Xenon1898

Quote: Originally posted by Samkk

Sulfuric acid production is described via equipment process flowsheet to easy understand the mechanism of the process

Just a nit, but this process flowsheet doesn't have an exit path shown for the acid product.

There are two points for take off.
They are the concentrated acid storage tank and the chamber acid storage tank.
They are not shown on the diagram but that is where they would be removed.
You pick the take off point based on the concentration you need.

DoctorZET - 16-4-2014 at 13:12

Remember the "Bordeaux mixture" ?
Everybody can go to a agricultural-shop (farm-shop) and buy some CuSO4 and Ca(OH)2 , who are the main ingredients of the "Bordeaux mixture"-a really good fungicide.

So, we can get a lot of CuSO4 (wich is not too expensive, just like the battery acid)

Then, I know 2 methods to make SO3 (sulfur trioxide) :

first involves a tin process:

1)if you don't have tin, take some soldering alloy (Pb+Sn) and bubble the molten metals (it melts at about 170-200*C) with Cl2. Meanwhile, distill the SnCl4 gas resulted.
Pb Sn + 3Cl2 --(200*C)--> PbCl2 + SnCl4^
Now you have a pure quaternary stannium salt : SnCl4 (a fuming liquid at room temperature, wich can be used in many other purposes)

2)take some SnCl4 and some dehydrated CuSO4 and mix them until the mixture become a dense opaque very pale greenish fluid:
SnCl4 (excess)(liquid) + 2CuSO4(solid) --(time)--> Sn(SO4)2(solid) + 2CuCl2(solid)

3)then distill the excess of SnCl4.

4)heat the remaining powder up to 200-250*C and collect the SO3 vapors:
Sn(SO4)2 --(200-250*C)--> SnO2 + 2 SO3^
(the CuCl2 impurities does no effect at all)

5)to re-use the tin in this process, heat the remaining powder (SnO2+CuCl2) with some H2 gas, in a tube, to reduce the tin(vi) oxide to tin metal and water, copper will be also reduced to Cu and HCl :
SnO2 + 2H2 --(190-210*C)--> Sn + 2 H2O^
CuCl2 + H2 --(150-170*C)--> Cu + 2 HCl^
now you can start again to transform Sn in SnCl4 ... and so on.

Second method is:

1) a boring school reaction:
Fe2(CO3)3(powder) + 3CuSO4(aq) --(some time)--> 3CuCO3(solid) + Fe2(SO4)3(aq)

2) Heat a bit the Fe2(SO4)3:
Fe2(SO4)3 --(480-500*C)--> Fe2O3 + 3 SO3^

3) Now you have to convert Fe2O3 to Fe2(CO3)3:
Fe2O3(powder) + 3[H2CO3](aq) --> Fe2(CO3)3(solid) + 3H2O
now you can start again to make iron(iii)sulfate...

~the problem with first method is that it requires more atention because of the stannic chloride, but the good part is that this method requires low temperatures...
~the problem with second method is that it requires higher temperatures (than glass can resist) and that's making a problem about collecting the toxic SO3 fumes...but if you have the materials, the overall process is simple to make.

[Edited on 16-4-2014 by DoctorZET]

blogfast25 - 26-4-2014 at 04:29

This post contains so much nonsense it's hard to tell where to start debunking it.

"SnCl4 (excess)(liquid) + 2CuSO4(solid) --(time)--> Sn(SO4)2(solid) + 2CuCl2(solid)"

That doesn't work as you've basically proved yourself.

"Fe2(CO3)3(powder) + 3CuSO4(aq) --(some time)--> 3CuCO3(solid) + Fe2(SO4)3(aq)"

Fe2(CO3) in all likelihood doesn't exist but if it did how this reaction is supposed to proceed remains a mystery. Presumably '(some time)' here means all of eternity.

"[...] but if you have the materials, the overall process is simple to make."

Yes and pigs will fly!

macckone - 26-4-2014 at 05:35

If you have copper sulfate, just heat the sulfate.
It decomposes just fine.

DoctorZET - 30-4-2014 at 04:17

Yes, I just discover that SnCl4 needs to be in H2SO4 solution at 100*C in order to react ... and Sn(SO4)2*2H20 always decompose to H2SO4 and SnO2 when heated just to 80*C.
And iron(iii)oxide remain the same, even I add it to a acidulated Na2CO3 solution. (just a few FeCO3 impurities appear along with FeO(OH) )
And "some time" means 1 to 10 minutes...(just if reactants actually exist

)
The first method is stupid, I know (H2SO4 is one of the incoming reagents to make H2SO4 because SnCl4 is decomposed only by warm sulfuric acid).
The second method is good, but only if I work with Fe2O3 as the main cycle product

aga - 8-9-2014 at 16:25

Quote: Originally posted by blogfast25

This post contains so much nonsense it's hard to tell where to start debunking it

AJKOER ! Wherefore art thou ?

AJKOER - 26-3-2015 at 08:29

Here is an idea for small quantity production of Sulfuric acid I may soon try when I have the sulfur. At this point, I will discuss only the theory and leave more precise possible ways to implement till later.

Concept: Heat dry (NH4)2SO4 to sublimation along with burning Sulfur in oxygen to form Sulfur dioxide in the presence of water, where the acid gas will be employed to neutralize the newly released ammonia gas.

Reactions:

(NH4)2SO42 + Heat (to 235 C) → 2 NH3 (g) + H2SO4 (g)

S8 + 8 O2 → 8 SO2

2 NH3 + H2O + SO2 → (NH4)2SO3

Along with the bisulfite, see "SO2 Removal by NH3 Gas Injection: Effects of Temperature and Moisture Content" by Hsunling Bai , Pratim Biswas , Tim C. Keener, first page available at http://pubs.acs.org/doi/abs/10.1021/ie00029a019?journalCode=iecred

Then collect the liquid containing Sulfuric acid, some unreacted Ammonium sulfate particles and Ammonium sulfite. Then filter, and mildly heat as, per the above reference, the sulfite impurity sublimes at 60 C.

One potential advantage of this method over say reacting H2SO3 with Cl2 or concentrated H2O2 or just conc HOCl (from aqueous NaOCl, CO2 gas and CaCl2 forming aqueous HOCl which can be further concentrated by repeated distillation of half of the starting very volatile Hypochlorous acid) is, while the quantity potentially produced is small, it could be fairly concentrated.
----------------------------------------

One can appropriately extend the sublimation method to use other neutralizing gases like Cl2, Cl2O, O3, possibly Singlet oxygen, ...

[Edited on 26-3-2015 by AJKOER]

AJKOER - 5-7-2015 at 11:28

Quote: Originally posted by AJKOER

.....
Reactions:

(NH4)2SO4 + Heat (to 235 C) → 2 NH3 (g) + H2SO4 (g)
......

No success here.

Further research indicates that the above reaction is incorrect in this instance. While gases created are many, the thermal decomposition of Ammonium sulfate is not a significant source of either SO3 or H2SO4 vapors mixed with NH3.

See, for example, http://pubs.acs.org/doi/abs/10.1021/i260036a001?journalCode=... and http://onlinelibrary.wiley.com/doi/10.1002/jctb.5010200408/a... .

[Edited on 5-7-2015 by AJKOER]

Assured Fish - 16-4-2016 at 22:29

H2SO4 using cation exchange membrane

Ok this is may be a completely bogus idea but would it be possible to use a cation exchange membrane to separate 2 chambers and then put HCL solution in one chamber and a solution of NaHSO4 in the other chamber and then because the cation exchange membrane will only allow cations to pass through it i.e. H+ and Na+ then theoretically both the HCL and Bisulfate would disassociate and you would generate sulfuric acid in the bisulfate solution and sodium chloride in the HCL solution in an equilibrium reaction.

HCL + Na+ <----> NaCl + H+
NaHSO4 + H+ <----> H2SO4 + Na+

After a certain period of time you would remove the liquids and separate the sodium bisulphate from sulfuric acid (fractional distillation 315*C to 337*C), unfortunately i don't know a hell of a lot about cation exchange membranes but in theory it should work provided its energetically feasible, but i have no idea where the equilibrium constant will sit and i am waaay to tired to bother sitting there for an hour trying to find the entropy of NaCl, NaHSO4, HCL and H2SO4 and then calculate it, but has anyone got a cation exchange membrane to see if this would work.

Cheers Fish

hissingnoise - 17-4-2016 at 02:17

Quote:

Ok this is may be a completely bogus idea . . .

Indeed, Na⁺ replaces H⁺ ─ not the other way round!

macckone - 17-4-2016 at 10:22

Sodium sulfate is commercially electrolyzed in a three chamber cell with cation and anion membranes forming sodium hydroxide and sulfuric acid. Similar should be possible with diaphragms to produce sodium hydroxide and sodium bisulfate. The bisulfate can then be heated to produce sulfur trioxide.

Tin man - 17-10-2016 at 21:26

Acording to a few chemistry textbooks I have read, Nitrogen dioxide oxidizes sulphur dioxide according to the following eqaution,
NO2+SO2>NO+SO3.
Then the NO can react with oxygen according to the following reaction
2NO+O2>2NO2.
My proposed rout would be as follows:
A two necked round bottem flask is filled with nitrogen dioxide, then a gas adaptor is placed in each neck. One gas adaptor is led to a source of oxygen and the other is led to a source of sulphur dioxide. The gas flow ratio is adjusted so that the gas in the round bottem flask remains a light brown colour, and SO3 can be seen condensing on the sides of the flask. Once you have collected enough SO3 to sufficiently scare yourself shitless, you can open the flask and pippet out the SO3( under a fume hood, with a dry glass pippet) and ampule it.
Please note that I just thought of this setup, and it is surely very dangerous. My excuse for not trying it yet is that I am not in possession of a fume hood or a two neck round bottem flask. If anyone should try this, best of regards, but please do be safe.

yakoot - 17-4-2017 at 11:46

Really it is a great work. Sulfuric acid now is produced by Contact Process. The first step for contact process is to melt the sulfur and then burn it in a furnace to form SO2 .

more information
http://sulfuricacidworld.blogspot.com.eg/

[Edited on 17-4-2017 by yakoot]

clearly_not_atara - 17-4-2017 at 12:48

Mmyes conc H2SO4 is by all reports an absolute bastard to make. Producing sulfur trioxide is nearly trivial by heating pyrosulfate but doing anything with the gas or simply remaining alive in its presence requires a great deal of care.

Amides form zwitterionic adducts with SO3 yes? Perhaps a lipophilic amide can be chosen such as N-benzoylpiperidine (logP = 2.2) or oleamide, which can be used to absorb a stream of SO3, forming the amide adduct R(=N+R'R")OSO3-, and then this is hydrolysed with just enough water to separate the amide. If the hydrolysis is kept cold you should mostly avoid hydrolysis of the amide. Presumably the amide can then be extracted with octanol or toluene or something.

Another possibility or possibly what can be used together with the amide process is heating dilute sulfuric acid to about 300 C or so to concentrate it to around 80-85% and then dissolving SO3 into this until it is truly concentrated. I don't know if this is comparably exothermic to adding SO3 to water; hopefully not.

Tin man - 18-5-2017 at 16:38

I can't think of any scenario in which sulfer trioxide could come in contact with an amide and not turn it into black goo. Are you thinking of sulfur dioxide?

[Edited on 19-5-2017 by Tin man]

clearly_not_atara - 20-5-2017 at 16:44

I no longer think that will work.

However I'm still sort of interested in ways to reduce/absorb the heat of hydration of SO3. Instead of dealing with strange chemicals the following rxn seems interesting:

CaSO4*2H2O (plaster/gypsum) + 2 SO3 >> CaSO4 + 2H2SO4

with significantly less heat released than the reaction with water, but still a lot. How to perform this though?

The difficulty is not in producing SO3, which happens by simply heating bisulfate, but in capturing it without destroying it or killing the experimenter.

Radium212 - 4-1-2018 at 01:53

Nice! Finally a way for people with little equipment to produce sulphuric acid at home. Thank you.

TheIdeanator - 30-3-2018 at 18:40

Codyslab did a video on this recently and got me thinking about improvements. It would seem to me that the dissolution rate of SO3 is limited by the surface area. Why not add in one of those piezoelectric humidifier elements to get some good spray going.

Refinery - 6-5-2020 at 02:42

Calcium sulfate with sodium silicate at significant temperature should result with SO3 production that can be absorbed to 98% sulfuric acid to yield oleum and that can be then diluted?

Other sulfates are also viable with lower decomp temp like copper, but the source price is too high. CaSO4 is extremely cheap, hence resulting product can just be discarded.

I suggested the same method with nitrates. Bulk Ca, Mg and Cu nitrates can be pyrolyzed for NO2 to produce HNO3, and I figure that they could be reconstituted with two step metathesis process from another source of nitrate?

RogueRose - 6-5-2020 at 05:08

What would the product be of the calcium sulfate and sodium silicate and what temp do you think you would need, I suspect pretty high, like melting steel high?

[Edited on 5-6-2020 by RogueRose]

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