ave369
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Sulfamic acid decomposition
Hello, it's me again, with another crazy idea how to prepare sulfuric acid.
Sulfamic acid is OTC as a descaling agent. I've read somewhere that among its thermal decomposition products are SO3 and H2O (along with SO2 and N2),
hence, if you decompose it in a retort and condense the results, you will get sulfuric acid.
But I can't get the equation of the decomposition reaction anywhere. Does anyone know?
Smells like ammonia....
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Hawkguy
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2H3NSO3 -> N2, SO2, H2SO4, 2H2
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Metacelsus
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There are many ways sulfamic acid could decompose. Most simply,
H3NSO3 -> NH3 + SO3 (basically, reverse of formation, so therefore unfavored)
My guess as to the approximate stoichiometry of the actual reaction would be:
3 NH3SO3 -> 3 H2O + N2 + 3 SO2 + NH3
I don't see any way hydrogen gas could be formed, mechanistically.
[Edited on 12-10-2015 by Cheddite Cheese]
[Edited on 12-10-2015 by Cheddite Cheese]
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softbeard
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I agree with C. Cheese.
The final products would probably be a mix of N2, H2O, SO2, and maybe some SO3 (as
H2SO4 with the H2O); ie. a mess.
You'd probably be better off trying to heat dried NaHSO4 and condense the SO3 as described in the oleum threads elsewhere on the
board. Of course, ideally, you need some ~95%+ H2SO4 to absorb the SO3.
Back to electrolyzing CuSO4 solutions...
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ave369
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Quote: Originally posted by softbeard  |
You'd probably be better off trying to heat dried NaHSO4 and condense the SO3 as described in the oleum threads elsewhere on the
board |
Already done that. Found out I was running out of test tubes too fast. That's why I'm looking for alternative methods so frenetically.
Smells like ammonia....
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deltaH
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Sulfamic acid reacts with oxidants to make sulfuric acid. When the oxidant is something like nitric or nitrous acid, then sulfuric acid is cleanly
formed with evolution of nitrous oxide and nitrogen gas, respectively, but such reagents are of course not practical.
However, the lesson is that you need to oxidise the amine group away.
What about electrolysing hot concentrated sulfamic acid solutions, hypothetically generating nitrogen gas and sulfuric acid at the anode and hydrogen
gas at the cathode? You wouldn't even need a seperator for this because I don't think you can reduce sulfamic acid easily to anything.
Lead would NOT be a suitable metal for the anode because lead sulfamate is very soluble  A lead cathode, however, should work well though in terms of compatibility with the acid, but it would have a high overpotential for
the hydrogen evolution reaction, but so what for the home chemist?
Manganese dioxide anode? Might also be a good electrocatalyst for the oxidation of the NH2- group?
Anyway, all hypothetical, but give it a try if you want.
[Edited on 13-10-2015 by deltaH]
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softbeard
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Copper would probably work ok as a cathode, but selecting the anode material would be tricky. Platinum, as usual, is the no. 1 contender; until you
see how much it costs. Graphite would work, but you'll need to filter off the graphite gunk suspended in your product. Maybe MMO anodes (coated Ti),
as used for chlorate cells, would work best.
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deltaH
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MMO cannot be used below pH 1... one big drawback of it. The coating dissolves off in strong acids quickly.
Graphite would be good, filtering is not a problem, but you'd get no catalytic effect with the graphite and might just end up making oxygen on the
anode instead of the desired oxidation of sulfamic acid, but it's worth a shot considering availability and cost!
[Edited on 14-10-2015 by deltaH]
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ave369
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I'm okay with graphite gunk. It's for a Zintl-Karyakin distillation anyway, graphite will not be distilled.
Smells like ammonia....
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softbeard
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| Quote: Originally posted by deltaH | | MMO cannot be used below pH 1... one big drawback of it. The coating dissolves off in strong acids quickly. |
Thanks Delta, I didn't know that. Something to keep in mind when using these MMO anodes. I had in mind the TiO2 in the MMO would be like a
fired ceramic; and so be virtually insoluble in acid. But that's not the case...
Seems like every one of these anode materials, except platinum, has a weakness.
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deltaH
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There are some specialised MMO's that might handle it, made from 'platinum oxides'. I have two small such plates and a 5A power supply, so I think I
might try this with a dilute sulfamic acid solution. The electrodes are too precious to me to try concentrated solutions 
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softbeard
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My guess is that 'platinic oxides' should be ok in dilute sulfamic acid... notice the 'should' emphasis. You just never know if some
kind of soluble Pt(sulfamic acid)++++ complex doesn't form (+4 for oxidation state is guess).
You'd hate to see your 'platinic oxides' anode dissolving before your eyes 
(nah, kidding, I doubt it... kinda...
On that note, not even solid Pt is completely safe. One of the parameters in perchlorate production is mg's of Pt loss per ton of perchlorate.
[Edited on 15-10-2015 by softbeard]
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deltaH
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Yeah, I feel the same, I love the electrodes too much to try more than 0.5M solution.
I don't trust that nature understands the theory 
That would ultimately yield something close to pH 1 and so should be fine.
I can dissolve some lime in enough vinegar to make a clear solution and then mix this with the product of electrolysis. If the solutions turns white,
then I'd say it's evidence you made some sulfuric acid.
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teodor
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I think the main result of decomposition is some adduct of sulfur trioxide.
Experimental:
An evaporation flask was charged with 15.87 g of sulfamic acid. Long arm and Dimroth cooler were attached. The receiving flask was placed in an
ethanol bath cooled to -20C. The same coolant had been continuously pumping through the cooler. In the opening of the receiving flask a drying tube
(CaCl2) was attached.
The acid was melted and immediately the decomposition was started. The heatgun was adjusted to give air with the temperature 580C.
The snow immediately started to form on a spiral of the cooler directly from gaseous phase.
The decomposition was not particularly fast and I had no patience to decompose all the charge, at the same time I decided that I should not increase
the temperature. I will be able to provide an additional information about how much acid was decomposed in the next report.
The arm was disconnected. Every time I opened apparatus a white cloud of irritating smoke was always formed, similar to that of oleum by its visual
appearance. Every time when I opened the apparatus a bit of oily liquid appeared on glass surface near the joints.
The "snow" didn't melt at the room temperature (12 - 17C). No slightly change after one day was noticed.
Next day I started to cool ethanol again, so now I kept the receiving flask at -35C and I supply warm water (approx 40C) through the spiral. All
“snow” sublimed to the bottom of the receiving flask without any visible liquid phase.
I attached a dropping funnel charged with distilled water.
Any drop of water was making a small “explosion” filling the flask with a white cloud and raising the temperature. The oily liquid was formed on
the bottom. It will be subject for further chemical analysis.
I noticed SO2 smell going from drying tube after every water contact.
My assumption that it could be some SO3 adduct consisted of SO3, SO2 and the nitrogen which is decomposing by water loosing SO2 part. But it could be
also some polymeric form of more or less pure SO3 also.
There is one unsolved problem yet. After melting the sulfamic acid is always liquifying in a form similar to epoxy resin. It adheres to glass and
unavoidable cracks the flask upon cooling even being spread when is still liquid as a thin layer.
Probably adding some high boiling point solvent to the flask before cooling could help, otherwise some acid resistant boat is probably needed to reuse
the flask.
So, that's all for now. And I express my thanks to woelen for his old present - the bottle of the sulfamic acid I used for this experiment.
P.S. As I already suggested many times, no information from the books with the words "industrial" or "technology" should be taken as granted. It's
obviously that no any water was distilled off from the decomposition.
[Edited on 23-12-2024 by teodor]
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teodor
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So, what is that mistery that wikipedia article gives incorrect information about the sulfamic acid decomposition products.
Wikipedia uses as the reference Yoshikubo, K.; Suzuki, M. (2000). "Sulfamic Acid and Sulfamates". Kirk-Othmer Encyclopedia of Chemical Technology.
I looked into this. It says that the decomposition starts at 209C. Already incorrect (I will show later why). Then it lists SO2, SO3, N2, H2O as
decomposition products. There is no reference for the data source. It's strange. So, doing my investigation of the literature further.
And then I found exactly the same temperature and "SO3, SO2, H2O, NH3, N2" in Ullmann's encyclopedia of industrial chemistry, another source wikipedia
likes to cite (on my opinion, without any good reason). But Ullmann's at least has the reference: E. Divers, T. Haga, J. Chem. Soc, 69 (1896) 1634.
The famous Edward Divers, inventor of Divers' solution and one of the biggest contributors to NOx chemistry (nitronium salts, sulfonated nitrogen
compounds). But 1896!
Yes, actually Divers was one of prominent researches of amidosulfonic acid and the paper published in 1896 contains “a summary, not elsewhere found,
of the work of others”.
“The only statement yet made, concerning the effect of heating amidosulphonic acid in the absence of water, is Berglund’s, that, when rapidly
heated, it is decomposed; SO2, N2, H2O and H2SO4 being produced. This is correct (but) ….”
And than Divers humble adds his own investigations which are remarkable, also by means of observancy. Basically, he describes the same effects I had
observed but it looks like he was seeing much much more details and tried much more condition variations.
The sloppiness of creators of those “industrial encyclopedias” is that citing Divers they even take information not from Divers himself but from
Berglund (the man who discovered the amidosulfonic (“sulfamic”) acid in 1876 ).
I’ve attached the article for historical interest. (I can also recommend his 2 articles “Imidosulphonates” 1892 and 1896 as well as more late
reviews of this class of compounds made by Audrieth in 1940 (highly recommended), Benson & Spillane in 1980 and Spillane in 2014).
But for real study of the decomposition product we should take Zeitschrift für Chemie 1966. I see, the german language and DDR. Rare person in
english-speeking world has interest to DDR nowadays. But it is there where the decomposition of sulfamic acid was studied with accuracy in 1966. So,
the article is attached and I believe I will present my english translation with some comments a bit later.
But yes, the product I condensed
according to the study, should be pure H2SO4. I’ve shown the proof (not completed yet without chemical analysis but with agreement to W. Wanek
publication) that in my reaction conditions I’ve get mostly SO3 or very concentrated oleum (in the form of SO3 polymer). So, where is H2O go I
don’t know yet.
Let’s read the both articles!
Attachment: divers1896.pdf (1.5MB)
This file has been downloaded 56 times
Attachment: 10.1002@zfch.19660061112.pdf (274kB)
This file has been downloaded 57 times
Attachment: audrieth1940.pdf (2.4MB)
This file has been downloaded 61 times
[Edited on 29-12-2024 by teodor]
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woelen
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Very interesting find! This opens up the option of making concentrated H2SO4 from an easy to obtain chemical, and only glassware is needed, not some
specially engineered high-temperature apparatus, which are needed for other methods of making SO3.
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teodor
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There is one problem: as other investigators mention, the speed of SO3 production is constant and slow, not related to the temperature. This has 2
side:
1. The rate of distillation is quite safe for any amateur lab. In case of breaking glassware the leak of H2SO4/SO3 fumes is manageable.
2. You need to wait.
I agree that it could be suitable for getting small amounts of SO3/making SO3 adducts for organic synthesis or using SO3 "in situ".
The residue also has interesting composition.
I would try to melt sulfamic acid/ammonium persulfate mix and check how the decomposition products change, but I don't have information about
explosiveness of such mixture.
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teodor
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The translation of the german article:
On nitrogen-sulfur compounds;
The thermolysis of amidosulfuric acid
by W. Wanek, Research Institute for Inorganic Chemistry Usti nad Labem (Czechoslovakia)
Dedicated to Prof. Dr. Skramovsky, on the occasion of his 65th birthday
In connection with the thermal behavior of amidosulfuric acid HSO3NH2, it is generally stated that this compound, which in the solid phase has the
formula H3N:SO3, melts at 209 °C with complete decomposition, without it having been possible to identify the products of this process or to explain
its mechanism. A systematic pursuit of this reaction yielded the results given below.
Differential thermal analysis always shows a strong endothermic effect at 178°C. If HSO3NH2 is carefully heated in a thin glass tube to 175 to
180°C, a cloudy melt is formed which actually only clears at 209°C. However, if the cloudy melt is heated for some time at 180°C with the exclusion
of moisture, two clearly distinguishable phases form which remain even after the melt has cooled. The paper chromatographic analysis of these two
phases according to Lehmann and Kempe [2] showed that the lighter, clear phase consists of undecomposed HSO3NH2 and sulfate, while the heavier, cloudy
phase contains imidodisulfate, sulfamide monosulfate, trisulfimide and higher N-S compounds remaining at the _starting point_ in addition to these two
compounds. Above 209C both phases mix, and at
260C the paper chromatogram shows only sulfate, high
N-S compounds and small amounts of imidodisulfate and
disulfimidosulfate [O3S(NHSO2)2O]2-.
At temperatures between 300 and 340°C, dense white mists escape from the melt, and the spots of the higher N-S compounds disappear on the paper
chromatogram, while the concentration of imidodisulfate and especially disulfimidosulfate increases.
The analytical evaluation of the melts and the resulting gaseous products indicates that H2SO4 mists are released from 180°C; this H2SO4 release
increases with temperature, but is still so slow even at 340°C that even after several days of testing, the end point of this reaction could not be
determined. Above 340 °C, further decomposition of the melt occurs with the escape of SO2, NH3, H2O and N2, which leads to complete destruction with
the formation of ammonium sulfate.
As can be seen from Table 1, in the melting process in the temperature range of 200 to 300 °C, regardless of the temperature, 50% of the amido
nitrogen of the HSO3NO2 is always converted into ammonium (determined by the formaldehyde method).
In order to achieve a quantitative sequence of the reactions taking place in the melt before total decomposition, the thermolysis was carried out in a
vacuum (10-5 mm Hg) in sealed two-leg tubes and the gaseous product was frozen out at -78°C. It was found that, as shown in Table 2, in
the entire temperature range from 178 to 300°C, 1 mol of HSO3NH2 is always released for 0.25 mol of H2SO4, which can be isolated and analyzed as a
white mass in the cooled leg of the apparatus. The paper chromatogram of the hard white melts always shows only sulfate and disulfimidosulfate as the
main components. As is clear from a detailed discussion of the question of the reaction course during the melting of HSO3NH2, which will be published
in a later paper, it is necessary to consider this acid as an intermediate in the _solvent-ansolvent_ (cosolvent?) systems.
H2SO4 <=> H2O + SO3 (1)
(NH2)2SO2 <=> NH3 + SO2NH (2)
which then in this system the formulation
2HSO3NH2 = H2SO4 * (NH2)2SO2 <=> H2O + NH3 + SO3 + SO2NH (3)
The primary reaction of decomposition will then proceed in principle according to (3), whereby the water molecules have a greater affinity for SO3
[3]. The equations for the melting process can then be
2HSO3NH2 = SO42- + O2SN- + 2H+ + NH4+ = H2SO4 + O2SNNH4 (4)
The sulfimide formed is unstable in these melts and reacts further with the H2SO4 to form sulfimide sulfates
[O2SNH]x + H2SO4 <=> HO[O2SNH]xSO2OH (5)
The degree of polymerization of the sulfimide sulfates in the melt is temperature-dependent and, as high-frequency titrations of the melts in
anhydrous dimethylformamide with anhydrous alcoholic NaOH solutions showed, decreases from a value of x = 6 at 200C to a value of x = 1 at 340C. To
demonstrate the primary formation of sulfimide, the HSO3NH2 was boiled in anhydrous pyridine. A reaction product was obtained from which a fraction
insoluble in dimethylformamide with the gross composition (HNSO2)5.5 * 2C6H5N and a fraction soluble in
this medium could be isolated. From this soluble fraction it was possible to prepare pure silver salts of tri- and tetrasulfimide.
Polysulfimide is therefore primarily formed in the melt, whereby equilibria form between the individual polysulfimides depending on the melting
temperature
HNSO2 <=> … <=> (HNSO2)3 <=> (HNSO2)4 <=> … <=
(HNSO2)x (6)
which in turn are in equilibrium with the solvent H2SO4 present
(HNSO2)v + H2SO4 <=> HOSO2(HNSO2)vOH (7)
If this solvent is distilled off (e.g. in a vacuum), the melt solidifies with simultaneous complete reaction of the sulfimide with the sulfate portion
that cannot be distilled off
4HSO3NH2 -> 2HNSO2 + H2SO4 + (NH4)2SO4 -> H4NSO3(HNSO2)2ONH4 + H2SO4 (8)
The degree of polymerization of the end product is determined by the amount of cations present (NH4+). As shown in a later work, at temperatures above
150 °C the reaction of trisulfimide with sulfates, amidosulfates and sulfamide proceeds smoothly to form the corresponding linear compounds
(imidodisulfate, sulfamide monosulfate, imidodisulfamide).
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I am very pleased that the final word of the decomposition study was said by the Czech / Slovak chemist because I also know very smart Czech and
Slovak amateur chemists by this forum.
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