Sciencemadness Discussion Board - Ammonium Pyrosulfate

Ammonium Pyrosulfate

Rosco Bodine - 24-10-2005 at 11:01

I have been doing some experiments with distillation of Azeotropic nitric acid from mixtures of a slight excess of theory 74 to 84% H2SO4 and NH4NO3 , which is leaving an ammonium bisulfate " niter cake " after the distillation residue cools and solidifes . The residual solid has some sulfuric acid , water , and a small amount of nitric acid as impurity . While still hot from the distillation it is entirely liquid , and I have been wondering if once
a distillation of HNO3 has completed if some further usefulness may be made of the material by removing the receiver containing the nitric acid , and continuing the heating of the residue to a much higher temperature to convert the niter cake to a pyrosulfate .

Or perhaps the material should be removed and recyrstallized before further processing if this is really necessary .

Has anybody done any experiments in conversion of the ammonium bisulfate niter cake to ammonium pyrosulfate ?

Or does anyone have any information regarding the decomposition temperature for the ammonium bisulfate ?

The interest in the ammonium pyrosulfate is as a dehydrating agent for sulfuric acid , producing oleum in the same way as does sodium pyrosulfate , if indeed this is possible that it works the same way .

Dave Angel - 24-10-2005 at 13:34

If memory serves, ammonium pyrosulphate is a reagent suggested in a patent dealing with acetic anhydride production via the pyrosulphate process - I'm sure I had a copy of that patent somewhere but I can't seem to locate it right now!

Anyway, I think there was a benefit in using ammonium acetate/pyrosulphate rather than alkali salts in the process... easy regeneration of the pyrosulphate perhaps; the exact details escape me right now, but I hope that gives you something to go on.

Rosco Bodine - 24-10-2005 at 14:06

A melting point for Ammonium Pyrosulfate is given as 236 C in US3885024 , along with a statement of the generalization of 250 - 300 C being used for decomposition of the bisulfates to the pyrosulfates .

I have found nothing else so far .

Rosco Bodine - 25-10-2005 at 07:59

The Ammonium Bisulfate in the niter cake
is a hot liquid melt at the end of the distillation of the nitric acid , having a melting point of 147 C which is lowered somewhat by the mixture with residual sulfuric acid which was used in about 8-10 % excess for the syntheis of the nitric acid . If left undiluted with any water , the residual hot liquid mixture will set up on cooling to a monolithic rock something like set concrete . Slowly reheating the
mixture will melt it to a liquid again . Diluting the hot residual liquid with about one third its volume of water seems to be about the right amount to avoid the formation of a monolith upon cooling and provide a normal crystallization of coarse separate crystals which can be broken up and filtered . I have some of the coarse crystals of Ammonium Bisulfate filtering now .

The source of intrigue for me about this materials possible value as a dehydrating agent for sulfuric acid is similar to the interest and line of thinking with regard to the discussion of metaphosphoric acid ,
although it is unclear if the equilibrium for the hoped reaction will be favorable at the
elevated temperature where the reactants and products will remain liquid .

The idea is that upon heating the ammonium bisulfate above its melting point of 147 C and to perhaps 300 C ,
that a boiling will appear as the material dehydrates to the pyrosulfate , which will also be liquid . Conversion should be complete when boiling stops and moisture ceases to come off the hot liquid
melt of Ammonium Pyrosulfate , m.p. 236 C
according to the patent .

The hope is that if H2SO4 is dripped into the stirred molten Ammonium Pyrosulfate , that SO3 will distill , at
some range of temperature near the
236 C , but less than the higher temperature where the Ammonium Bisulfate byproduct would be cycled back to the pyrosulfate by heat driven dehydration .

If this is possible , it would be more thermally efficient than the similar process using the hydration of metaphosphoric acid , regenerated from orthophosphoric acid , as the range of extremes in temperature required for cycling the dehydration would be less , and there
would not be the problem of glassware being attacked .

This seems simple enough and is probably just one of those reactions that works fine on paper and is " too good to be true " in the reaction flask where the proof is in the pudding for the hoped reaction

Anyway , I mean to find out what is the story on this possible reaction , since the literature has provided no answers and leads directly to the question of whether or not this ammonium pyrosulfate scheme
for SO3 is workable .

Eclectic - 25-10-2005 at 15:15

Isn't runaway decomposition a potential problem using ammonium nitrate as your nitrate source (especially if there is ANY chloride contamination)?

Rosco Bodine - 25-10-2005 at 17:15

Quote:

Originally posted by Eclectic
Isn't runaway decomposition a potential problem using ammonium nitrate as your nitrate source (especially if there is ANY chloride contamination)?

I have seen no problem when using pure ammonium nitrate and 92% drain cleaner along with 1.260 electrolyte for the azeotropic nitric acid . Several experiments with distillation of HNO3 at normal pressure have given very smooth distillations with no red fumes and produce an almost clear nitric acid , so long as the process is not exposed to direct sunlight . For a 2 liter distilling flask , 700 to 800 grams of NH4NO3 is a comfortable amount so long as the heating is done using a good temperature controlled mantle . I use an Electromantle Stirrer Mantle , which allows good control by a solid state rate controller and solid state electromagnetic
rotating field drive for the rare earth stirbar

Anyway , I had at one time years ago worked out the exact proportions for azeotropic nitric from ammonium nitrate , but I can't find my notes so I am repeating the work . From the data from
three test distillation runs , it looks like the correct proportions are near to the following : for 700 grams of NH4NO3
( pure uncoated prills ) , add 200 ml of electrlolyte ( 1.260 d. H2SO4 ) and then add 500 ml 92% drain cleaner H2SO4 .
This should produce something over half a liter of azeotropic or slightly stronger HNO3 , 100% of the theoretical would amount to 558.7 ml of 69.95% HNO3 .

The distillation curve is variable because of the non-idealized mixture , and the solvent effect of the acids on the brine of salts in the reaction mixture , which interacts in a sort of oscillation of temperature across a few degrees of range which is actually above the boiling point of the azeotrope . So you may see stillhead temperatures anywhere from
124 to 130 , as the azeotrope comes off the mixture a bit superheated even at its gentlest carryover . The mixture tends to
simmer away with a slow rise in temperature well above the azeotrope until it accumulates to a certain level and then breaks free , distilling over in a steady stream for several seconds as the temperature falls from the boiling away .
And then the distillate stops coming over
except isolated drop by drop until another
temperature rise and accumulation occurs , whereupon the process repeats ,
in a sort of gentle oscillation of charge and discharge of the accumulating nitric in the distillation mixture . A cloth towel wrapped around the upper part of the flask and connecting adapters is secured with clothespins to prevent the heat loss from the glass which will otherwise see almost all of the distillate reflux instead of carrying over to the condenser .

There are patents concerning the distillation of nitric acid even as d 1.5 from
NH4NO3 and H2SO4 in high concentration , and this has been done safely without any decomposition of the NH4NO3 . I haven't used NH4NO3 myself for d 1.5 nitric , but since the temperature is actually much lower for the higher concentration of HNO3 , and I know it is good for azeotropic nitric , then I will probably try the process for the d. 1.5 nitric . It may even be safer to distill the
d. 1.5 nitric than the d. 1.41 nitric because of the lower temperature .

But no , I don't plan to add any ZnCl to the mixture

, nor would I recommend it .

[Edited on 26-10-2005 by Rosco Bodine]

neutrino - 25-10-2005 at 17:18

Any chloride should come out as HCl gas when mixed with sulfuric acid.

Eclectic - 26-10-2005 at 05:26

Neutrino: Chloride ion catalyzes the decomposition of ammonium nitrate into N2O in acidic aqueous solutions at a fairly low temperature, less than 100 C, and since it's catalytic, it doesn't take much.

Rosco: Do you have your patent numbers at hand?

Rosco Bodine - 26-10-2005 at 05:51

DE280967

GB129305

Info From Chris the Great in Uncoventional Shaped Charges thread:

Eclectic - 26-10-2005 at 07:40

I was just flipping through some of my info when I stumbled across this completely by accident. It is from The Chemistry and Technology of Explosives, Vol III, page 462

quote:
An unconfirmed hypothesis was also formulated that ammonium nitrate in the presence of sulfuric acid undergoes dehydration to the formation of nitramine (Vol. III) which is a strong and unstable explosive compound.
A double salt of ammonium nitrate and sulphate, defined by the formula 2NH4NO3-(NH4)2SO4, is capable of detonating with a rate of 1000-1400 m/sec provided an exceptionally strong initiator and completely hermetic confinement are used. However, the detonation train of this double salt has a tendency to break up after travelling a short distance.

Going to the Nitramine entry (page 15) we see some more interesting stuff:

quote:
Nitramine is believed to occur in an ammonium nitrate solution in an excess of concentrated sulfuric acid as a result of dehydration of this salt:
NH4NO3 --> NH2NO2 + H2O
Davis and Adams [8] report the following experimental observations in support of this supposition. On heating a solution of ammonium nitrate in sulfuric acid to 150*C, nitric acid cannot be distilled, but nitrous oxide is evolved, probably from the decomposition of nitramine. If, however, the solution is kept for a long time between 90 and 120*C, nitric acid can be obtained by distillion. The authors' conjecture is that the addition of water to the nitramine takes place according to reaction (16):
NH2NO2 + H2O --> NH3 + HNO3

Rosco Bodine - 26-10-2005 at 10:25

Well those who are not believers can use sodium nitrate , but for me it will always be that experimental results are better than conjecture about what distills from mixtures of H2SO4 and NH4NO3 in the presence of a limited amount of water .

HNO3 .....you either got it and know it ,
or you think not and don't , for never having tried

Eclectic - 26-10-2005 at 11:37

I'm sure your method works, I am just rasing cautionary issues for those who might want to use this method and havn't done the research you have.

In plain english: If the sulfuric is too strong, the heat is too high, and the ammonium nitrate is not very pure, the mixture can deflagrate or detonate.

Joeychemist - 26-10-2005 at 17:04

Although it is off topic, the attached file should help to clear up the H2S04+NH4NO3=Nitramine issue.

Attachment: dehydration of ammonium nitrate.pdf (221kB)
This file has been downloaded 999 times

Rosco Bodine - 28-10-2005 at 14:12

NH4HSO4 mol. wt. 115.11

[(NH4)2]S2O7 mol. wt. 212.22

(2) NH4HSO4 ------> [(NH4)2]S2O7 + HOH

230.22 grams of NH4HSO4 ( 2 moles ) was heated gradually to 400 C . The material first melted and then as the temperature rose decomposition set in as a gentle boiling with evolution of steam and also some acrid smoke with a choking odor . The boiling away gradually ceased
and a clear melt resulted at about 380 C .

A weight loss of 14.3 grams was measured for the dehydration , which
would have given a maximum theoretical loss of 18 grams ( 1 mole ) of water for
complete conversion to the pyrosulfate .
So it appears that the reaction is only easily heat driven to 79.4% completion ,
resulting in an equilibrium mixture with the residue of unconverted acid sulfate which is a melt solidifying at about 160 C .

The melt 79.4% ammonium pyrosulfate residue has not yet been tested for reactivity with sulfuric acid .

Rosco Bodine - 2-11-2005 at 18:34

An experiment was performed to observe the result of a distillation at normal pressure of HNO3 from a mixture of 92% H2SO4 and uncoated pure prills of NH4NO3 .

From 650 grams of NH4NO3 and 650 ml of 92% H2SO4 was distilled 300 ml of very pale yellow , light ale colored HNO3 which
had a measured density of 1.5003 ~96.8% HNO3 . The distillation proceeded smoothly at a delivery rate of 1 to 2 drops per second with very little supplemental heating required , with the vapor coming over beginning ar 86 C and
gradually increasing in temperature to 121 C at the time the heating was stopped , 300 ml of distillate representing
an 87% yield based upon NH4NO3.

-----material below edited in from old board-----

It being overdue that I put the arguments to rest about the distillation of 1.5 nitric from drain cleaner 92% SA , and NH4NO3 prills , I have confirmed this by experiment .

2 kilos of highly pure and uncoated NH4NO3 prills were dried at 80C and so little weight loss was observed that this was unnecessary preparation .

650 grams of the NH4NO3 prills were placed in a 2,000 ml RB 24/40 flask ,
which was equipped with a 12 X 40 mm rare earth stirbar , and heated by an Electromantle stirring mantle . The 3 way distillation adapter was equipped with
a top thermometer , and side discharged
into a Pope Scientific flexible teflon connecting tube to an angle adapter on the top of a 600 mm Graham condenser cooled by recirculated ice water . The discharge of the condenser was fitted with a side tubulation straight vacuum adapter having a vent line for remote discharge of any fumes from the closed system , and a 500 ml Erlenmeyer receiving flask was pinch clamped to
the discharge of the straight adapter .

650 ml of 92% drain cleaner SA was poured through a long stem funnel inserted through the vertical opening
in the distillation adapter , and the thermometer and 24/40 bushing adapter
was then inserted to close the apparatus . The stirrer was started in
auto-reversing mode with heating at 30%
to slowly bring up the heat and gradually dissolve the prills , and then the stirring was set to maximum in single direction mode . The heating was very gradual to allow for any offgassing which typically precedes such distillations with a
" false boil " , which requires a temporary cessation of all heating to manage . This preliminary offgassing and crud layer formation occurred as expected , and dissipated quickly after 10 minutes , whereafter heating was resumed at 16% to initiate the distillation and reduced back to 10% as the mixture began to foam . The reaction goes through a sort of induction period where it simply foams with a head of foam about
12 - 18 mm thick having the appearance of effervescence like the head on a mug of ale .....not a typical boiling , and the
vapor refluxes fully within the flask , not
reaching the thermometer . Dry terrycloth
toweling was wrapped around the upper part of the flask and the distillation head to insulate it against heat loss , and in a few minutes the thermometer began to indicate rising vapor , with the first distillate coming over at 86 C and the cooling pump was started . The distillation
proceeds very smoothly almost from its own heat of reaction with very little supplemental heating at 10 to 14 % on
the rate controller . The nitric acid comes over at a rate of 1 to 2 drops per second
and the vapor is somewhat superheated
in the reaction flask , as the reaction producing the nitric acid is occuring in a range of estimated 105 - 120 C , well above the boiling point of the nitric acid
evolving as vapor from the reaction mixture . The distillation proceeds with
a gradual increase of the temperature of
the distilling vapor , at 100 C early in the distillation , rising to 105 C @ 150 ml in the receiver , 108 C @ 170 ml , 112 C @ 200 ml , 115 C @ 225 ml , 119 C @ 280 ml , 121 C @ 300 ml .....whereupon the
heating was ended and the distillation stopped , accumulated product ~ 87% of
theory based upon NH4NO3 .

100 ml of the HNO3 was carefully weighed using a Class A volumetric flask which was carefully dried and tared in advance , using an Ohaus 311 quad beam scale . The density of the HNO3 was determined to be 1.5003 which is about 96.8% HNO3 . The color of the acid is about like a light ale , very pale yellow .

[Edited on 11-3-2005 by Polverone]

Fleaker - 3-11-2005 at 15:21

Excellent report! 87% is an acceptable yield (~10%+ than similar distillations I performed with tec. grade potassium nitrate and 93% sulfuric acid in 24/40 glass setup).

I do have several questions for you, Rosco. First, would not a 0 deg. C vacuum distillation be preferred (assuming one has an adequate pump and trap)?

Second, in a stoichiometric distillation several months ago (again using KNO3) I inadvertently left the 'lab' to do some errand only to return and see the distillation temperature at 150C (and rising). Needless to say, I quickly turned down the heating mantle. When I went to check the level of the receiving flask I noticed two products in two distinct layers. One product was apparently nitric acid (perhaps produced before the temperature increase) and was floating on top but the other product on the bottom was a bright green! I can only guess that some other nitrogen oxide was formed. I did separate the two using a separatory funnel. Nitric acid was clearly a product (oily, fuming, and having the odor of nitric) while the other product was even thicker and having an even more distinct smell than HNO3.

I'd be interested in your (or anyone's) suggestion of what this bizzare product is. I've done quite a few distillations of near azeotropic nitric acid but I have never seen anything like it. I can only speculate that the higher temperature may have lead to the formation of something more favorable than HNO3.

Thank you.

edit: adjective choice

[Edited on 5-11-2005 by Fleaker]

Rosco Bodine - 3-11-2005 at 16:31

Quote:

Originally posted by Fleaker
Excellent report! 87% is an acceptable yield (~10%+ than similar distillations I performed with tec. grade potassium nitrate and 93% sulfuric acid in 24/40 glass setup).

When the ratio of H2SO4 and H2O to NH4NO3 is optimum the yield is possible to be 98 to 99% based on the NH4NO3 , and I have in the past gotten that good yield on the azeotropic nitric , but I cannot locate the notation which I made at the time on the correct proportions .
Quantitative yield may also be possible for the d. 1.5 HNO3 , but I am not sure about the optimum proportions for it either , but I do believe a bit more of a slightly stronger H2SO4 would be of benefit there , and likely 96% H2SO4 would be fine . But really the yield of 87% is fine for most purposes . I think the probable best still size for a more worthwhile distillation run would be a 3000 ml or perhaps even a 5000 ml instead of the 2000 ml which I am using .
2000 ml is about the smallest size you would want to use for the distillation ,
unless you just need a thimble full of HNO3 product for all your labor . Anyway the KNO3 is a less desirable nitrate if distillation of HNO3 is intended , because of the insolubility and increased volume of the acid sulfate byproduct in comparison to NaNO3 or NH4NO3 which are better precursors .

Quote:

I do have several questions for you, Rosco. First, would not a 0 deg. C vacuum distillation be preferred (assuming one has an adequate pump and trap)?

No . In my estimation that is junk science and myth about HNO3 distillation . Sunlight and high heat and too fast a distillation are the things you need to avoid in terms of getting a pure product ,
not atmospheric pressure . HNO3 is sufficiently stable to be distilled without any great decomposition if it is done gently and slowly , and the heat needed to drive the reaction forward is pretty closely matched to the reaction condition at atmospheric pressure . If there was any benefit for vacuum distillation it would likely be at something in the range of one half to two thirds of 1 Atm of pressure ,
definitely not a hard vacuum . Now for redistillation of the HNO3 later , that may be a time when the higher vacuum would be of benefit ......but not from the original reaction mixture of H2SO4 and a nitrate .
You see the distillation is limited by the reaction producing the HNO3 which is being distilled , so the vacuum at this stage is not very helpful , and can even upset what would otherwise be a very well behaved reaction , with the inevitable pressure fluctuations causing oscillations of instability in the boiling of the mixture , introducing more trouble than it is worth . Stick with atmospheric
pressure , it is more predictable what the reaction mixture will do .

Quote:

Second, in a stoichiometric distillation several months ago (again using KNO3) I inadvertently left the 'lab' to do some errand only to return and see the distillation temperature at 150C (and rising). Needless to say, I quickly turned down the heating mantle. When I went to check the level of the receiving flask I noticed two products in two distinct layers. One product was apparently nitric acid (perhaps produced before the temperature increase) and was floating on top but the other product on the bottom was a bright green! I can only guess that some other nitrogen oxide was formed. I did separate the two using a separatory funnel. Nitric acid was clearly a product (thick, fuming, and having the odor of nitric) while the other product was even thicker and having an even more distinct smell than HNO3.

I'd be interested in your (or anyone's) suggestion of what this bizzare product is. I've done quite a few distillations of near azeotropic nitric acid but I have never seen anything like it. I can only speculate that the higher temperature may have lead to the formation of something more favorable than HNO3.

Thank you.

The reaction mixture probably foamed and carried over through the condenser and mixed with the distillate . You can pour the mixed crud in your receiver back into the reaction flask and stopper it while you rinse out the condenser and connecting tubes and restart the entire distillation again to clean up the product and salvage the run . Generally you
really have to watch a KNO3 reaction because of this potential complication of
the foaming of the insoluble bisulfate which forms a slurry of solids in the acid mixture .

This also limits the batch size since you have to allow for the expansion of the bubbling solids , a mud which considerably increases the volume of the reaction mixture and makes a boilover more likely than from a completely fluid reaction mixture . This is where the use
of NH4NO3 really provides advantage ,
the tradeoff being it is a much slower distillation .

Microtek - 4-11-2005 at 01:44

Good quality HNO3 of high concentration shouldn't be "thick" ( I'm assuming you use that word in the meaning "viscous" ). HNO3 of 95%+ is as mobile as water. Also, when measuring the density of the distilled HNO3, it is important to purge all the NOx from the acid, otherwise you can get unrealistically high values ( a clear, pale yellow acid with little or no water content can have a density of more than 1.52 g/cc due to the NOx ).
For removing the NOx, I use an aquarium pump to blow air through a vertical 40 cm tube filled with CaCl2 flakes ( with cotton wool at both ends to catch any dust ). The air is then led through a bubbler made from PE or teflon into the HNO3. I heat the nitric to 50 C and maintain the bubbling until the acid is as clear as pure water ( about 5 minutes depending on batch size and air flow ).

Rosco Bodine - 4-11-2005 at 05:46

The effect of daylight on the decomposition of highly concentrated HNO3 is significant . I have watched
water clear HNO3 accumulating in the receiver and before 100 ml has accumulated , a slight yellow tint begins to appear simply because of the photosensitivity . Even in a shaded area
where there is indirect daylight , a clear sample of 97% HNO3 will take on a pale yellow tint in 30 minutes , so the stuff is
definitely photosensitive . And if exposed to direct sunlight , red fumes will appear in seconds above the liquid and it will gradually darken through the yellow range all the way to red . I have found it is essential to store the highly concentrated HNO3 in the dark and very cold , preferably in a glass bottle having an attached teflon seal and kept in a freezer at -10 C .

The very pale yellow tint is practically an unavoidable decomposition trace which can be most easily eliminated with a pinch of urea or urea nitrate . I have not encountered any problems in syntheses with the pale yellow acid , so I would suppose it depends on how demanding is your requirement for absolute purity for a particular reaction at the very start , whether the acid should be absolutely colorless . I have doubts that this level of perfection for colorless HNO3 is required for most syntheses .

The concentrated HNO3 is quite a good solvent for silicone high vacuum grease amd will wash it from a greased joint as easily as will toluene wash off vaseline , something which is good to keep in mind with regards to your glassware . The joints should not be left assembled for an extended time after exposure to concentrated HNO3 or its vapors , or they are likely to become frozen .

Joints which have become frozen
can be separated by applying a constant parting tension of about 4 kilograms using a suspended weight , and hitting the opposite sides of the outer joint simultaneously with a medium high blast of flame from two propane torches , so that the outer joint is heated very suddenly and intensely . The surface area below should be padded to prevent the breakage of the extracted part when it pops free under the tension .

Axt - 4-11-2005 at 09:00

Talking about nitric distillation in the sun...

<center><img src="http://www.sciencemadness.org/scipics/axt/solar-distiller.jpg"></center>

......sorry

(it didn't work at all)

Rosco Bodine - 5-11-2005 at 08:17

Hmmmm , yeah I think know what that
apparatus is .....recognized it right off .

It's one of those gadgets that an academic on horseback holds by one end and swings round and round like a lariat ,
and then throws it to entangle the feet
of fleeing bush chemists

update

Rosco Bodine - 26-10-2006 at 11:38

The preliminary experiment with the ammonium bisulfate
as having potential usefulness as an intermediate for
the production of oleum , has not yet been further explored .

So whether this line of experimentation for a simplified route to oleum will be fruitful or not remains to be seen .

Today while doing a bit of googling concerning sulfamic acid
I ran across this information which may be pertinent concerning how urea may be a useful component in such
a scheme .

The article linked below shows the reactions for urea and sulfuric acid which leads to sulfamic acid and ammonium bisulfate intermediate products , and further heating to decomposition producing several things including SO3 .

http://en.wikipedia.org/wiki/Sulfamic_acid

I have not yet plotted the course of these urea / sulfamic acid decomposition related reactions and how they might interact with the ammonium pyrosulfate ....but at first glance
it appears that urea could facilitate the process by virtue of
the formation of sulfamic acid as an intermediate and its
subsequent decomposition on further heating .

What is being sought is a scheme for production of oleum under milder conditions than some of the other known methods .

This scheme also may also have usefulness as a method
for acetic anhydride from ammonium acetate , via a regenerable dehydration reagent , as mentioned by
Dave Angel above in the first reply in this thread .

[Edited on 26-10-2006 by Rosco Bodine]

Rosco Bodine - 19-11-2006 at 08:10

Very strange . Immediately after I linked above the page on
wikipedia where I saw this interesting role of urea in the
formation of sulfamic acid , and noted the possible usefulness
of this in schemes for production of SO3 at higher temperatures where decomposition would occur ..........

the wiki article was * edited * and the specific lines were
removed from the current page .......very strange .

The following lines were the ones that had caught my notice ,
and were edited from the current page .

This makes me even more curious about the reaction than
I was before

portion of article removed :

Synthesis

NH3 + SO3 → H3NSO3

NH2OH + SO2 → H3NSO3

According to Greenwood and Earnshaw, the industrial synthesis uses urea as the source of NH3:

(H2N)2CO + 2 H2SO4 → CO2 + H3NSO3 + NH4HSO4

( version with subscripts below )

==Synthesis==

NH3 + SO3 → H3NSO3
NH2OH + SO2 → H3NSO3

According to Greenwood and Earnshaw, the industrial synthesis uses urea as the source of NH3:
(H2N)2CO + 2 H2SO4 → CO2 + H3NSO3 + NH4HSO4

Aqua_Fortis_100% - 7-2-2008 at 06:58

Sorry by bringing up this topic again..

But so simply urea and sulfuric acid reacted together will give sulfamic acid (and ammonium bissulfate)? looks good for me, since sulfamic acid here come only pure in large , non - comfortable amounts..
Anyway, how to isolate this from the mix?

Also, can anyone also say if method proposed by BASF in Ac2O thread is reliable to produce ammonium pyrosulfate?

Patent attached

[Edited on 7-2-2008 by Aqua_Fortis_100%]

Attachment: US3885024.pdf (86kB)
This file has been downloaded 916 times

LSD25 - 20-5-2008 at 22:12

I just finished reading this article (attached) and thought it might be best placed here.

Apparently the authors used a propane flame (oxygen defficient) to burn both ammonium sulfate and ammonium bisulfate to give predominantly SO2 with some SO3 (<5%). Of rather more interest is the design of the burner - they used a steel bell (well sounds like it) and a propane burner within that bell to generate the gasses from a solution of the salts (which was sprayed into the flame). The principal design was to cut down on the amount of SO3 and NOx produced. Seems like if they had to design a process to cut down on these, it likely wouldn't be too hard to reverse engineer the process to generate significant quantities of them (although it is hard to see how to contain the SO3 that might be produced, perhaps PTFE lining on the interior of the burner vessel? Well above the flame? What about glass - what temperature would the gas be at? Perhaps if the pyrolysis is happening in the actual flame-cone, then the gas temperature could be sufficiently low as to allow for less durable containers?)

[Edited on 21-5-2008 by LSD25]

Attachment: Evaluation.Combustion.Process.for.Feedstock.AmmoniumSulfate.AmmoniumBisulfate.pdf (201kB)
This file has been downloaded 1168 times

12AX7 - 20-5-2008 at 23:25

Propane, C3H8, is rich in hydrogen, an element which burns to steam, H2O. SO3 isolation will be problematic at best, but the SO2 could be seperated and ran through other processes.

If all the gasses were collected and compressed, you'd get quite a bit of acidic water (H2SO3 / H2SO4), whatever SO2 was produced, and CO2, either in solution or gaseous. Probably moderate fractionation would be necessary to seperate them...

Tim

LSD25 - 21-5-2008 at 20:40

I forgot to attach the article I cited, I believe it dealt with some of the points you raised regarding the potential output

[Edited on 21-5-2008 by LSD25]

Attachment: Evaluation.Combustion.Process.for.Feedstock.AmmoniumSulfate.AmmoniumBisulfate.pdf (201kB)
This file has been downloaded 667 times