AndersHoveland
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Catalytic decomposition of Ammonium Nitrate using Nitric Oxide
On a hunch, I decided to investigate whether nitric oxide can catalyze the decomposition of ammonium nitrate at room temperature. 
Hypothesis
NH4NO3 + H2O <==> NH4OH + HNO3
(2)HNO3 + NO <==> H2O + (3)NO2
I have read of an experiment where bubbling in nitrogen dioxide can (very slowly) oxidize solutions of ammonium perchlorate to free perchloric acid.
I would also like to mention that a source I found in the literature claims that nitric oxide alone has no effect on ammonium perchlorate solutions,
so possible oxidation of ammonium ions directly by nitric oxide should not be a concern.
Experiment
20 grams of dry technical grade NH4NO3 was weighed out and dissolved in 100mL distilled water. I will emphasise that the prills of regent appeared
completely dry, and that they were immediately weighed after removal from the air tight container.
Nitric oxide was prepared according to the following procedure:
by reaction of sodium nitrite with dilute hydrochloric acid, with ferrous sulfate, FeSO4, also present to react with, and prevent the escape of, the
nitrogen dioxide. The reactions are:
(2)NaNO2 + (2)H2SO4 + (2)FeSO4 --> Fe2(SO4)3 + Na2SO4 + (2)H2O + NO
The initial temperature in the NH4NO3 solution was measured at 17°C.
The nitric oxide was passed into a stoppered flask of the NH4NO3 solution. Care was taken to exclude air from the flasks. The flasks were flushed with
carbon dioxide immediately prior to being connected.
The nitric oxide was slowly bubbled into the solution for 60 minutes. More regents were added into the nitric oxide-generating flask intermittently as
required to maintain the gas flow over the time of the procedure.
After 60 minutes, the flask was disconnected, with the stopper still on. The air space above flask containing the NH4NO3 was flushed out carbon
dioxide, then quickly connected to another flask containing a plain water solution with a little sodium sulfite dissolved, such that air would not be
able to enter the first flask. Next, the flask containing the NH4NO3 solution was boiled down on a hot plate to nearly a quarter of its initial
volume, the vaporized water venting off into the second flask. The NH4NO3 solution, still hot, was allowed to partially cool, then unstoppered, thus
now exposing it to air after all the nitric oxide had been expelled. The flask was refitted for warm (80-90°C) distillation under reduced pressure.
Care has taken to ensure the flask was only gently heated, and the hot plate temperature was at no time allowed to reach 100°C. Every few minutes a
drop of water was placed on the hot plate to ensure it would not instantaneously boil away. Crystals of NH4NO3 began to form in the solution, and
after approximately 20 minutes, the NH4NO3 was obtained in apparent anhydrous form. Just prior to the complete disappearance of all the liquid, the
heat setting was set down to a lower setting, and the last remaining quantities of water where removed. Care was taken to avoid any possible thermal
decomposition of the NH4NO3.
The original flask weighed 49.022 grams.
After addition of the NH4NO3 it weighed 69.100 grams.
After the reaction, the flask containing the NH4NO3 weighed 67.432 grams.
Conclusion
I made the best attempts possible to exclude all air from reacting with the nitric oxide. It seems that 1.668 grams of the original 20.078 grams of
ammonium nitrate was decomposed over the course of the reaction.
This experiment had several weaknesses, namely that some of the nitric oxide likely spontaneously disproportionated into nitrous oxide and nitrogen
dioxide, which would obscure the results. A better experiment would have been to measure the volume of the entering and escaping gases to determine if
more gas was emitted than was introduced. And in retrospect, I probably should have dried the NH4NO3 under vacuum at the beginning of the experiment,
just to be certain it was completely anhydrous.
[Edited on 22-12-2011 by AndersHoveland]
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bbartlog
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Nice to see some experimental work from you! The decomposition seems quite plausible, since gassing the solution as you describe should create
ammonium nitrite (even if only transiently), which is known to be unstable under acid conditions.
The less you bet, the more you lose when you win.
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497
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I think the rate of disproportionation at atmospheric pressure is negligible. http://www.ncbi.nlm.nih.gov/pubmed/12653205
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AJKOER
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AndersHoveland:
A wild idea. My recollection from a dated source was that ammonia will react (perhaps even made to burn?) in the presence of various oxides of
nitrogen (NO, for example). So if one could safely extract the NH3 from NH4NO3 (acidify or heat?) and initiate a reaction with NO (spark/heat), a
significant reaction my be evident (with considerable precautions as one may simply initiate an explosion).
As a sidebar, came across good reference on making Chloride free HOCl that I believe you express an interest. Essentially, make an aqueous version of
HOCl (say CO2 acting on moist Ca(ClO)2) and strip out the Cl2O/HOCl with a polar solvent. Possible solvent mentioned include ketones, nitriles and
esters (see page 557). Interestingly CCl4 extracts Cl2O, but not HOCl (per page 546).
SOURCE:
DICHLORINE MONOXIDE, HOCl, HYPOCHLORITES Vol. 8, page
LINK:
http://www.questscan.com/?tmp=redir_bho_bing&prt=Qstscan...
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AndersHoveland
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Ammonium nitrate is tends to be very stable, and relatively unreactive apart from its mild acidity, at ambient temperatures. It seems to be quite
remarkable that it can be induced to decompose at room temperature.
Nitric oxide has a remarkably long lifespan considering that it is actually a reactive radical. Ammonium ions are not easily oxidized, and nitrate
ions tend to be relatively inert (except at high temperatures or under extremely acidic conditions). The fact that NO is a reactive radical seems to
be able to effectively catalyze the oxidation of the ammonium ions by the nitrate ions, in a two-step process.
I am not quite sure what the decomposition products from ammonium nitrate are, when catalysed by this method. It is quite possible that one molecule
of NO is only able to catalyse a finite number of NH4NO3 pairs until it is itself consumed, which is to say that the NO is regenerated in the reaction
at a lower rate than it reacts.
Also to mention that NO can only reduce acidified nitrate ions. Under more alkaline conditions, the reaction is actually the opposite, which
is to say that NO2 oxidizes nitrite to nitrate, leaving NO.
[Edited on 24-12-2011 by AndersHoveland]
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AJKOER
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I think it would be interesting to see if there is any difference upon working with the dry decomposing NH4NO3 in an atmosphere of NOx than the
aqueous reaction (both in yield and path).
As an example of different results, take the case of Ca(ClO)2 acted upon Chlorine. Dry yields Cl2O and CaCl2, while aqueous also forms Ca(ClO3)2.
Reference see page 561 at:
www.scribd.com/doc/30121142
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AndersHoveland
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The reaction rate for this experiment was very slow. I would think the reaction rate between solid NH4NO3 and gaseous NO, in the absence of water,
would be negligible.
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AJKOER
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Of course you are right except that NOx can (and has been) catalyzed to react with NH3. See, for example, Wikipedia's Ammonia remediation comment:
"For remediation of gaseous emissions
Ammonia is used to scrub SO2 from the burning of fossil fuels, and the resulting product is converted to ammonium sulfate for use as fertilizer.
Ammonia neutralizes the nitrogen oxides (NOx) pollutants emitted by diesel engines. This technology, called SCR (selective catalytic reduction),
relies on a vanadia-based catalyst.[38]" (where [38] links to an article).
I will let you ponder if a SCR is at all workable here.
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plante1999
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I remember seeing a post a while back where some user showed a video where ammonium nitrate analytically decomposed and asked what was the catalyst.
After some test and research, I found that a mixture of potassium ferrocyanide and potassium dichromate will catalyze ammonium dichromate
decomposition. A pinch of the two salts mixed with 10 g of ammonium nitrate make a suitable mixture. The ammonium nitrate is fairly easy to ligth and
will burn ferously leaving a green glass. It will make a lot of fumes that readily disperse, maybe because they are mostly water vapors.
A very impressive demonstration. If mixed with carbon, as in the original ammonpulver black powder, it may be possible to attain much faster burning
rate then what we observe when testing the original 85 ammonium nitrate 15 carbon ratio.
I never asked for this.
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papaya
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Quote: Originally posted by plante1999  | I remember seeing a post a while back where some user showed a video where ammonium nitrate analytically decomposed and asked what was the catalyst.
After some test and research, I found that a mixture of potassium ferrocyanide and potassium dichromate will catalyze ammonium dichromate
decomposition. A pinch of the two salts mixed with 10 g of ammonium nitrate make a suitable mixture. The ammonium nitrate is fairly easy to ligth and
will burn ferously leaving a green glass. It will make a lot of fumes that readily disperse, maybe because they are mostly water vapors.
A very impressive demonstration. If mixed with carbon, as in the original ammonpulver black powder, it may be possible to attain much faster burning
rate then what we observe when testing the original 85 ammonium nitrate 15 carbon ratio. |
Must be my thread: https://www.sciencemadness.org/whisper/viewthread.php?tid=24...
You wrote "a mixture of potassium ferrocyanide and potassium dichromate will catalyze ammonium dichromate decomposition", did you mean - "ammonium
nitrate" instead of ".. dichromate" ? If it's about AN decomposition, please give details like percentages,etc., in particular do you need both salts
together or just one is enough, is ammonium dichromate a possible substitute, can potassium ferricyanide be used instead of ferro one, do you need to
mix them in solution or dry powders will also work.. Because I'm still interested in that topic. BTW, ferro/ferricyanides cannot be counted as real
catalysts, because they are good fuels.
EDIT: It will be better if you answer to this in the related thread, as this discussion is off-topic here. Thanks!
[Edited on 9-8-2013 by papaya]
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