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Ral123
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[*] posted on 27-3-2013 at 06:15


For a long time I thought the easiest way to make nitrogen gas at home is NaNO2+NH4Cl. I added the two solutions in a small flask and with a little warming, bubbles formed. Off course they didn't sustain fire. Ammonium nitrite seemed nice oxygen balanced explosive to me some time ago. Now I can't imagine it stable enough for whatever job.
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[*] posted on 27-3-2013 at 07:32


Quote: Originally posted by Ral123  
.... Now I can't imagine it stable enough for whatever job.


Actually, NH4NO2 is not practical explosive on several levels. First, it has a half life of around 2 hours (see http://books.google.com/books?id=oAlxIVRzrN8C&pg=PA560&a... ), which translates if you did create 8 grams (not recommended), at the end of 2 hours you have 4 grams left, in 4 hours, 2 grams and at the end of 6 hours a single gram. Next, it behaves badly (shock and heat sensitive even if in aqueous solutions). And finally, it is not just considered an explosive, but a high explosive.
------------------------------------------------------

Also, those forming Ammonium nitrite via NH3 and H2O2 in the presence of a lot of NaOH (or Na2CO3, which are cited catalysts) may also, by NaNO2 formation, be able to avoid the decomposition issue associated with NH4NO2. A source cites the reaction (see http://books.google.com/books?id=6OcDAAAAMBAJ&pg=PA277&a... ):

NH4NO2 + NaOH --> NaNO2 + H2O + NH3 (g)

It is also conceivable in a highly alkaline condition that actually mostly NaNO2, and not half NaNO3, is created. My source: "Mechanism of the NO2 conversion to NO2- in an alkaline solution" by Chen X, Okitsu K, Takenaka N, Bandow H. at the Department of Applied Materials Science, Graduate School of Engineering, Osaka Prefecture University, to quote from the abstract:

"The reaction of NO2 and NaOH aqueous solution at room temperature was studied for elucidating the behavior of gaseous NO2 in an alkaline solution. Experimental runs related to NO2 absorption have been carried out in various pH solutions. The nitrite and nitrate ions formed in these absorption solutions were quantitatively analyzed. In the case of pH 5-12, both of the nitrite and nitrate ions were formed simultaneously. On the other hand, only the nitrite ion was formed when the pH of the absorption solution was higher than 13. In this paper, a new reaction mechanism was proposed to explain the selective formation of nitrite ion in the 10 M alkaline solution. In order to confirm the new reaction mechanism, H2(18)O was used as part of the absorption solution for detecting oxygen gas production. The amounts of reaction products: (18)O(18)O, (18)O(16)O and (16)O(16)O, were quantitatively determined. It was confirmed that the new reaction proceeds mainly in the 10 M alkaline solution."

Link: http://www.ncbi.nlm.nih.gov/pubmed/15636532

Now, those targeting NaNO3 and not NaNO2 could try adding add NaOCl or HOCl as:

NaNO2 + HOCl --> NaNO3 + HCl

However, it is important first to let a warm solution vent out any excess NH3 still presence (and unreacted) due to avoid the possible formation of any chloramines (so perform outdoors in any event). The latter are toxic and NCl3 (a yellow oily liquid that only slowly undergoes hydrolysis) is an even more ill tempered explosive than NH4NO2 whose formation is dependent on an excess HOCl concentration and low pH conditions (note the formation of HCl with the reaction of HOCl on NaNO2). Some reaction sequences:

NH3 + HOCl <--> NH2Cl + H2O
NH2Cl + HOCl --> NHCl2 + H2O
2NH2Cl <--> NHCl2 + NH3
NHCl2 + HOCl --> NCl3 + H2O

to cite a few of the reactions (for more details, see page 16 at http://books.google.com/books?id=zAND55GZet4C&pg=PA16&am... ). Note, this reaction sequence implies that in acidic environments, the hydrolysis of NH2Cl may be preferred forming more HOCl to react with any existing NHCl2 to form some NCl3. So even without intentional excess hypochlorous acid, low pH could foster some NCl3 formation.

[EDIT] On second thought, I now have stronger reservations on the use of NaOCl or HOCl in the possible presence of NH3. The interaction of any formed NH2Cl and NaNO2 may proceed as follows (speculation):

NaNO2 + NH2Cl --?--> NH2NO2 + NaCl

where the possible formation of nitramide, or a voilent decomposition, is viewed as problematic. If this reaction is successful, replacing NaNO2 with AgNO2 to produce an insoluble AgCl may be a preferrable procedure. [EDIT][EDIT] See Franklyn's prior (from 2007) comments/research on this very idea titled "nitramide from chloramine and nitrite" and also the response by AndersHoveland at http://www.sciencemadness.org/talk/viewthread.php?tid=6042 . Also, a source (Canis, C., Rev. Chim. Minerale, 1964, 1, 521) notes that "Nitramide is quite unstable and various reactions in which it is formed are violent. Attempts to prepare it by interaction of various nitrates and sulfamates showed
that the reactions became explosive at specific temperatures." Also with alkalies (source: Thiele, J. et al., Ber., 1894, 27, 1909) "A drop of conc. alkali solution added to solid nitramide causes a flame and explosive decomposition" and also "Nitramide decomposes explosively on contact with conc. sulfuric acid." Caution: Chloramine is known to attenuated the posionous properties of other compounds (for example, with CH3OH), so exercise appropriate safety in this chloramine/nitrite combination.

In place of NaOCl and HOCl, one may be abe to employ H2O2, but be aware that hydrogen peroxide in the presence of select organics and certain metal oxides/halides (metals include Al, Ti, Ce,..) may form an oxynitrite, and not the nitrate.


[Edited on 28-3-2013 by AJKOER]

[Edited on 28-3-2013 by AJKOER]
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[*] posted on 27-3-2013 at 10:27


This evening I went looking at my NH4NO2 solution and see what remains of this. All water now has evaporated and there is a very small amount of solid left. The funny thing is, the solid forms a very thin layer and this layer has the shape of bubbles. I added a small amount of water to the solid and the solid does not dissolve quickly. Next, I added a small amount of dilute H2SO4 and when this is done, no visible reaction occurs. The air above the solid becomes very pale brown, it is just visible, and there is the smell of NO2, unmistakenly, but weak.

So, the conclusion of my experiment is that making solid NH4NO2 from a solution of NH4NO2 in water is not possible. The material decomposes on evaporation of the water. The remaining solid must be KClO4 and/or NH4ClO4, remaining from the perchlorate added. Traces of nitrite were still left in the solid, but not more than traces.

So, in this case the books are right. NH4NO2 cannot be prepared from aqueous solution, it decomposes to water and N2 when the solution is allowed to evaporate.




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[*] posted on 27-3-2013 at 13:54


Quote: Originally posted by Antiswat  
today i cleaned a 0.5 L PVC bottle with NO2 from Cu + HNO3..
i decided to open a bottle of 25% NH4OH aswell
NO2 + NH4OH > NH4NO2
or am i wrong?
the brown gas went very fast into white smoke alike other acid reactions in air with ammonia..
but... i have searched around on here to see if anybody have been wanting to make ammonium nitrite.. apparently not, seemingly its ''acute toxic'' that doesnt sound very pleasant..
anyways my thought about NH4NO2 is that it should be possible to make quite easily by NO2 + NH3 or even NH4OH
it could then be used to make other nitrites as i also read that it very quickly goes into H2O and N2...

NH4OH + NO2 > NH4NO2
NH4NO2 + Na2CO3 > (NH4)2CO3 + NaNO2
i really love ammonium carbonate's property to decompose in hot water..
what do you guys think.. useful or trash? the Na2CO3 could be dissolved in weak NH4OH btw


Here is the cited preparation per Wikipedia (http://en.wikipedia.org/wiki/Ammonium_nitrite ):

"Ammonium nitrite forms naturally in the air and can be prepared by the absorption of equal parts nitrogen dioxide and nitric oxide upon aqueous ammonia.[2]

It can also be prepared by oxidizing ammonia with ozone or hydrogen peroxide, or in a precipitation reaction of barium or lead nitrite with ammonium sulfate, or silver nitrite with ammonium chloride, or ammonium perchlorate with potassium nitrite. The precipitate is filtered off and the solution concentrated. It forms colorless crystals which are soluble in water and decompose on heating or in the presence of acid, with the formation of nitrogen.[3] Ammonium nitrite solution is stable at higher pH and lower temperature. If there is any decrease in pH lower than 7.0, It may lead to explosion. It is desirable to maintain pH by adding ammonia solution. The mole ratio of Ammonium Nitrite to Ammonia must be above 10% mole ratio.

NH4NO2 → N2 + 2 H2O "

Wiki also notes that "Ammonium nitrite may explode at a temperature of 60–70 °C.[2] It decomposes more quickly when a concentrated solution than when it is a dry crystal."
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[*] posted on 28-3-2013 at 04:24


Quote: Originally posted by woelen  
This evening I went looking at my NH4NO2 solution and see what remains of this. All water now has evaporated and there is a very small amount of solid left. The funny thing is, the solid forms a very thin layer and this layer has the shape of bubbles. I added a small amount of water to the solid and the solid does not dissolve quickly. Next, I added a small amount of dilute H2SO4 and when this is done, no visible reaction occurs. The air above the solid becomes very pale brown, it is just visible, and there is the smell of NO2, unmistakenly, but weak.

So, the conclusion of my experiment is that making solid NH4NO2 from a solution of NH4NO2 in water is not possible. The material decomposes on evaporation of the water. The remaining solid must be KClO4 and/or NH4ClO4, remaining from the perchlorate added. Traces of nitrite were still left in the solid, but not more than traces.

So, in this case the books are right. NH4NO2 cannot be prepared from aqueous solution, it decomposes to water and N2 when the solution is allowed to evaporate.


yes, that seems to confirm its unstability, but as ajkroer stated ammonia makes it more stable, so possibly bubbling NO2 through ammonia would be able to keep it stable for abit, but also HNO3 would form with the water and from that NH4NO3




~25 drops = 1mL @dH2O viscocity - STP
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https://en.wikipedia.org/wiki/Solubility_table
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[*] posted on 28-3-2013 at 04:28


Quote: Originally posted by AJKOER  
Quote: Originally posted by Antiswat  
today i cleaned a 0.5 L PVC bottle with NO2 from Cu + HNO3..
i decided to open a bottle of 25% NH4OH aswell
NO2 + NH4OH > NH4NO2
or am i wrong?
the brown gas went very fast into white smoke alike other acid reactions in air with ammonia..
but... i have searched around on here to see if anybody have been wanting to make ammonium nitrite.. apparently not, seemingly its ''acute toxic'' that doesnt sound very pleasant..
anyways my thought about NH4NO2 is that it should be possible to make quite easily by NO2 + NH3 or even NH4OH
it could then be used to make other nitrites as i also read that it very quickly goes into H2O and N2...

NH4OH + NO2 > NH4NO2
NH4NO2 + Na2CO3 > (NH4)2CO3 + NaNO2
i really love ammonium carbonate's property to decompose in hot water..
what do you guys think.. useful or trash? the Na2CO3 could be dissolved in weak NH4OH btw


Here is the cited preparation per Wikipedia (http://en.wikipedia.org/wiki/Ammonium_nitrite ):

"Ammonium nitrite forms naturally in the air and can be prepared by the absorption of equal parts nitrogen dioxide and nitric oxide upon aqueous ammonia.[2]

It can also be prepared by oxidizing ammonia with ozone or hydrogen peroxide, or in a precipitation reaction of barium or lead nitrite with ammonium sulfate, or silver nitrite with ammonium chloride, or ammonium perchlorate with potassium nitrite. The precipitate is filtered off and the solution concentrated. It forms colorless crystals which are soluble in water and decompose on heating or in the presence of acid, with the formation of nitrogen.[3] Ammonium nitrite solution is stable at higher pH and lower temperature. If there is any decrease in pH lower than 7.0, It may lead to explosion. It is desirable to maintain pH by adding ammonia solution. The mole ratio of Ammonium Nitrite to Ammonia must be above 10% mole ratio.

NH4NO2 → N2 + 2 H2O "

Wiki also notes that "Ammonium nitrite may explode at a temperature of 60–70 °C.[2] It decomposes more quickly when a concentrated solution than when it is a dry crystal."


yes.. it seems that its not very well preferred to have it as a solid and its perhaps even impossible.. if it would be possible it would probably take some lower temperatures and vacuum
but still i have an idea that NH4OH + Na2CO3 added in, would make NaNO2 as the NH4NO2 would react instantly with the Na2CO3
possibly not a great purity, but.. it would be there..

if this was made in huge amounts the NaNO2 could probably be purified by recrystallization




~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
http://www.trimen.pl/witek/calculators/stezenia.html
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[*] posted on 29-3-2013 at 05:26


OK, using some NaOH (or Na2CO3) as a catalyst is probably a good idea based on this old (1905) report (see page 242 at http://books.google.com/books?pg=PA242&lpg=PA242&dq=... from Journal Chemical Society, London, Volume 88, Part 2), to quote:

"Electrolytic Oxidation of Ammonia to Nitrites. Erich Muller and Fritz Spitzer (Ber., 1905, 38, 778—782. Compare Traube and Biltz, Abstr., 1904, ii, 727).—In the presence of a small amount of sodium hydroxide, ammonia may be oxidised electrolytically to nitrite even in the absence of copper compounds.

In the presence of copper hydroxide and sufficient alkali, the oxidation of ammonia to nitrite does not cease suddenly when the nitrite concentration has reached a certain value, but appears to proceed quite independently of the nitrite concentration. In these experiments, the oxidation was allowed to proceed for a comparatively short time only, so that the amount of alkali present was not greatly reduced. The formation of nitrite is intimately connected with the amount of alkali present, and when no sodium 'hydroxide is present, but only ammonia, nitrite, and copper hydroxide, it is found that the nitrite is transformed into nitrate more rapidly than the ammonia into nitrite, and thus the concentration of the nitrite tends to decrease.

Nitrogen is also formed during the oxidation. J. J. S."

Apparently replacing NaOH with Cu(OH)2 favors the formation of nitrate (caution: could include copper ammonium nitrate [EDIT] may be right on this speculation, see http://www.pyrosociety.org.uk/forum/topic/3303-electrolysis-... ) over nitrites.


[Edited on 30-3-2013 by AJKOER]
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[*] posted on 30-3-2013 at 05:12


OK, I can across an incidental reference (source, see http://pubs.acs.org/doi/abs/10.1021/ed019p230 ) that appears to confirm one of Wikipedia's cited preparations of NH4NO2 to quote:

"can be prepared by the absorption of equal parts nitrogen dioxide and nitric oxide upon aqueous ammonia"

but, more interestingly, a dry cloud of ammonium nitrite (reputedly more chemically stable) seems to be formed when NO and NO2 comes into contact with gaseous ammonia in the presence of water vapor. This parallels the reaction of NH3 and HCl (forming the characteristic NH4Cl white smoke), which apparently requires moisture as well.

A small scale suggested procedure could be to flow moist NO and O2 (or air) into a chamber containing a bottle of open aqueous ammonia. The chamber walls should collect some of the NH4NO2 powder. One could also elect to place a NO generator (say dilute HNO3 and Cu) into the vessel and allow air contact.

Obvious advantages of this procedure is its simplicity with the formation of dry NH4NO2 of relatively high purity.

Disadvantage: To quote Axt (see http://www.sciencemadness.org/talk/viewthread.php?tid=11958 ): "NH4NO2 and NH4BrO3 seem to be of no practical value as high explosives due to their low chem. stability and high sensitivity to mech. and thermal shock." and also on NH4NO2 its "shock sensitivity is 6.5 times that of cyclotrimethylenetrinitramine". So separation of the powdered NH4NO2 from the walls of the vessel may be an issue, especially in the presence of organic matter (from a brush, etc.).
------------------------------------------------------------------------

I now believe per this reference on NaNO2 (see http://www.guidechem.com/reference/dic-15333.html ), to quote:

"Heating a mixture of an ammonium salt with a nitrite salt causes a violent explosion on melting, owing to formation and decomposition of ammonium nitrite. Salts of other nitrogenous bases behave similarly. Mixtures of ammonium chloride and sodium nitrite are used as commercial explosives."

so the aqueous solution of NH4NO2, upon addition of NaOH or Na2CO3, benefits from a elevated pH, but as long as NH4+ and NO2- are present, there is risk of explosion if the pH is not monitored. Also, any dry mix salt obtained should be considered as an explosive.


[Edited on 30-3-2013 by AJKOER]
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[*] posted on 30-3-2013 at 17:22


hm.. it seems weird that NaNO2 and NH4Cl can be used as a COMMERCIAL EXPLOSIVE..?
never heard about it.. cant imagine what use it could be if it needs to be melted first, but something indicates that its not really that good to heat it up alot, apparently..
unless if you really want it to go bang ofcourse..

but if im not wrong, RDX's sensitivity isnt really .. well existing so to say.. have anybody ever succeeded detonating it by shock? (:
i think this should be considered also.. collecting it from the walls of the vessel shouldnt be a problem if the sensitivity is 6.5 times as high as RDX's talking about shock, but its potentially explosive, interesting interesting..
think i might make some NaNO2 by NaNO3 + C and then attempt to heat a tiny amount with ammonium chloride..

but yes using NaOH makes more sense, as Na2CO3 could potentially react with the ammonia forming ammonium carbonate, not sure tho.

i read abit on a thread that suddenly showed up about electrolysis of ammonia with h2o2 and so on, turns out that it forms ammonium nitrite IIRC.. they state in one of your links aswell that nitrogen is formed during the process (instant decomposition, on small scale?)

but another link here in the thread lead to an indian scientist of some sort, that had concluded that ammonium nitrite wasnt really that unstable as thought of before

i think what causes the NH4NO2 to go bang when heated (detonation we cant really conclude without evidence for it) is because it acts abit like NH4NO3, just abit less stable, kind of like KClO3 compared to KClO4, i guess.
meaning it decomposes more violently at lower temperatures and is more unstable generally..
but just leading some air into NH3 atmosphere with some NOx forming.. could be done..




~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
http://www.trimen.pl/witek/calculators/stezenia.html
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[*] posted on 31-3-2013 at 06:30


My take is that heating dry NH4Cl simply forms vapors of NH3 and HCl. Then, HCl on NaNO2 forms HNO2 (or, more precisely H2O+NO+NO2) and the latter vapors form NH4NO2 on contact with NH3, as I have indicated previously. In essence, NH4Cl/NaNO2 removes some of the stability (or practicality) issues of ammonium nitrite itself although I would have qualms on the direct contact between NH4Cl and NaNO2 in the possible presence of moisture. Also, I would doubt if NH4Cl/NaNO2 was mixed with an organic or sulfide, or at least, not NaNO2 itself as the prior referenced MSDS on NaNO2 noted an immediate explosion hazard.
-------------------------------------------------------------------

A related idea for an experiment, react NH4Cl and H2O2 in the presence of a little NaOH. In the form of aqueous dilute solutions, I would expect a decomposition reaction forming N2 and leaving HCl and a little NaCl. For concentrated NH4Cl and H2O2 with a touch of NaOH in a sealed vessel, I would not be surprised on the possibility of an explosion due to the formation/decomposition of NH4NO2.

Other paths on the reaction of heated aqueous NH4Cl and H2O2 are possible. One could parallel the action of HOCl on NH4Cl, with the formation of chloramine, NH2Cl. But, the action of NH2Cl vapors on dry NaNO2 may be explosive as well, with the formation (or failed formation) of nitroamine, NH2NO2.


[Edited on 31-3-2013 by AJKOER]

[Edited on 31-3-2013 by AJKOER]
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[*] posted on 31-3-2013 at 18:49


Quote: Originally posted by Antiswat  

i decided to open a bottle of 25% NH4OH as well
NO2 + NH4OH > NH4NO2
or am i wrong?
the brown gas went very fast into white smoke alike other acid reactions in air with ammonia..
i also read that it very quickly goes into H2O and N2...


Actually, the gas phase reaction is:

(2)NO2 + (2)NO + (4)NH3 --> (2)NH4NO2 + (2)H2O + (2)N2

Dry ammonia gas reacts with the nitrogen dioxide and nitric oxide, at room temperature. (reaction investigated by Marcellin Berthelot)

Solid ammonium nitrite inside a tube explodes if heated on a water bath to between 60-70 °C. And the substance gradually decomposes at room temperature, slower if cold, or faster in aqueous solutions, forming nitrogen gas.

If you instead bubbled NO2 into aqueous NH4OH solution, I would imagine you would get a different reaction:

2NH4OH + 2NO2 --> NH4NO3 + NH4NO2

Nitrogen dioxide can also oxidize aqueous solutions of ammonium perchlorate, but the reaction tends to be very slow. Even at boiling temperatures, it takes around an hour. The perchlorate does not take part in the reaction.
"The production of Perchloric Acid", Thesis by Horrace C. Adams, California Institute of Technology, (1925)

(concentrated nitric acid generally cannot oxidize ammonium ions, although this subject is too complicated to go into detail here. The interesting thing is that concentrated nitric acid is a stronger oxidizing agent than nitrogen dioxide, but the exception here has to do with the reaction mechanism, since the active species responsible for the oxidizing power in nitric acid is protonated, carrying a positive charge, and thus ammonium ions are effectively rendered inert because they also carry a positive charge and repel in the solution.)

[Edited on 1-4-2013 by AndersHoveland]




I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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[*] posted on 1-4-2013 at 05:52


Quote: Originally posted by AndersHoveland  
Quote: Originally posted by Antiswat  

i decided to open a bottle of 25% NH4OH as well
NO2 + NH4OH > NH4NO2
or am i wrong?
the brown gas went very fast into white smoke alike other acid reactions in air with ammonia..
i also read that it very quickly goes into H2O and N2...


Actually, the gas phase reaction is:

(2)NO2 + (2)NO + (4)NH3 --> (2)NH4NO2 + (2)H2O + (2)N2

Dry ammonia gas reacts with the nitrogen dioxide and nitric oxide, at room temperature. (reaction investigated by Marcellin Berthelot)

Solid ammonium nitrite inside a tube explodes if heated on a water bath to between 60-70 °C. And the substance gradually decomposes at room temperature, slower if cold, or faster in aqueous solutions, forming nitrogen gas.

If you instead bubbled NO2 into aqueous NH4OH solution, I would imagine you would get a different reaction:

2NH4OH + 2NO2 --> NH4NO3 + NH4NO2
.......


First, I think it is interesting that it is reported that dry ammonia proceeds to react here as per one of my readings, I recall that dry ammonia apparently does not react with dry hydrogen chloride gas.

Also, if one opened a box of dilute aqueous ammonia, I would expect the presence of some moisture as well, so perhaps:

2 NH3 + (2 NO2 + H2O) -->2 NH3 + (HNO2 + HNO3) --> NH4NO2 + NH4NO3

and as NH4NO3 (or any unreacted NO2 for that matter) in the presence of moisture is acidic, any NH4NO2 formed is most likely not stable either. This could translate into a partial or complete decomposition into N2 (as suggested by Antiswat ) or even a possible detonation of the solid ammonium nitrite depending on temperature (in agreement with the reference cited by AndersHoveland) and/or the presence of dust, etc.

[EDIT] With respect to my prior comments on NH4Cl/NaNO2 where the heated NH4Cl forms only dry NH3 and HCl (which is assumed to react with the dry NaNO2 to form NO, NO2 and some water vapor), AndersHoveland cited reaction gives me more understanding on why the heated NH4Cl/NaNO2 behaves, in effect, like an energetic compound. More explicitly, dry NO and NO2 can also form NH4NO2 in the presence of only dry NH3.

[Edited on 1-4-2013 by AJKOER]
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[*] posted on 1-4-2013 at 06:30


BTW, if anyone needs some pure (99.9%) NaNO2, just drop me a line. All I ask is you cover the shipping.

Always be careful with nitrites, especially if you have made a nitrosamine. Highly teratogenic. Ascorbic acid, apparently, mitigates the ability of small amounts of nitrites (added to hot dogs, for instance), to cause any harm. NaNO2 is added to beef products to stop the beef from browning due to nitrogen fixation bacterium.




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[*] posted on 3-4-2013 at 07:25


i like the idea of NH4NO2 going off violently, what can be made can be controlled, and if small can be controlled very likely more can be controlled
SM2: very tempting.. ill consider it if i forget to keep thought on my economy in the near future (;

anyways guys im quite sure i did manage to form some NH4NO2 as i took dilute ammonia, then lead NO2 (+ NO?) through the ammonia, i got the NO2 from reaction of an somewhat unidentified metal, probably copper with nitric acid
when the ammonia somehow due to a physical law that has several times annoyed me i very suddenly saw a change in colour (tetraammine formation) and im currently trying also to get some TACN out of it (which i understand as not having an actually explosive property
but that N2 might be formed as a side reaction, and not being the NH4NO2 decomposing is pretty relevant..
i saw lots of white smoke forming, ofcourse more violent at the start but it seemed quite dense as it formed some layer with the air in the bottler, where the air was on top (??)
im evaporating the solution at very low heat, on top of my computer, not any sharp UV or anything.. hope to get some results.. (:




~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
http://www.trimen.pl/witek/calculators/stezenia.html
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