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Author: Subject: Toying with NCl3 idea tidbit
Morgan
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[*] posted on 9-8-2011 at 08:47
Toying with NCl3 idea tidbit


"An amine-based AGIL has been studied elsewhere,4,5 but neither lasing nor positive gain was reported. We developed a numerical simulation of the NCl3+HI+H2/H reaction system,6 and found that the injection order of the three gas species is the key to achieving positive gain. Other parameters that can be optimized include the flow rates of the gasses, the operating pressure, and the position of the optical axis."
http://spie.org/x26878.xml?ArticleID=x26878

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hissingnoise
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[*] posted on 9-8-2011 at 09:50


You don't "toy with NCl<sub>3</sub>' - it toys with you!
That's if you can call blowing off fingers and gouging eyes out "toying!

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Morgan
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[*] posted on 9-8-2011 at 11:31


Quote: Originally posted by hissingnoise  
You don't "toy with NCl<sub>3</sub>' - it toys with you!
That's if you can call blowing off fingers and gouging eyes out "toying!


From the good old days.
http://www.lateralscience.co.uk/oil/index.html
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AndersHoveland
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[*] posted on 9-8-2011 at 11:48


It was found that NCl3 forms much more readily when ammonium sulfate is used, instead of ammonium chloride. This is because a solution of ammonium sulfate has a higher pH.
When preparing NCl3, only VERY tiny quantities (a few drops) should ever be prepared at one time. Ideally, a plastic container should be used. Always wear plastic safety goggles!
Do not try storing the substance, it will expolode on its own, unless dissolved in carbon tetrachloride.

That laser article was interesting; the main problem of course would be producing the NCl3 on site, so probably not the most practical proposal.

[Edited on 9-8-2011 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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The WiZard is In
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[*] posted on 9-8-2011 at 13:33


Quote: Originally posted by Morgan  
"An amine-based AGIL has been studied elsewhere,4,5 but neither lasing nor positive gain was reported. We developed a numerical simulation of the NCl3+HI+H2/H reaction system,6 and found that the injection order of the three gas species is the key to achieving positive gain. Other parameters that can be optimized include the flow rates of the gasses, the operating pressure, and the position of the optical axis."
http://spie.org/x26878.xml?ArticleID=x26878



Nitrogen chloride is considered to be one of the most dangerous bodies to
handle, owing to the facility with which it explodes, by shock, friction, or contact
with various bodies.

M. Berthelot 1892

Nitrogen trichloride was discovered by my good friend, physician and professor
of physics at the École Polytechnique Pierre Louis Dulong. I first meet him at
Berthollet’s home at Arcueil just south of Paris, where Berthollet had settled
following his return from Napoleon’s abortive Egyptian campaign. All the greats
meet their; Berthollet’s neighbour Laplace, Arago, Bérand, Biot, Amédée
Berthollet (Claude’s son), Chaptal, Collet-Desostils, de Candolle, Gay-Lussac,
Humboldt, Malus, Poisson and Thernard. Napoleon showed his approval of our
meetings by allowing the use of the title “Société d’Arcueil” for our gatherings.

Pierre D. first published notice of his discovery in Schweigger’s J. Chem. Pharm.
8, 32 (1812). Shortly there after he lost and eye and three fingers when a sample
exploded in his laboratory!! Indeed, our mutual friend Humphry Davy was also
severely injured although happily not maimed by an unexpected explosion of a
small quantity of Pierre’s “une nouvelle substance detonnante”. [Later Gay-
Lussac and Thernard suffered from inhaling hydrogen fluoride fumes.]

“The preparation and handling of this compound requires the greatest care.
Every vessel employed must be washed by alkali-lye in order to free it from
grease; even grease from the fingers may cause an explosion. The substance
[yellow oily liquid] is very liable to spontaneous explosion, and thick gloves, and a
face shield are indispensable.” It is also possible to cause it to explode by
exposure to strong sun light or the light of a magnesium flame!"


The West American scientist Official Organ of the San Diego Society of Natural History
By Charles Russell Orcutt
1888

MARTYRS OF SCIENCE.
The scientific investigator, and notably the experimenting chemist, incurs dangers
that would be likely to appall the most valiant fighters, and he meets them calmly
and deliberately, not in hot blood and excitement. Familiarity with danger breeds
recklessness, if not contempt, and the chemist will coolly taste with the tip of his
tongue any unknown liquid that may be handed to him. Prussic acid is about the only
thing he will not put into his mouth, but he can recognize that without tasting it. He
will make all sorts of chemical combinations, and try them to see if they are explosive
or otherwise dangerous, and sometimes he discovers in a very unpleasant way that
his new product is loaded. Dulong, in 1811 discovered chloride of nitrogen, and when
he recovered consciousness, amid the wreck of all the glassware in his laboratory, he
made a note with his unwounded hand, of the fact that chloride of nitrogen is highly
explosive. About a year afterward, in trying to ascertain the exact conditions
necessary to explode the compound, he lost an eye and two fingers. Davy, in the
same year, had a similar adventure with the new explosive. A tube containing a small
quantity was suddenly shivered to atoms without any apparent cause, and a piece of
the glass struck Davy in the corner of the right eye, disabling him from further
immediate experiments. In July, 1813, Davy set about finding out what was the
matter with chloride of nitrogen, and attempted an analysis by mercury. The stuff
went off again as usual, but Davy had protected his eyes with a thick plate of glass,
and he was wounded only in the head and hands. Faraday was an assistant to Davy
at this time. He was holding a small tube containing a few grains of the chloride
between his thumb and finger, when a sudden and wholly unprovoked explosion
occurred, stunning him and badly tearing his hand.


djh
----
The explosion removed the windows,
the door and most of the chimney.
It was the sort of thing you expected in
the Street of Alchemists. The neighbours
preferred explosions, which were at least
identifiable and soon over. They were better
than the smells, which crept up on you.

Terry Pratchett
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hissingnoise
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[*] posted on 10-8-2011 at 00:56


Berthollet and Berthelot, both pioneers of energetic science used to confuse me no end when I first encountered their names!
Such coincidence . . .

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AJKOER
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[*] posted on 17-8-2011 at 03:32


I would send the small quantity of NCl3 vapors directly into the intended reaction mixture, but be prepared for the consequences.

Make no attempt to collect or store it. Note, even a few millimeters of the oily yellowish mixture can produce a sizable explosion to induce injury. In other words, being underestimated (as it is, in fact, one of the most powerful explosive) the chemists over produces it with injurious results.

For those who foolishly would condense the vapor, also avoid strong light (certain uv frequency especially) and exposure to organic compounds (including dust?).

Good Luck.
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MrTechGuy1995
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[*] posted on 21-9-2011 at 16:08


Learn from my idiotic experimentation: http://www.youtube.com/watch?v=YX5b_f0W3x0
This is the explosive decomposition of it when it was in solution.
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AndersHoveland
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[*] posted on 21-9-2011 at 17:40
Cl2 + NH3


The reaction of NH3 solution with Cl2 is highly variable, depending on the pH and reactant ratios.

If conditions are too acidic, there will not be any reaction.
If conditions are only slightly acidic, then dangerously sensitive explosive droplets of NCl3 will form. It is actually an equilibrium reaction:
NCl3 + (4)HCl <==> NH4Cl + (3)Cl2

If conditions are mildly basic, chloramine gas will tend to form.
If conditions are more basic, and an excess of ammonia is used, some hydrazine will form.

Under the acidic reactions, the presence of excess ammonia will lead to nitrogen gas being formed. Under basic conditions, the situation is just the opposite, if there is not an excess of ammonia, more of the ammonia will be oxidized to nitrogen.

Chloramine is somewhat chemically unstable, it will gradually decompose to elemental nitrogen after a short time, much faster if it is more concentrated.

(3)NH2Cl --> NH4Cl + (2)HCl + N2

Reaction of bromine with ammonia has a similar chemistry to that of chlorine. Bromamine and dibromamine can form. The main difference, however, is that NBr3 cannot be made in the same way that NCl3 can. (Researchers have actually prepared NBr3, but its synthesis is much more complicated and involves other obscure regents)

"Real" nitrogen triiodide, like NBr3, is also very difficult to prepare. The product of NH4OH and elemental iodine, which is sometimes incorrectly reffered to as "nitrogen triiodide", actually has a formula NI3*NH3, where the adduct is actually polymerized by cross-linking of the iodine atoms.

[Edited on 22-9-2011 by AndersHoveland]
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Endimion17
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[*] posted on 21-9-2011 at 18:45


Quote: Originally posted by MrTechGuy1995  
Learn from my idiotic experimentation: http://www.youtube.com/watch?v=YX5b_f0W3x0
This is the explosive decomposition of it when it was in solution.


No, it was not in the solution. It is not soluble in aqueous solutions. You made a few droplets at most, which detonated because of the agitation. If the solution detonated, you wouldn't have your video... or eyes, skull?

But you're true about the "idiotic" part. Making nitrogen trichloride by bubbling chlorine through the solution? What were you thinking?

If chlorine is used, it is done by the Dulong-Davy-Thénard-Berthollet method (I'm mentioning all of the four guys because I don't know whose is it originally, perhaps Dulong's).
A quantity of chlorine in a tube/flask is inverted over a solution of ammonium chloride and the droplets are collected in a small lead plate, but any ordinary watch glass would be ok if squeaky clean.

The whole setup is not moving. Only the droplets are falling down.
Davy said he was injured by a quantity "scarcely as large as a grain of mustard seed". I find that hard to believe. He was either lying, or made this mistake, acting like Johnny here.




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AndersHoveland
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[*] posted on 21-9-2011 at 18:57


I tried passing chlorine into a solution of ammonium nitrate, using plastic bottles, taped drinking straws, and a plastic container. There was no formation of any oily droplets.

Perhaps this is because ammonium nitrate is not a good ammonium salt to use, as it might be too acidic, or perhaps contact with plastic caused the NCl3 to degrade faster than it formed.

Supposedly, ammonium sulfate works much better than ammonium chloride, since its solutions are less acidic, and one of the ammonias can more freely come off.

It may likely be safer to use a plastic container, rather than glass. If the NCl3 detonates, it will cause the glass to shatter, and sharp fragments will be hurled out.

[Edited on 22-9-2011 by AndersHoveland]
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MrTechGuy1995
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[*] posted on 21-9-2011 at 18:57


Quote: Originally posted by Endimion17  
Quote: Originally posted by MrTechGuy1995  
Learn from my idiotic experimentation: http://www.youtube.com/watch?v=YX5b_f0W3x0
This is the explosive decomposition of it when it was in solution.


No, it was not in the solution. It is not soluble in aqueous solutions. You made a few droplets at most, which detonated because of the agitation. If the solution detonated, you wouldn't have your video... or eyes, skull?

But you're true about the "idiotic" part. Making nitrogen trichloride by bubbling chlorine through the solution? What were you thinking?

If chlorine is used, it is done by the Dulong-Davy-Thénard-Berthollet method (I'm mentioning all of the four guys because I don't know whose is it originally, perhaps Dulong's).
A quantity of chlorine in a tube/flask is inverted over a solution of ammonium chloride and the droplets are collected in a small lead plate, but any ordinary watch glass would be ok if squeaky clean.

The whole setup is not moving. Only the droplets are falling down.
Davy said he was injured by a quantity "scarcely as large as a grain of mustard seed". I find that hard to believe. He was either lying, or made this mistake, acting like Johnny here.


Ya, I know It's stupid, maybe it's what it takes to graphically document things. Seeing as I don't have proper equipment. But after this fuck up, I don't intend to do this again, unless I have some sort of plexiglass shield, and Proper glassware.

[Edited on 22-9-2011 by MrTechGuy1995]
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AndersHoveland
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[*] posted on 21-9-2011 at 19:05


Another potential way to make NCl3 is to mix NH4Cl with TCCA (pool chlorinator).

TCCA works by being able to partially hydrolyse into HOCl and cyanuric acid.

One simple formula to express the formation of nitrogen trichloride is:

NH3 + (3)HOCl <==> NCl3 + (3)H2O
(the reaction can go both ways)

Although this formula is, in fact, accurate, the whole reaction is actually more complicated than this. The formula is just an easy way to conceptualize what is happening.

[Edited on 22-9-2011 by AndersHoveland]
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AndersHoveland
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[*] posted on 29-5-2012 at 23:57


Here is an interesting article on some of the unusual chemistry of NCl3, apparently it reacts with nitric oxide at -130 C to form nitrous oxide,
http://pubs.acs.org/doi/abs/10.1021/ja01374a015
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AJKOER
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[*] posted on 31-5-2012 at 07:37


A point I would add further than adding heat to your Ammonium chloride (or sulfate) will definitely aid in the creation of NCl3 (but not necessarily required) given its a highly endothermic compound. The reference below cites the following reaction in Chloroform:

NH4Cl + Cl2 --> NCl3 + 4 HCl

where the concentration of NCl3 in CHCl3 should not exceed 18% else one has an extremely explosive substance ( I would always act like this is the case).

In recent years I have seen the suggested use of NCl3 in dilute Chloroform solutions in organic synthesis as it also is a highly reactive compound. Before providing the link, I strongly advised all to read the author's warning message about working with micro quantities and safety measures, and even NOT to repeat many of the experiments cited.

http://www.chemexplore.net/BookP4s.pdf
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