plante1999
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The road to nitronium nitrate and to "nitrateum"
I had a idea few days ago, the idea was to make the equivalent of oleum, but with nitric acid. First I was in need to make nitronium nitrate. After
search in books, I found a procedure suitable for me, using silver nitrate and chlorine, so I made 20ml of conc. HNO3 and dissolved silver, drying the
solution to get the salt.
Now I will need to make more nitric acid,and make the nitronium nitrate.
What do you think of the "nitrateum" idea?
The silver nitrate:
<img src="http://i1103.photobucket.com/albums/g469/plante1999/DSC00076_zps22be3630.jpg" width="600" />
<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: reduced
image size(s)]
[Edited on 20.1.14 by bfesser]
I never asked for this.
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Lambda-Eyde
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Quote: Originally posted by plante1999  | | First I was in need to make nitronium nitrate. After search in books, I found a procedure suitable for me, using silver nitrate and chlorine, so I
made 20ml of conc. HNO3 and dissolved silver, drying the solution to get the salt. |
Source?
"Nitrateum" being...?
This just in: 95,5 % of the world population lives outside the USA
You should really listen to ABBAPlease drop by our IRC channel: #sciencemadness @ irc.efnet.org
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Mailinmypocket
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The equivalent of oleum, but with nitric acid? Do you mean fuming nitric acid? Oleum is a solution of SO3 in sulfuric acid... So by nitrateum (is the
"eum" added at the end just like olEUM?) do you mean nitric acid with excess NO3 in it? I have never heard of this before, where did you read about
this material?
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hissingnoise
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| Quote: | | I have never heard of this before, where did you read about this material? |
Wiki has this.
I doubt the method produces the anhydride in quantity . . .
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Mixell
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I think he means a solution of dinitrogen pentoxide (N2O5) in nitric acid, which is usually structured as a salt (nitronium nitrate).
But N2O5 is pretty unstable, decomposing into nitrogen dioxide and oxygen.
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woelen
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I can imagine that a solution of N2O5 in nitric acid is more stable than the pure salt, but I do not know the solubility of N2O5 in HNO3. This is
something which plante1999 should try.
There is one big difference though: N2O5 can explode, apparently without any visible reason and SO3 cannot explode. This most likely is the reason why
a solution of N2O5 in HNO3 is not an article of commerce while a solution of SO3 in H2SO4 is.
So, plante1999, please be careful with this!! After your experience with HNO3/H2O2 you must not think of an explosion of HNO3/N2O5 resulting in
getting splashes of this material on your skin.
If you have P4O10, then you could try adding some HNO3 to this and see if you can get the resulting material dissolved in HNO3.
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plante1999
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More data:
I dissolved 14.5 g of silver to make the nitrate.
When I talk about ''nitrateum'' I actually talk about the possibility of a solution of nitronium nitrate in anhydrous nitric acid, the name refering
to a solution of sulphur trioxide in anhydrous sulphuric acid, cvommonly named oleum. I got the idea recently, I read many things about it, and many
source are contradictory about its decomposition. Some say it is quite stable, other say it spontaneusly detonate. The objective is to experiment the
propretie of the hypotetical material.
I will, of course use face shield and gloves, and try on a small scale. I already have enough of silver nitrate stain... I did wear gloves until the
salt was dry, but is seam even AgNO3 dust make stains.
I do not own P2O5, thats is the main reason why I need to use silver nitrate.
Reference for silver nitrate reaction with chlorine:
P.199: http://books.google.ca/books?id=8ocDAAAAQAAJ&pg=PA200&am...
P. 157 : http://books.google.ca/books?id=7QwNAAAAYAAJ&pg=PA157&am...
[Edited on 13-3-2013 by plante1999]
[Edited on 13-3-2013 by plante1999]
[Edited on 13-3-2013 by plante1999]
I never asked for this.
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Vargouille
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Nitric pentoxide in nitric acid is not unknown, up to 35-40% solutions. It is the subject of this article, and is a nitric pentoxide source used in this article. Incidentally, it may be more efficient to produce a solution of N2O5 in HNO3 by way of electrolysis of a solution of NO2 in nitric
acid, as this patent, borrowing from a Soviet patent, suggests.
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AndersHoveland
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| Quote: Originally posted by plante1999 | | to make nitronium nitrate. After search in books, I found a procedure suitable for me, using silver nitrate and chlorine, |
I have read of this alleged reaction before somewhere, but have severe doubts that it actually works. I have a feeling it is not so simple.
| Quote: | Deville first isolated this oxide by decomposing silver nitrate with dry chlorine. Meyer obtained it later from nitric acid by dehydrating with
phosphorus pentoxide.
Dry chlorine reacts with silver nitrate at 95° C., and as soon as the action has started the mixture is cooled to 50°-60° C. The nitrogen pentoxide
evolved is separated from the oxygen by condensing in a U-tube immersed in a freezing mixture. No corks or rubber joints may be used owing to the
corrosive action of the gas:
4AgNO3 + 2Cl2 = 4AgCl + 2N2O5 + O2.
It is also produced by the reaction between nitryl chloride and silver nitrate:
NO2Cl + AgNO3 = AgCl + N2O5.
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The first equation is also mentioned in a brief entry on page 145 of The Medical student's manual of chemistry
By Rudolph August Witthaus , Sixth Edition, (1906)
| Quote: |
"When Cl is passed over silver nitrate at 60 °C, or when pure nitric acid (63 parts) is mixed with phosphorous pentoxide (71 parts) at low
temperature, N2O5 is formed."
Chemical lecture notes, (taken from Prof. C.O. Curtman's lectures at the St. Louis College of Pharmacy), Henry Milton Whelpley, Charles O.
Curtman, (1895), p104
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| Quote: |
"Nitryl chloride is prepared most conveniently by reacting chlorosulfonic acid with anhydrous nitric acid at 0°.
An older preparation method involves passing dry chlorine gas over dry silver nitrate heated to about 100 °C. the gaseous reaction products are
allowed to cool to low temperature. After several hours, nitryl condenses to a pale yellowish-brown liquid. Chlorine is removed be purging with CO2."
2 AgNO3 + 2 Cl2 → 2 NO2Cl + 2 AgCl + O2
Handbook of Inorganic Chemicals, Pradyot Patnaik
http://site.iugaza.edu.ps/bqeshta/files/2010/02/94398_16.pdf
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I would have thought that concentrated nitric acid would just oxidize the chloride ions to chlorine, but perhaps this does not work on AgCl.
2 NaCl + 4 HNO3 --> 2 NaNO3 + 2 H2O + 2 NO2 + Cl2
In the compound AgCl, the chloride ions act as ligands to the silver, explaining why it is insoluble in water. It can be viewed as more a covalently
bonded compound, and indeed the presence of Ag+ ions significantly changes the redox values of equations involving the chloride ion. (for example,
FeCl3 solutions should actually be able to oxidize the surface of silver)
In other words, the formation of AgCl is highly favorable, and may be what is driving the reaction.
[Edited on 19-3-2013 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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plante1999
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In books, I saw that NO2Cl reacted with AgNO3 to make nitronium nitrate, it is the reaction that make it all in fact.
I never asked for this.
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blogfast25
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| Quote: Originally posted by AndersHoveland |
I would have thought that concentrated nitric acid would just oxidize the chloride ions to chlorine, but perhaps this does not work on AgCl.
2 NaCl + 4 HNO3 --> 2 NaNO3 + 2 H2O + 2 NO2 + Cl2
[Edited on 19-3-2013 by AndersHoveland] |
The Standard Reduction Potentials don't bear this out:
Cl-/Cl2 Eox = -1.36 V
NO3-/NO Ered = +0.96 V
Ecell = - 0.4 V, Delta G > 0, no reaction.
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Random
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Dinitrogen pentoxide sounds like scary stuff
what about chlorinating AgNO3 in CHCl3? stay safe though certainly doesnt look harmless
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AndersHoveland
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The nitronium ion, NO2+ has a reduction potential of at least 1.6 V.
Concentrated solutions of nitric acid contain nitronium ions in equilibrium. This is why nitric acid acts as an oxidizer, and why more concentrated
nitric acid is more reactive, while very dilute nitric acid is hardly an oxidizer at all.
The oxidation of chloride by concentrated nitric acid may well be an endothermic reaction, but that does not mean it will not take place at ambient
temperatures. There may well be some equilibrium.
Just to remove any further doubts you may have, Br-/Br2 is -1.06 V, yet concentrated nitric acid can still oxidize it. So obviously the standand
reduction value for nitric acid can be a little misleading.
Indeed, one source describes the equilibrium reaction between nitric acid and bromide ions:
"Kinetic study of the autocatalytic nitric acid-bromide reaction and its reverse, the nitrous acid-bromine reaction", Istvan Lengyel , Istvan Nagy
, Gyorgy Bazsa, J. Phys. Chem., 1989, 93 (7), pp 2801–2807
Just try heating a solution of sodium nitrate in hydrochloric acid, brown fumes will come out.
If you were wondering how exactly the formation of AgCl effects the reduction potentials:
Ag+ + e- --> Ag = +0.80 V
AgCl + e- --> Ag + Cl- = +0.223 V
So the formation/precipitation of AgCl will make a reaction 0.577 V more favorable. That is theoretically enough to overcome the oxidation of chloride
ions by nitronium ions, so one would not expect the following reaction to be favorable:
2 AgCl + 2 NO2+ --> 2 Ag+ + 2 NO2 + Cl2
N2O5 acts as a source of NO2+ cations in reactions, thus the reduction potential calculations suggest N2O5 should be able to
exist perfectly well in the presence of AgCl.
[Edited on 20-3-2013 by AndersHoveland]
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franklyn
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Gleaned from this cited reference Handbook of Inorganic Chemicals
Dinitrogen Tetroxide reacts with Ozone to yield Dinitrogen Pentoxide & Oxygen
2 NO2 + O3 => N2O5 + O2
The liquid needs to be cooled at least below room temperature as
N2O4 boils at just 21 ºC and solidifies below - 9 ºC
Pure N2O5 melts at 30 ºC and decomposes above 45 ºC.
Dinitrogen Pentoxide , the anhydride of Nitric acid , is also obtained by treating
very cold Nitric acid (-10 ºC ) with Diphosphorus Pentoxide.
2 HNO3 + P2O5 => 2 HPO3 + N2O5
I'm just wondering out loud here but what about Sulfer Trioxide
2 HNO3 + SO3 => H2SO4 + N2O5
Partitioning with CH2Cl2 and chilling should then precipitate crystal N2O5
.
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AndersHoveland
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Perhaps, I am not really sure. But too much SO3 would further react with the N2O5.
At least one reference in the literature makes me suspect that SO3 can not be used to obtain N2O5, at least not by direct reaction:
| Quote: |
Nitronium hydorgen disulfate was prepared by treating nitric acid with more than two molecular portions of sulphur trioxide in nitromethane solution,
from which the salt crystallised:
HNO3 + 2SO3 = (NO2+)(HS2O7‒ ‒)
The same salt resulted from all attempts to prepare nitronium hydrogen sulfate.
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| Quote: |
Normal nitronium disulfate was also produced in the reaction between nitric acid and sulphur trioxide, but it could not thus be obtained free from the
hydrogen disulphate. It was prepared in pure form by treating dinitrogen pentoxide with less than two molecules of sulphur trioxide:
N2O5 + 2SO3 = (NO2+)2(S2O7‒ ‒)
Normal nitronium trisulfate was obtained in a pure state when dinitrogen pentoxide was treated with more than three molecular proportions of sulfur
trioxide:
N2O5 + 3 SO3 =
(NO2+)2(S3O10- -)
No more than three molecules of sulphur trioxide could be induced to enter into reaction with dinitrogen pentoxide.
Chemistry of nitronium salts: Isolation of some nitronium salts, D. R. Goddard, E. D. Hughes, C. K. Ingold, J. Chem. Soc., 1950
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Sulphur trioxide reacts with nitrogen pentoxide in carbon-tetrachloride solution, with the formation of a crystalline precipitate melting at 124° to
128° C.
Nitric acid dissolves nitrogen pentoxide, and a definite compound, 2HNO3.N2O5, has been obtained which is liquid at ordinary temperatures but
solidifies at 5° C.
[Edited on 21-3-2013 by AndersHoveland]
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madscientist
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| Quote: | | Sulphur trioxide reacts with nitrogen pentoxide in carbon-tetrachloride solution, with the formation of a crystalline precipitate melting at 124° to
128° C. |
The reaction of SO<sub>3</sub> with CCl<sub>4</sub> was used by the Italians in WWI to produce phosgene for the war effort.
This reaction was also accidentally rediscovered by Gilbert Stork (in my mind, the greatest chemist to ever live) in the 70s, which nearly cost him
his life and that of a colleague. If a mind of that caliber can make an oversight that serious, do you really think you are prepared?
| Quote: | “I was at Columbia, and Bruce Ganem, who is now a professor at Cornell, was in a lab across the way from my office. I found a bottle of SO3, which
is not that stable and had crystallized inside the bottle, which normally looks like Karo syrup, like molasses, and you pour it through a small
opening. This thing couldn't be poured out and the question was, how do I get rid of this stuff? And so the idea was to find some solvent, some inert
solvent, dissolve it, and pour it gently into ice. And as the solvent, I decided on carbon tetrachloride...
“To this date, I don't know what happened. There may have been a metallic impurity somewhere that catalyzed, ripping out one of the chlorines from
CCl4 in this extremely acidic medium... it was bubbling furiously, the bottle cracked in the hood. Black crap was coming all over the place, and I
could detect what I was convinced was the smell of phosgene.
“And I remember the dilemma that I had. I thought, 'Should I tell Bruce that he will probably die during the night or should I keep it quiet and
just see what happens?' Well, the truth is, it was probably low enough in concentration, that nothing happened. But I remember I was really
frantically concerned.” |
Angew. Chem. Int. Ed. 2012, 51, 3012 – 3023
If you can smell phosgene, you are in great danger. Without a fumehood that exhausted far from his vicinity, he would've been toast.
[Edited on 27-3-2013 by madscientist]
I weep at the sight of flaming acetic anhydride.
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The_Davster
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Keep it in the freezer (-20) and you can keep it for months no problem.
| Quote: Originally posted by woelen |
There is one big difference though: N2O5 can explode, apparently without any visible reason and SO3 cannot explode. |
Source? I and others have made it from NO2 and ozone. Many gram scale. Never a problem. I would suspect this is from the presence of organic
impurities.
Make it just before you use it though. Every time you take it out of the freezer to manipulate, as it warms it decomposes releasing brown fumes of
NO2.
Fun stuff. Nitrates almost anything
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woelen
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I have read about the instability of N2O5 in many different text books. Older ones from before WW II, but also newer ones (e.g. Chemistry of the
Elements by Earnshaw). Some texts even talk about spontaneous apparently without reason violent or even explosive decomposition of N2O5 and all texts
state that the material can only be stored for a limited amount of time (e.g. days, or at most weeks if stored in the cold).
But as I already wrote in another thread (the thread about ammonium nitrite), books sometimes can be wrong. Your experience with real N2O5 tells me
more than texts in books, even if these books are from respectable publishers and authors.
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