Adas
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Reaction between SO3 and nitrates
Hello. This question came to my mind today, and I can't find any answers for this.
My question is: Will this reaction make N2O5? Because I think it's a very useful compound if you can make it in significant quantities (>5g)
Thanks for your time.
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hissingnoise
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No is the short answer!
SO<sub>3</sub> dehydrates HNO<sub>3</sub>; (HNO<sub>3</sub> + SO<sub>3</sub> --->
NO<sub>2</sub><sup>+</sup> +
HSO<sub>4</sub><sup>-</sup>.
[Edited on 23-12-2011 by hissingnoise]
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Lambda-Eyde
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Hissingnoise, your equation doesn't... Equate.  Also, dinitrogen pentoxide
is the anhydride of nitric acid and can, according to Wiki, be prepared by the action of phosphorus pentoxide on nitric acid.
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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Adas
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Quote: Originally posted by hissingnoise  | No is the short answer!
SO<sub>3</sub> dehydrates HNO<sub>3</sub>; (HNO<sub>3</sub> + SO<sub>3</sub> --->
NO<sub>2</sub> + HSO<sub>4</sub><sup>-</sup>.
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But the anhydride of HNO3 is N2O5. And I am talking about nitrates, not HNO3. But, of course, if the temperature is too high, it decomposes to O2 and
NO2.
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smuv
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| Quote: | | Hissingnoise, your equation doesn't... Equate |
That equation is in everyone's organic chemistry textbook to rationalize electrophilic aromatic substitution w. mixed acids. I guess that doesn't
mean it's right...
"Titanium tetrachloride…You sly temptress." --Walter Bishop
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hissingnoise
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Fixed it L-E . . . (Christmas spirit --- hic?)
P<sub>4</sub>O<sub>10</sub> is a potent enough desiccant to even form some SO<sub>3</sub> in
H<sub>2</sub>SO<sub>4</sub>.
But for synthesis N<sub>2</sub>O<sub>5</sub> is prepared by either electrolytic oxidation of HNO<sub>3</sub> or by
N<sub>2</sub>O<sub>4</sub> oxidation by ozone.
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Adas
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| Quote: Originally posted by hissingnoise | Fixed it L-E . . . (Christmas spirit --- hic?)
P<sub>4</sub>O<sub>10</sub> is a potent enough desiccant to even form some SO<sub>3</sub> in
H<sub>2</sub>SO<sub>4</sub>.
But for synthesis N<sub>2</sub>O<sub>5</sub> is prepared by either electrolytic oxidation of HNO<sub>3</sub> or by
N<sub>2</sub>O<sub>4</sub> oxidation by ozone.
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But I was asking if the reaction I suggested is possible or not. Thanks anyways.
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hissingnoise
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IIRC, N<sub>2</sub>O<sub>5</sub> was first prepared by reacting AgNO<sub>3</sub> with dry
Cl<sub>2</sub> . . .
Good luck with that if you want to try it!
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AndersHoveland
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The reaction of sodium nitrate with sulfur trioxide, assuming the liquid sulfur trioxide is able to dissolve the solid sodium nitrate, will most
likely result in the formation of nitronium and pyrosulfate ions. The reaction could best be represented by this equation:
(2)NaNO3 + (4)SO3 --> Na2S2O7 + (NO2)2S2O7
| Quote: Originally posted by hissingnoise | N<sub>2</sub>O<sub>5</sub> was first prepared by reacting AgNO<sub>3</sub> with dry Cl<sub>2</sub> .
. .
Good luck with that if you want to try it!
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That is one reaction I am unable to believe. I would have thought the reaction would be just the opposite. Do you have a reference for this?
[Edited on 24-12-2011 by AndersHoveland]
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hissingnoise
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| Quote: |
That is one reaction I am unable to believe.
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The truth is out there man?
Somewhere?
[Edited on 25-12-2011 by hissingnoise]
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AndersHoveland
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Yes, silver chloride is an insoluble precipitate, and that would tend to shift the reaction equilibrium. But one would think the oxidizing nature of
N2O5 would be a more significant factor.
While 60% conc nitric acid does not appear to dissolve AgCl, I would think that much more concentrated HNO3 acid, possibly with concentrated sulfuric
acid, would oxidize it.
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AndersHoveland
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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.
Sulphur trioxide reacts with nitrogen pentoxide in carbon-tetrachloride solution, with the formation of a crystalline precipitate melting at 124° to
128° C.
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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Adas
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| Quote: Originally posted by AndersHoveland | 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.
Sulphur trioxide reacts with nitrogen pentoxide in carbon-tetrachloride solution, with the formation of a crystalline precipitate melting at 124° to
128° C. |
The percipitate is probably dinitronium sulfate. What else?
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AJKOER
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Per Atomistry.com (http://nitrogen.atomistry.com/nitrogen_pentoxide.html ) notes that Deville first isolated it by decomposing silver nitrate with dry chlorine.
Later, Meyer obtained it later from nitric acid by dehydrating with phosphorus pentoxide. On preparations, to quote:
"1. 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.
2. The most convenient method is by the dehydration of nitric acid. This is first obtained pure by repeated distillation with concentrated sulphuric
acid, and bubbling dry air through the final distillate in order to remove oxides of nitrogen. 200 grams of this "white fuming acid " are put into a
2-litre flask with a side arm at right angles to the neck, and 400 grams of phosphorus pentoxide are slowly added during cooling until pasty. The
flask is then attached by means of sealing-wax to three wash-bottles, the first containing glass-wool and phosphorus pentoxide, and the others being
empty, and immersed in a freezing mixture of ice and salt. The flask is heated on a water-bath to 60°-70° C. and a slow stream of air (dried by
concentrated sulphuric acid) is passed through the paste. The nitrogen pentoxide collects in the wash-bottles as a slightly yellow brittle solid:
2HNO3 = N2O5 + H2O
3. Nitrogen trioxide and nitrogen tetroxide are both oxidised by ozone to nitrogen pentoxide.
4. A mixture of nitrogen and oxygen can be converted into nitrogen pentoxide by means of the silent electric discharge in the presence of ozone."
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condennnsa
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Atochemistry is an awesome website, ajoker!! I did not know about it
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AndersHoveland
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| Quote: Originally posted by Adas | | 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 percipitate is probably dinitronium sulfate. What else? |
I suspect it may be dinitronium pyrosulfate, or, not unlikely, another compound with an even higher equivalent ratio of SO3.
This should be able to answer all your questions:
| 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.
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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(It is really off topic here, but we read so often about nitromethane being used as a solvent with strong acids in the presence of nitrating media;
actually it is very vulnerable to oxidation in alkaline solution. Nitromethane can also undergo disproportionation when heated in aqueous acidic
solution in one of the reactions which bears the name of Victor Meyer. The exact chemistry of these reactions are too complicated to discuss here, but
I just wanted to mention it.)
[Edited on 20-3-2013 by AndersHoveland]
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