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Author: Subject: Reaction between SO3 and nitrates
Adas
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[*] posted on 23-12-2011 at 08:45
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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[*] posted on 23-12-2011 at 09:39


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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[*] posted on 23-12-2011 at 09:55


Hissingnoise, your equation doesn't... Equate. :P 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.



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Adas
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[*] posted on 23-12-2011 at 09:57


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>;).




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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[*] posted on 23-12-2011 at 10:32


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...:(




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hissingnoise
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[*] posted on 23-12-2011 at 10:46


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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[*] posted on 23-12-2011 at 12:26


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.




But I was asking if the reaction I suggested is possible or not. Thanks anyways.




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hissingnoise
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[*] posted on 24-12-2011 at 09:30


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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[*] posted on 24-12-2011 at 15:55


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!


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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[*] posted on 25-12-2011 at 03:26


Quote:

That is one reaction I am unable to believe.

The truth is out there man?



Somewhere?






[Edited on 25-12-2011 by hissingnoise]
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[*] posted on 25-12-2011 at 10:28


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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[*] posted on 19-3-2013 at 14:57


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.




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[*] posted on 20-3-2013 at 08:11


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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[*] posted on 20-3-2013 at 10:13


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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[*] posted on 20-3-2013 at 12:23


Atochemistry is an awesome website, ajoker!! I did not know about it
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[*] posted on 20-3-2013 at 13:12


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



(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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