Dinitrogen pentoxide (also known as nitrogen pentoxideornitric anhydride) is the chemical compound with the formulaN2O5. It is one of the binary nitrogen oxides, a family of compounds that only contain nitrogen and oxygen. It exists as colourless crystals that sublime slightly above room temperature, yielding a colorless gas.[4]
Dinitrogen pentoxide is an unstable and potentially dangerous oxidizer that once was used as a reagent when dissolved in chloroform for nitrations but has largely been superseded by nitronium tetrafluoroborate (NO2BF4).
N2O5 is a rare example of a compound that adopts two structures depending on the conditions. The solid is a salt, nitronium nitrate, consisting of separate nitronium cations[NO2]+ and nitrate anions[NO3]−; but in the gas phase and under some other conditions it is a covalently-bound molecule.[5]
Pure solid N2O5 is a salt, consisting of separated linear nitronium ionsNO+2 and planar trigonal nitrate anions NO−3. Both nitrogen centers have oxidation state +5. It crystallizes in the space group D4 6h (C6/mmc) with Z = 2, with the NO−3 anions in the D3h sites and the NO+2 cations in D3d sites.[8]
The vapor pressure P (in atm) as a function of temperature T (inkelvin), in the range 211 to 305 K (−62 to 32 °C), is well approximated by the formula
being about 48 torr at 0 °C, 424 torr at 25 °C, and 760 torr at 32 °C (9 °C below the melting point).[9]
In the gas phase, or when dissolved in nonpolar solvents such as carbon tetrachloride, the compound exists as covalently-bonded molecules O2N−O−NO2. In the gas phase, theoretical calculations for the minimum-energy configuration indicate that the O−N−O angle in each −NO2 wing is about 134° and the N−O−N angle is about 112°. In that configuration, the two −NO2 groups are rotated about 35° around the bonds to the central oxygen, away from the N−O−N plane. The molecule thus has a propeller shape, with one axis of 180° rotational symmetry (C2) [10]
When gaseous N2O5 is cooled rapidly ("quenched"), one can obtain the metastable molecular form, which exothermically converts to the ionic form above −70 °C.[11]
Another laboratory process is the reaction of lithium nitrateLiNO3 and bromine pentafluorideBrF5, in the ratio exceeding 3:1. The reaction first forms nitryl fluorideFNO2 that reacts further with the lithium nitrate:[8]
BrF5 + 3 LiNO3 → 3 LiF + BrONO2 + O2 + 2 FNO2
FNO2 + LiNO3 → LiF + N2O5
The compound can also be created in the gas phase by reacting nitrogen dioxideNO2orN2O4 with ozone:[13]
2 NO2 + O3 → N2O5 + O2
However, the product catalyzes the rapid decomposition of ozone:[13]
Dinitrogen pentoxide reacts with water (hydrolyses) to produce nitric acidHNO3. Thus, dinitrogen pentoxide is the anhydride of nitric acid:[11]
N2O5 + H2O → 2 HNO3
Solutions of dinitrogen pentoxide in nitric acid can be seen as nitric acid with more than 100% concentration. The phase diagram of the system H2O−N2O5 shows the well-known negative azeotrope at 60% N2O5 (that is, 70% HNO3), a positive azeotrope at 85.7% N2O5 (100% HNO3), and another negative one at 87.5% N2O5 ("102% HNO3").[15]
Dinitrogen pentoxide eventually decomposes at room temperature into NO2 and O2.[17][13] Decomposition is negligible if the solid is kept at 0 °C, in suitably inert containers.[8]
Decomposition of dinitrogen pentoxide at high temperatures[edit]
Dinitrogen pentoxide between high temperatures of 600 and 1,100 K (327–827 °C), is decomposed in two successive stoichiometric steps:
N2O5 → NO2 + NO3
2 NO3 → 2 NO2 + O2
In the shock wave, N2O5 has decomposed stoichiometrically into nitrogen dioxide and oxygen. At temperatures of 600 K and higher, nitrogen dioxide is unstable with respect to nitrogen oxideNO and oxygen. The thermal decomposition of 0.1 mM nitrogen dioxide at 1000 K is known to require about two seconds.[19]
Decomposition of dinitrogen pentoxide in carbon tetrachloride at 30 °C[edit]
Apart from the decomposition of N2O5 at high temperatures, it can also be decomposed in carbon tetrachlorideCCl4 at 30 °C (303 K).[20] Both N2O5 and NO2 are soluble in CCl4 and remain in solution while oxygen is insoluble and escapes. The volume of the oxygen formed in the reaction can be measured in a gas burette. After this step we can proceed with the decomposition, measuring the quantity of O2 that is produced over time because the only form to obtain O2 is with the N2O5 decomposition. The equation below refers to the decomposition of N2O5inCCl4:
2 N2O5 → 4 NO2 + O2(g)
And this reaction follows the first order rate law that says:
Decomposition of nitrogen pentoxide in the presence of nitric oxide[edit]
N2O5 can also be decomposed in the presence of nitric oxideNO:
N2O5 + NO → 3 NO2
The rate of the initial reaction between dinitrogen pentoxide and nitric oxide of the elementary unimolecular decomposition.[21]
Dinitrogen pentoxide, for example as a solution in chloroform, has been used as a reagent to introduce the −NO2 functionality in organic compounds. This nitration reaction is represented as follows:
N2O5 + Ar−H → HNO3 + Ar−NO2
where Ar represents an arene moiety.[22] The reactivity of the NO+2 can be further enhanced with strong acids that generate the "super-electrophile" HNO2+2.
In this use, N2O5 has been largely replaced by nitronium tetrafluoroborate[NO2]+[BF4]−. This salt retains the high reactivity of NO+2, but it is thermally stable, decomposing at about 180 °C (into NO2F and BF3).
Dinitrogen pentoxide is relevant to the preparation of explosives.[7][23]
In the atmosphere, dinitrogen pentoxide is an important reservoir of the NOx species that are responsible for ozone depletion: its formation provides a null cycle with which NO and NO2 are temporarily held in an unreactive state.[24]Mixing ratios of several parts per billion by volume have been observed in polluted regions of the nighttime troposphere.[25] Dinitrogen pentoxide has also been observed in the stratosphere[26] at similar levels, the reservoir formation having been postulated in considering the puzzling observations of a sudden drop in stratospheric NO2 levels above 50 °N, the so-called 'Noxon cliff'.
Variations in N2O5 reactivity in aerosols can result in significant losses in tropospheric ozone, hydroxyl radicals, and NOx concentrations.[27] Two important reactions of N2O5 in atmospheric aerosols are hydrolysis to form nitric acid[28] and reaction with halide ions, particularly Cl−, to form ClNO2 molecules which may serve as precursors to reactive chlorine atoms in the atmosphere.[29][30]
N2O5 is a strong oxidizer that forms explosive mixtures with organic compounds and ammonium salts. The decomposition of dinitrogen pentoxide produces the highly toxic nitrogen dioxide gas.
^Simon, Arndt; Horakh, Jörg; Obermeyer, Axel; Borrmann, Horst (1992). "Kristalline Stickstoffoxide — Struktur von N2O3 mit einer Anmerkung zur Struktur von N2O5". Angewandte Chemie (in German). 104 (3). Wiley: 325–327. Bibcode:1992AngCh.104..325S. doi:10.1002/ange.19921040321.
^ abcdeWilson, William W.; Christe, Karl O. (1987). "Dinitrogen pentoxide. New synthesis and laser Raman spectrum". Inorganic Chemistry. 26 (10): 1631–1633. doi:10.1021/ic00257a033.
^McDaniel, A. H.; Davidson, J. A.; Cantrell, C. A.; Shetter, R. E.; Calvert, J. G. (1988). "Enthalpies of formation of dinitrogen pentoxide and the nitrate free radical". The Journal of Physical Chemistry. 92 (14): 4172–4175. doi:10.1021/j100325a035.
^Parthiban, S.; Raghunandan, B.N.; Sumathi, R. (1996). "Structures, energies and vibrational frequencies of dinitrogen pentoxide". Journal of Molecular Structure: Theochem. 367: 111–118. doi:10.1016/S0166-1280(96)04516-2.
^ abcHolleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils (ed.), Inorganic Chemistry, translated by Eagleson, Mary; Brewer, William, San Diego/Berlin: Academic Press/De Gruyter, ISBN0-12-352651-5
^Osborne, Bruce A.; Marston, George; Kaminski, L.; Jones, N.C; Gingell, J.M; Mason, Nigel; Walker, Isobel C.; Delwiche, J.; Hubin-Franskin, M.-J. (2000). "Vacuum ultraviolet spectrum of dinitrogen pentoxide". Journal of Quantitative Spectroscopy and Radiative Transfer. 64 (1): 67–74. Bibcode:2000JQSRT..64...67O. doi:10.1016/S0022-4073(99)00104-1.
^ abcYao, Francis; Wilson, Ivan; Johnston, Harold (1982). "Temperature-dependent ultraviolet absorption spectrum for dinitrogen pentoxide". The Journal of Physical Chemistry. 86 (18): 3611–3615. doi:10.1021/j100215a023.
^Schott, Garry; Davidson, Norman (1958). "Shock Waves in Chemical Kinetics: The Decomposition of N2O5 at High Temperatures". Journal of the American Chemical Society. 80 (8): 1841–1853. doi:10.1021/ja01541a019.
^Lloyd, L.; Wyatt, P. A. H. (1955). "The vapour pressures of nitric acid solutions. Part I. New azeotropes in the water–dinitrogen pentoxide system". J. Chem. Soc.: 2248–2252. doi:10.1039/JR9550002248.
^Wilkins, Robert A.; Hisatsune, I. C. (1976). "The Reaction of Dinitrogen Pentoxide with Hydrogen Chloride". Industrial & Engineering Chemistry Fundamentals. 15 (4): 246–248. doi:10.1021/i160060a003.
^Gruenhut, N. S.; Goldfrank, M.; Cushing, M. L.; Caesar, G. V.; Caesar, P. D.; Shoemaker, C. (1950). "Nitrogen(V) Oxide (Nitrogen Pentoxide, Dinitrogen Pentoxide, Nitric Anhydride)". Inorganic Syntheses. Inorganic Syntheses. pp. 78–81. doi:10.1002/9780470132340.ch20. ISBN9780470132340.
^Frenck, C.; Weisweiler, W. (2002). "Modeling the Reactions Between Ammonia and Dinitrogen Pentoxide to Synthesize Ammonium Dinitramide (ADN)". Chemical Engineering & Technology. 25 (2): 123. doi:10.1002/1521-4125(200202)25:2<123::AID-CEAT123>3.0.CO;2-W.
^Schott, Garry; Davidson, Norman (1958). "Shock Waves in Chemical Kinetics: The Decomposition of N2O5 at High Temperatures". Journal of the American Chemical Society. 80 (8): 1841–1853. doi:10.1021/ja01541a019.
^Wilson, David J.; Johnston, Harold S. (1953). "Decomposition of Nitrogen Pentoxide in the Presence of Nitric Oxide. IV. Effect of Noble Gases". Journal of the American Chemical Society. 75 (22): 5763. doi:10.1021/ja01118a529.
^Talawar, M. B. (2005). "Establishment of Process Technology for the Manufacture of Dinitrogen Pentoxide and its Utility for the Synthesis of Most Powerful Explosive of Today—CL-20". Journal of Hazardous Materials. 124 (1–3): 153–64. doi:10.1016/j.jhazmat.2005.04.021. PMID15979786.
^Finlayson-Pitts, Barbara J.; Pitts, James N. (2000). Chemistry of the upper and lower atmosphere : theory, experiments, and applications. San Diego: Academic Press. ISBN9780080529073. OCLC162128929.
^Rinsland, C.P. (1989). "Stratospheric N2O5 profiles at sunrise and sunset from further analysis of the ATMOS/Spacelab 3 solar spectra". Journal of Geophysical Research. 94: 18341–18349. Bibcode:1989JGR....9418341R. doi:10.1029/JD094iD15p18341.
^Kelleher, Patrick J.; Menges, Fabian S.; DePalma, Joseph W.; Denton, Joanna K.; Johnson, Mark A.; Weddle, Gary H.; Hirshberg, Barak; Gerber, R. Benny (2017-09-18). "Trapping and Structural Characterization of the XNO2·NO3− (X = Cl, Br, I) Exit Channel Complexes in the Water-Mediated X− + N2O5 Reactions with Cryogenic Vibrational Spectroscopy". The Journal of Physical Chemistry Letters. 8 (19): 4710–4715. doi:10.1021/acs.jpclett.7b02120. PMID28898581.
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