Chemistry:Dinitrogen pentoxide
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| Names | |
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| IUPAC name
Dinitrogen pentoxide
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| Other names
Nitric anhydride
Nitronium nitrate Nitryl nitrate DNPO Anhydrous nitric acid | |
| Identifiers | |
3D model (JSmol)
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| ChEBI | |
| ChemSpider | |
| EC Number |
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PubChem CID
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| Properties | |
| N2O5 | |
| Molar mass | 108.01 g/mol |
| Appearance | white solid |
| Density | 2.0 g/cm3[1] |
| Boiling point | 33 °C (91 °F; 306 K) sublimes[1] |
| reacts to give HNO3 | |
| Solubility | soluble in chloroform negligible in CCl4 |
| −35.6×10−6 cm3 mol−1 (aq) | |
| 1.39 D | |
| Structure[2] | |
| Hexagonal, hP14 | |
| P63/mmc No. 194 | |
a = 0.54019 nm, c = 0.65268 nm
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Formula units (Z)
|
2 |
| planar, C2v (approx. D2h) N–O–N ≈ 180° | |
| Thermochemistry[3] | |
Heat capacity (C)
|
143.1 J K−1 mol−1 (s) 95.3 J K−1 mol−1 (g) |
Std molar
entropy (S |
178.2 J K−1 mol−1 (s) 355.7 J K−1 mol−1 (g) |
Std enthalpy of
formation (ΔfH⦵298) |
−43.1 kJ/mol (s) +13.3 kJ/mol (g) |
Gibbs free energy (ΔfG˚)
|
113.9 kJ/mol (s) +117.1 kJ/mol (g) |
| Hazards | |
| Main hazards | strong oxidizer, forms strong acid in contact with water |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Related compounds | |
| Nitrous oxide Nitric oxide Dinitrogen trioxide Nitrogen dioxide Dinitrogen tetroxide | |
Related compounds
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Nitric acid |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
 verify (what is   ?)
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| Infobox references | |
Dinitrogen pentoxide (also known as nitrogen pentoxide or nitric anhydride) is the chemical compound with the formula N
2O
5. 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dangerous oxidizer that once was used as a reagent when dissolved in chloroform for nitrations but has largely been superseded by nitronium tetrafluoroborate (NO
2BF
4).
N
2O
5 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 [NO
2]+
and nitrate anions [NO
3]−
; but in the gas phase and under some other conditions it is a covalently-bound molecule.[5]
History
N
2O
5 was first reported by Deville in 1840, who prepared it by treating silver nitrate (AgNO
3) with chlorine.[6][7]
Structure and physical properties
Pure solid N
2O
5 is a salt, consisting of separated linear nitronium ions NO+
2 and planar trigonal nitrate anions NO−
3. Both nitrogen centers have oxidation state +5. It crystallizes in the space group D46h (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 (in kelvin), in the range 211 to 305 K (−62 to 32 °C), is well approximated by the formula
- [math]\displaystyle{ \ln P = 23.2348 - \frac{7098.2}{T} }[/math]
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 O
2N–O–NO
2. In the gas phase, theoretical calculations for the minimum-energy configuration indicate that the O–N–O angle in each –NO
2 wing is about 134° and the N–O–N angle is about 112°. In that configuration, the two –NO
2 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 N
2O
5 is cooled rapidly ("quenched"), one can obtain the metastable molecular form, which exothermically converts to the ionic form above −70 °C.[11]
Gaseous N
2O
5 absorbs ultraviolet light with dissociation into the free radicals nitrogen dioxide NO
2•
and nitrogen trioxide NO
3•
(uncharged nitrate). The absorption spectrum has a broad band with maximum at wavelength 160 nm.[12]
Preparation
A recommended laboratory synthesis entails dehydrating nitric acid (HNO
3) with phosphorus(V) oxide:[11]
- P
4O
10 + 12 HNO
3 → 4 H
3PO
4 + 6 N
2O
5
Another laboratory process is the reaction of lithium nitrate LiNO
3 and bromine pentafluoride BrF
5, in the ratio exceeding 3:1. The reaction first forms nitryl fluoride FNO
2 that reacts further with the lithium nitrate:[8]
- BrF
5 + 3 LiNO
3 → 3 LiF + BrONO
2 + O
2 + 2 FNO
2 - FNO
2 + LiNO
3 → LiF + N
2O
5
The compound can also be created in the gas phase by reacting nitrogen dioxide NO
2 or N
2O
4 with ozone:[13]
- 2 NO
2 + O
3 → N
2O
5 + O
2
However, the product catalyzes the rapid decomposition of ozone:[13]
- 2 O
3 + N
2O
5 → 3 O
2 + N
2O
5
Dinitrogen pentoxide is also formed when a mixture of oxygen and nitrogen is passed through an electric
discharge.[8] Another route is the reactions of Phosphoryl chloride POCl
3 or nitryl chloride NO
2Cl with silver nitrate AgNO
3[8][14]
Reactions
Dinitrogen pentoxide reacts with water (hydrolyses) to produce nitric acid HNO
3. Thus, dinitrogen pentoxide is the anhydride of nitric acid:[11]
- N
2O
5 + H
2O → 2 HNO
3
Solutions of dinitrogen pentoxide in nitric acid can be seen as nitric acid with more than 100% concentration. The phase diagram of the system H
2O−N
2O
5 shows the well-known negative azeotrope at 60% N
2O
5 (that is, 70% HNO
3), a positive azeotrope at 85.7% N
2O
5 (100% HNO
3), and another negative one at 87.5% N
2O
5 ("102% HNO
3").[15]
The reaction with hydrogen chloride HCl also gives nitric acid and nitryl chloride NO
2Cl:[16]
- N
2O
5 + HCl → HNO
3 + NO
2Cl
Dinitrogen pentoxide eventually decomposes at room temperature into NO
2 and O
2.[17][13] Decomposition is negligible if the solid is kept at 0 °C, in suitably inert containers.[8]
Dinitrogen pentoxide reacts with ammonia NH
3 to give several products, including nitrous oxide N
2O, ammonium nitrate NH
4NO
3, nitramide NH
2NO
2 and ammonium dinitramide NH
4N(NO
2)
2, depending on reaction conditions.[18]
Decomposition of dinitrogen pentoxide at high temperatures
Dinitrogen pentoxide between high temperatures of 600 and 1,100 K (327–827 °C), is decomposed in two successive stoichiometric steps:
- N
2O
5 → NO
2 + NO
3 - 2 NO
3 → 2 NO
2 + O
2
In the shock wave, N
2O
5 has decomposed stoichiometrically into nitrogen dioxide and oxygen. At temperatures of 600 K and higher, nitrogen dioxide is unstable with respect to nitrogen oxide NO 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
Apart from the decomposition of N
2O
5 at high temperatures, it can also be decomposed in carbon tetrachloride CCl
4 at 30 °C (303 K).[20] Both N
2O
5 and NO
2 are soluble in CCl
4 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 O
2 that is produced over time because the only form to obtain O
2 is with the N
2O
5 decomposition. The equation below refers to the decomposition of N
2O
5 in CCl
4:
- 2 N
2O
5 → 4 NO
2 + O
2(g)
And this reaction follows the first order rate law that says:
- [math]\displaystyle{ -\frac{d[\mathrm{A}]}{dt} = k [\mathrm{A}] }[/math]
Decomposition of nitrogen pentoxide in the presence of nitric oxide
N
2O
5 can also be decomposed in the presence of nitric oxide NO:
- N
2O
5 + NO → 3 NO
2
The rate of the initial reaction between dinitrogen pentoxide and nitric oxide of the elementary unimolecular decomposition.[21]
Applications
Nitration of organic compounds
Dinitrogen pentoxide, for example as a solution in chloroform, has been used as a reagent to introduce the –NO
2 functionality in organic compounds. This nitration reaction is represented as follows:
- N
2O
5 + Ar–H → HNO
3 + Ar–NO
2
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, N
2O
5 has been largely replaced by nitronium tetrafluoroborate [NO
2]+
[BF
4]−
. This salt retains the high reactivity of NO+
2, but it is thermally stable, decomposing at about 180 °C (into NO
2F and BF
3).
Dinitrogen pentoxide is relevant to the preparation of explosives.[7][23]
Atmospheric occurrence
In the atmosphere, dinitrogen pentoxide is an important reservoir of the NO
x species that are responsible for ozone depletion: its formation provides a null cycle with which NO and NO
2 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 NO
2 levels above 50 °N, the so-called 'Noxon cliff'.
Variations in N
2O
5 reactivity in aerosols can result in significant losses in tropospheric ozone, hydroxyl radicals, and NO
x concentrations.[27] Two important reactions of N
2O
5 in atmospheric aerosols are hydrolysis to form nitric acid[28] and reaction with halide ions, particularly Cl−
, to form ClNO
2 molecules which may serve as precursors to reactive chlorine atoms in the atmosphere.[29][30]
Hazards
N
2O
5 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.
References
- ↑ 1.0 1.1 Haynes, p. 4.76
- ↑ Simon, Arndt; Horakh, Jörg; Obermeyer, Axel; Borrmann, Horst (1992). "Kristalline Stickstoffoxide — Struktur von N2O3 mit einer Anmerkung zur Struktur von N2O5" (in de). Angewandte Chemie (Wiley) 104 (3): 325–327. doi:10.1002/ange.19921040321. Bibcode: 1992AngCh.104..325S.
- ↑ Haynes, p. 5.29
- ↑ Connell, Peter Steele. (1979) The Photochemistry of Dinitrogen Pentoxide. Ph. D. thesis, Lawrence Berkeley National Laboratory.
- ↑ Angus, W.R.; Jones, R.W.; Phillips, G.O. (1949). "Existence of Nitrosyl Ions (NO+) in Dinitrogen Tetroxide and of Nitronium Ions (NO2+) in Liquid Dinitrogen Pentoxide". Nature 164 (4167): 433. doi:10.1038/164433a0. PMID 18140439. Bibcode: 1949Natur.164..433A.
- ↑ Deville, M.H. (1849). "Note sur la production de l'acide nitrique anhydre". Compt. Rend. 28: 257–260. https://archive.org/details/comptesrendusheb28acad/page/257.
- ↑ 7.0 7.1 Agrawal, Jai Prakash (2010). High Energy Materials: Propellants, Explosives and Pyrotechnics. Wiley-VCH. p. 117. ISBN 978-3-527-32610-5. https://books.google.com/books?id=rqZROysoS7QC&pg=PA117. Retrieved 20 September 2011.
- ↑ 8.0 8.1 8.2 8.3 8.4 Wilson, 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.
- ↑ 11.0 11.1 11.2 Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils, ed., Inorganic Chemistry, San Diego/Berlin: Academic Press/De Gruyter, ISBN 0-12-352651-5
- ↑ Osborne, Bruce A.; Marston, George; Kaminski, L.; Jones, N.C; Gingell, J.M; Mason, Nigel; Walker, Isobel C.; Delwiche, J. et al. (2000). "Vacuum ultraviolet spectrum of dinitrogen pentoxide". Journal of Quantitative Spectroscopy and Radiative Transfer 64 (1): 67–74. doi:10.1016/S0022-4073(99)00104-1. Bibcode: 2000JQSRT..64...67O.
- ↑ 13.0 13.1 13.2 Yao, 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. ISBN 9780470132340.
- ↑ 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.
- ↑ Jaime, R. (2008). Determinación de orden de reacción haciendo uso de integrales definidas. Universidad Nacional Autónoma de Nicaragua, Managua.
- ↑ 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.
- ↑ Bakke, Jan M.; Hegbom, Ingrid; Verne, Hans Peter; Weidlein, Johann; Schnöckel, Hansgeorg; Paulsen, Gudrun B.; Nielsen, Ruby I.; Olsen, Carl E. et al. (1994). "Dinitrogen Pentoxide--Sulfur Dioxide, a New Nitration System". Acta Chemica Scandinavica 48: 181–182. doi:10.3891/acta.chem.scand.48-0181.
- ↑ 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. PMID 15979786.
- ↑ Finlayson-Pitts, Barbara J.; Pitts, James N. (2000). Chemistry of the upper and lower atmosphere : theory, experiments, and applications. San Diego: Academic Press. ISBN 9780080529073. OCLC 162128929.
- ↑ Wang, Haichao; Lu, Keding; Chen, Xiaorui; Zhu, Qindan; Chen, Qi; Guo, Song; Jiang, Meiqing; Li, Xin et al. (2017). "High N2O5 Concentrations Observed in Urban Beijing: Implications of a Large Nitrate Formation Pathway". Environmental Science and Technology Letters 4 (10): 416–420. doi:10.1021/acs.estlett.7b00341.
- ↑ 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. doi:10.1029/JD094iD15p18341. Bibcode: 1989JGR....9418341R.
- ↑ Macintyre, H. L.; Evans, M. J. (2010-08-09). "Sensitivity of a global model to the uptake of N2O5 by tropospheric aerosol". Atmospheric Chemistry and Physics 10 (15): 7409–7414. doi:10.5194/acp-10-7409-2010. Bibcode: 2010ACP....10.7409M.
- ↑ Brown, S. S.; Dibb, J. E.; Stark, H.; Aldener, M.; Vozella, M.; Whitlow, S.; Williams, E. J.; Lerner, B. M. et al. (2004-04-16). "Nighttime removal of NOx in the summer marine boundary layer" (in en). Geophysical Research Letters 31 (7): n/a. doi:10.1029/2004GL019412. Bibcode: 2004GeoRL..31.7108B.
- ↑ Gerber, R. Benny; Finlayson-Pitts, Barbara J.; Hammerich, Audrey Dell (2015-07-15). "Mechanism for formation of atmospheric Cl atom precursors in the reaction of dinitrogen oxides with HCl/Cl− on aqueous films" (in en). Physical Chemistry Chemical Physics 17 (29): 19360–19370. doi:10.1039/C5CP02664D. PMID 26140681. Bibcode: 2015PCCP...1719360H. https://escholarship.org/content/qt3087m4xv/qt3087m4xv.pdf?t=oubfuu.
- ↑ 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. PMID 28898581.
Cited sources
- Haynes, William M., ed (2016). CRC Handbook of Chemistry and Physics (97th ed.). CRC Press. ISBN 9781498754293.
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