Dinitrogen pentoxide

Dinitrogen pentoxide is the chemical compound with the formula N2O5, also known as nitrogen pentoxide or nitric anhydride. It is one of the binary nitrogen oxides, a family of compounds that only contain nitrogen and oxygen. It exists as colourless crystals that melt at 41 °C. Its boiling point is 47 °C, and sublimes slightly above room temperature,[1] yielding a colorless gas.[2]

Dinitrogen pentoxide
Full structural formula with dimensions
Ball-and-stick model
Names
IUPAC name
Dinitrogen pentaoxide
Other names
Nitric anhydride
Nitronium nitrate
Nitryl nitrate
DNPO
Anhydrous nitric acid
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.030.227
EC Number
  • 233-264-2
UNII
Properties
N2O5
Molar mass 108.01 g/mol
Appearance white solid
Density 1.642 g/cm3 (18 °C)
Melting point 41 °C (106 °F; 314 K) [1]
Boiling point 47 °C (117 °F; 320 K) sublimes
reacts to give HNO3
Solubility soluble in chloroform
negligible in CCl4
−35.6·10−6 cm3/mol (aq)
1.39 D
Structure
hexagonal
planar, C2v (approx. D2h)
N–O–N ≈ 180°
Thermochemistry
178.2 J K−1 mol−1 (s)
355.6 J K−1 mol−1 (g)
−43.1 kJ/mol (s)
+11.3 kJ/mol (g)
114.1 kJ/mol
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
Nitric acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

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dangerous oxidizer that once was used as a reagent when dissolved in chloroform for nitrations but has largely been superseded by NO2BF4 (nitronium tetrafluoroborate).

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.[3]

HistoryEdit

N2O5 was first reported by Deville in 1840, who prepared it by treating AgNO3 with Cl2.[4][5]

Structure and physical propertiesEdit

Pure solid N2O5 is a salt, consisting of separated linear nitronium ions NO2+ and planar trigonal nitrate anions NO3. 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.[6]

The vapor pressure P (in torr) as a function of temperature T (in kelvin), in the range 211 to 305 K, 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 degrees below the melting point).[7]

In the gas phase, or when dissolved in a nonpolar solvents such as CCl4, the compound exists as covalently bound molecules O2N–O–NO2. 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) [8]

When gaseous N
2
O
5
is cooled rapidly ("quenched"), one can obtain the metastable molecular form, which exothermically converts to the ionic form above −70 °C.[9]

Gaseous N
2
O
5
absorbs ultraviolet light with dissociation into the 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.[10]

PreparationEdit

A recommended laboratory synthesis entails dehydrating nitric acid (HNO3) with phosphorus(V) oxide:[9]

P4O10 + 12 HNO3 → 4 H3PO4 + 6 N2O5

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:[6]

BrF
5
+ 3LiNO
3
→ 3LiF + BrONO
2
+ O2 + 2FNO2
FNO2 + LiNO
3
LiF + N
2
O
5

The compound can also be created in the gas phase by reacting nitrogen dioxide NO
2
or N
2
O
4
with ozone:[11]

2NO
2
+ O
3
N
2
O
5
+ O
2

However, the product catalyzes the rapid decomposition of ozone:[11]

2O
3
+ N
2
O
5
→ 3O
2
+ N
2
O
5

Dinitrogen pentoxide is also formed when a mixture of oxygen and nitrogen is passed through an electric discharge.[6] Another route is the reactions of POCl
3
or NO
2
Cl
with AgNO
3
[6]

ReactionsEdit

Dinitrogen pentoxide reacts with water (hydrolyses) to produce nitric acid HNO
3
. Thus, dinitrogen pentoxide is the anhydride of nitric acid:[9]

N2O5 + H2O → 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
2
O
N
2
O
5
shows the well-known negative azeotrope at 60% N
2
O
5
(that is, 70% HNO
3
), a positive azeotrope at 85.7% N
2
O
5
(100% HNO
3
), and another negative one at 87.5% N
2
O
5
("102% HNO
3
").[12]

The reaction with hydrogen chloride HCl also gives nitric acid and nitryl chloride NO
2
Cl
:[13]

N
2
O
5
+ HClHNO
3
+ NO
2
Cl

Dinitrogen pentoxide eventually decomposes at room temperature into NO2 and O2.[14][11] Decomposition is negligible if the solid is kept at 0 °C, in suitably inert containers.[6]

Dinitrogen pentoxide reacts with ammonia NH
3
to give several products, including nitrous oxide N
2
O
, ammonium nitrate NH
4
NO
3
, nitramide NH
2
NO
2
and ammonium dinitramide NH
4
N(NO
2
)
2
, depending on reaction conditions.[15]

ApplicationsEdit

Nitration of organic compoundsEdit

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.[16] The reactivity of the NO2+ can be further enhanced with strong acids that generate the "super-electrophile" HNO22+.

In this use, N
2
O
5
has been largely replaced by nitronium tetrafluoroborate [NO
2
]+[BF
4
]. This salt retains the high reactivity of NO2+, but it is thermally stable, decomposing at about 180 °C (into NO2F and BF3).

Dinitrogen pentoxide is relevant to the preparation of explosives.[5][17]

Atmospheric occurrenceEdit

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.[18] Mixing ratios of several ppbv have been observed in polluted regions of the night-time troposphere.[19] Dinitrogen pentoxide has also been observed in the stratosphere[20] 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.[21] Two important reactions of N2O5 in atmospheric aerosols are: 1) Hydrolysis to form nitric acid[22] and 2) Reaction with halide ions, particularly Cl, to form ClNO2 molecules which may serve as precursors to reactive chlorine atoms in the atmosphere.[23][24]

HazardsEdit

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.

ReferencesEdit

  1. ^ a b Emeleus (1 January 1964). Advances in Inorganic Chemistry. Academic Press. pp. 77–. ISBN 978-0-12-023606-0. Retrieved 20 September 2011.
  2. ^ Peter Steele Connell The Photochemistry of Dinitrogen Pentoxide. Ph. D. thesis, Lawrence Berkeley National Laboratory.
  3. ^ W. Rogie Angus, Richard W. Jones, and Glyn O. Phillips (1949): "Existence of Nitrosyl Ions (NO+
    ) in Dinitrogen Tetroxide and of Nitronium Ions (NO+
    2
    ) in Liquid Dinitrogen Pentoxide". Nature, volume 164, pages 433–434. doi:10.1038/164433a0
  4. ^ M.H. Deville (1849). "Note sur la production de l'acide nitrique anhydre". Compt. Rend. 28: 257–260.
  5. ^ a b Jai Prakash Agrawal (19 April 2010). High Energy Materials: Propellants, Explosives and Pyrotechnics. Wiley-VCH. pp. 117–. ISBN 978-3-527-32610-5. Retrieved 20 September 2011.
  6. ^ a b c d e William W. Wilson and Karl O. Christe (1987): "Dinitrogen Pentoxide. New Synthesis and Laser Raman Spectrum". Inorganic Chemistry, volume 26, pages 1631-1633. doi:10.1021/ic00257a033
  7. ^ A. H. McDaniel, J. A. Davidson, C. A. Cantrell, R. E. Shetter, and J. G. Calvert (1988): "Enthalpies of formation of dinitrogen pentoxide and the nitrate free radical". Journal of Physical Chemistry, volume 92, issue 14, pages 4172-4175. doi:10.1021/j100325a035
  8. ^ S. Parthiban, B. N. Raghunandan, and R.Sumathi (1996): "Structures, energies and vibrational frequencies of dinitrogen pentoxide". Journal of Molecular Structure: THEOCHEM, volume 367, pages 111-118. doi:10.1016/S0166-1280(96)04516-2
  9. ^ a b c Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils (ed.), Inorganic Chemistry, translated by Eagleson, Mary; Brewer, William, San Diego/Berlin: Academic Press/De Gruyter, ISBN 0-12-352651-5
  10. ^ Bruce A. Osborne, George Marston, L. Kaminski, N. C. Jones, J. M. Gingell, Nigel Mason, Isobel C. Walker, J. Delwiche, and M.-J. Hubin-Franskin (2000): "Vacuum ultraviolet spectrum of dinitrogen pentoxide". Journal of Quantitative Spectroscopy and Radiative Transfer, volume 64, issue 1, pages 67-74. doi:10.1016/S0022-4073(99)00104-1
  11. ^ a b c Francis Yao, Ivan Wilson, and Harold Johnston (1982): "Temperature-dependent ultraviolet absorption spectrum for dinitrogen pentoxide". Journal of Physical Chemistry, volume 86, issue 18, pages 3611-3615. doi:10.1021/j100215a023
  12. ^ L. Lloyd and P. A. H. Wyatt (1955): "The vapour pressures of nitric acid solutions. Part I. New azeotropes in the water–dinitrogen pentoxide system". Journal of the Chemical Society (Resumed), volume 1955, pages 2248-2252.doi:10.1039/JR9550002248
  13. ^ Robert A. Wilkins Jr. and I. C. Hisatsune (1976): "The Reaction of Dinitrogen Pentoxide with Hydrogen Chloride". Industrial & Engineering Chemistry Fundamentals, volume 15, issue 4, pages 246-248. doi:10.1021/i160060a003
  14. ^ Nitrogen(V) Oxide. Inorganic Syntheses. 3. 1950. pp. 78–81.
  15. ^ C. Frenck and W. Weisweiler (2002): "Modeling the Reactions Between Ammonia and Dinitrogen Pentoxide to Synthesize Ammonium Dinitramide (ADN)". Chemical Engineering & Technology, volume 25, issue 2, pages 123-128. doi:10.1002/1521-4125(200202)25:2<123::AID-CEAT123>3.0.CO;2-W
  16. ^ Jan M. Bakke and Ingrd Hegbom (1994): "Dinitrogen pentoxide-sulfur dioxide, a new nitration system". Acta chemica scandinavica, volume 48, issue 2, pages 181-182. doi:10.3891/acta.chem.scand.48-0181
  17. ^ Talawar, M. B.; et al. (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.
  18. ^ 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.
  19. ^ HaiChao Wang; 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.
  20. ^ C.P. Rinsland; et al. (1989). "Stratospheric N205 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.
  21. ^ 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. ISSN 1680-7324.
  22. ^ Brown, S. S.; Dibb, J. E.; Stark, H.; Aldener, M.; Vozella, M.; Whitlow, S.; Williams, E. J.; Lerner, B. M.; Jakoubek, R. (2004-04-16). "Nighttime removal of NOx in the summer marine boundary layer". Geophysical Research Letters. 31 (7): n/a. doi:10.1029/2004GL019412. ISSN 1944-8007.
  23. ^ 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" (PDF). Physical Chemistry Chemical Physics. 17 (29): 19360–19370. Bibcode:2015PCCP...1719360H. doi:10.1039/C5CP02664D. ISSN 1463-9084. PMID 26140681.
  24. ^ 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. ISSN 1948-7185. PMID 28898581.