Dinitrogen pentoxide

{{chembox | Verifiedfields = changed | Watchedfields = changed | verifiedrevid = 451483489 | Name = | ImageFile = Dinitrogen pentoxide.svg | ImageClass = skin-invert-image | ImageSize = 200px | ImageName = Full structural formula with dimensions | ImageFile1 = Dinitrogen-pentoxide-3D-balls.png | ImageClass1 = bg-transparent | ImageSize1 = 200px | ImageName1 = Ball-and-stick model | IUPACName = Dinitrogen pentoxide | OtherNames = Nitric anhydride Nitronium nitrate Nitryl nitrate DNPO Anhydrous nitric acid | Section1 = {{Chembox Identifiers | ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}} | ChemSpiderID = 59627 | InChI = 1/N2O5/c3-1(4)7-2(5)6 | ChEBI_Ref = {{ebicite|correct|EBI}} | ChEBI = 29802 | SMILES = [O-][N+](=O)O[N+]([O-])=O | SMILES_Comment = gas phase | SMILES1 = [O]=[N+]=[O].[N+](=O)([O-])[O-] | SMILES1_Comment = solid phase | InChIKey = ZWWCURLKEXEFQT-UHFFFAOYAN | StdInChI_Ref = {{stdinchicite|correct|chemspider}} | StdInChI = 1S/N2O5/c3-1(4)7-2(5)6 | StdInChIKey_Ref = {{stdinchicite|correct|chemspider}} | StdInChIKey = ZWWCURLKEXEFQT-UHFFFAOYSA-N | CASNo = 10102-03-1 | CASNo_Ref = {{cascite|correct|CAS}} | UNII_Ref = {{fdacite|changed|FDA}} | UNII = 6XB659ZX2W | PubChem = 66242 | EINECS = 233-264-2 }} | Section2 = {{Chembox Properties | Formula = N2O5 | MolarMass = 108.01 g/mol | Appearance = white solid | Density = 2.0 g/cm3<ref name=r1>Haynes, p. 4.76</ref> | Solubility = reacts to give HNO3 | SolubleOther = soluble in chloroform negligible in CCl4 | BoilingPtC = 33 | BoilingPt_notes = sublimes<ref name=r1/> | Dipole = 1.39 D | MagSus = {{val|-35.6e-6|u=cm3 mol−1}} (aq) }} | Section3 = {{Chembox Structure | Structure_ref =<ref>{{cite journal | last1=Simon | first1=Arndt | last2=Horakh | first2=Jörg | last3=Obermeyer | first3=Axel | last4=Borrmann | first4=Horst | title=Kristalline Stickstoffoxide — Struktur von N2O3 mit einer Anmerkung zur Struktur von N2O5 | journal=Angewandte Chemie | publisher=Wiley | volume=104 | issue=3 | year=1992 | doi=10.1002/ange.19921040321 | pages=325–327 | bibcode=1992AngCh.104..325S | language=de}}</ref> | CrystalStruct = Hexagonal, hP14 | SpaceGroup = P63/mmc No. 194 | LattConst_a = 0.54019 nm | LattConst_c = 0.65268 nm | UnitCellFormulas =2 | MolShape = planar, ''C''2v (approx. ''D''2h) N–O–N ≈ 180° }} | Section4 = {{Chembox Thermochemistry | Thermochemistry_ref = <ref>Haynes, p. 5.29</ref> | DeltaHf = −43.1 kJ/mol (s) +13.3 kJ/mol (g) | DeltaGf = 113.9 kJ/mol (s) +117.1 kJ/mol (g) | Entropy = 178.2 J K−1 mol−1 (s) 355.7 J K−1 mol−1 (g) | HeatCapacity = 143.1 J K−1 mol−1 (s) 95.3 J K−1 mol−1 (g) }} | Section7 = {{Chembox Hazards | MainHazards = strong oxidizer, forms strong acid in contact with water | NFPA-H = 4 | NFPA-F = 0 | NFPA-R = 2 | NFPA-S = W OX | FlashPt = Non-flammable }} | Section8 = {{Chembox Related | OtherFunction = Nitrous oxide Nitric oxide Dinitrogen trioxide Nitrogen dioxide Dinitrogen tetroxide | OtherFunction_label = nitrogen oxides | OtherCompounds = Nitric acid }} }} '''Dinitrogen pentoxide''' (also known as '''nitrogen pentoxide''' or '''nitric anhydride''') is the chemical compound with the formula '''{{chem2|N2O5}}'''. It is one of the binary nitrogen oxides, a family of compounds that contain only nitrogen and oxygen. It exists as colourless crystals that sublime slightly above room temperature, yielding a colorless gas.<ref name=conn1979>Connell, Peter Steele. (1979) ''[https://escholarship.org/uc/item/5dn1m436 The Photochemistry of Dinitrogen Pentoxide]''. Ph. D. thesis, Lawrence Berkeley National Laboratory.</ref>

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 ({{chem2|NO2BF4}}).

{{chem2|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 {{chem2|[NO2]+}} and nitrate anions {{chem2|[NO3]-}}; but in the gas phase and under some other conditions it is a covalently-bound molecule.<ref name=angus1949>{{cite journal|title=Existence of Nitrosyl Ions (NO+) in Dinitrogen Tetroxide and of Nitronium Ions (NO2+) in Liquid Dinitrogen Pentoxide|year=1949 |doi=10.1038/164433a0|pmid=18140439 |author1=Angus, W.R. |author2=Jones, R.W. |author3=Phillips, G.O. |journal=Nature |volume=164 |issue=4167 |page=433 |bibcode=1949Natur.164..433A |s2cid=4136455 }}</ref>

==History== {{chem2|N2O5}} was first reported by the French chemist Henri Deville in 1840, who prepared it by treating silver nitrate ({{chem2|AgNO3}}) with chlorine.<ref>{{cite journal|author=Deville, M.H. |journal= Compt. Rend. |volume=28 |year=1849|title=Note sur la production de l'acide nitrique anhydre|url=https://archive.org/details/comptesrendusheb28acad/page/257|pages=257–260}}</ref><ref name=b1>{{cite book|author=Agrawal, Jai Prakash |title=High Energy Materials: Propellants, Explosives and Pyrotechnics|url=https://books.google.com/books?id=rqZROysoS7QC&pg=PA117|access-date=20 September 2011|date=2010|publisher=Wiley-VCH|isbn=978-3-527-32610-5|page=117}}</ref>

==Structure and physical properties== Pure solid {{chem2|N2O5}} is a salt, consisting of separated linear nitronium ions {{chem2|NO2+}} and planar trigonal nitrate anions {{chem2|NO3-}}. Both nitrogen centers have oxidation state +5. It crystallizes in the space group ''D''{{su|p=4|b=6''h''}} (''C''6/''mmc'') with ''Z'' = 2, with the {{chem2|NO3-}} anions in the ''D''3''h'' sites and the {{chem2|NO2+}} cations in ''D''3''d'' sites.<ref name=wils1982/>

The vapor pressure ''P'' (in atm) as a function of temperature ''T'' (in kelvin), in the range {{cvt|211|to|305|K|C}}, is well approximated by the formula :<math> \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).<ref name=mcda1988>{{cite journal|doi=10.1021/j100325a035|title=Enthalpies of formation of dinitrogen pentoxide and the nitrate free radical |year=1988 |last1=McDaniel |first1=A. H. |last2=Davidson |first2=J. A. |last3=Cantrell |first3=C. A. |last4=Shetter |first4=R. E. |last5=Calvert |first5=J. G. |journal=The Journal of Physical Chemistry |volume=92 |issue=14 |pages=4172–4175 }}</ref>

In the gas phase, or when dissolved in nonpolar solvents such as carbon tetrachloride, the compound exists as covalently-bonded molecules {{chem2|O2N\sO\sNO2}}. In the gas phase, theoretical calculations for the minimum-energy configuration indicate that the {{chem2|O\sN\sO}} angle in each {{chem2|\sNO2}} wing is about 134° and the {{chem2|N\sO\sN}} angle is about 112°. In that configuration, the two {{chem2|\sNO2}} groups are rotated about 35° around the bonds to the central oxygen, away from the {{chem2|N\sO\sN}} plane. The molecule thus has a propeller shape, with one axis of 180° rotational symmetry (''C''2) <ref name=parth1996>{{cite journal|doi=10.1016/S0166-1280(96)04516-2|title=Structures, energies and vibrational frequencies of dinitrogen pentoxide |year=1996 |last1=Parthiban |first1=S. |last2=Raghunandan |first2=B.N. |last3=Sumathi |first3=R. |journal=Journal of Molecular Structure: Theochem |volume=367 |pages=111–118}}</ref>

When gaseous {{chem2|N2O5}} is cooled rapidly ("quenched"), one can obtain the metastable molecular form, which exothermically converts to the ionic form above −70 °C.<ref name=Holl/>

Gaseous {{chem2|N2O5}} absorbs ultraviolet light with dissociation into the free radicals nitrogen dioxide {{chem2|NO2^{•}|}} and nitrogen trioxide {{chem2|NO3^{•}|}} (uncharged nitrate). The absorption spectrum has a broad band with maximum at wavelength 160 nm.<ref name=osbo2000>{{cite journal|doi=10.1016/S0022-4073(99)00104-1|title=Vacuum ultraviolet spectrum of dinitrogen pentoxide |year=2000 |last1=Osborne |first1=Bruce A. |last2=Marston |first2=George |last3=Kaminski |first3=L. |last4=Jones |first4=N.C |last5=Gingell |first5=J.M |last6=Mason |first6=Nigel |last7=Walker |first7=Isobel C. |last8=Delwiche |first8=J. |last9=Hubin-Franskin |first9=M.-J. |journal=Journal of Quantitative Spectroscopy and Radiative Transfer |volume=64 |issue=1 |pages=67–74 |bibcode=2000JQSRT..64...67O}}</ref>

==Preparation== A recommended laboratory synthesis entails dehydrating nitric acid ({{chem2|HNO3}}) with phosphorus(V) oxide:<ref name=Holl>{{Holleman&Wiberg}}</ref> :{{chem2|P4O10 + 12 HNO3 → 4 H3PO4 + 6 N2O5}}

Another laboratory process is the reaction of lithium nitrate {{chem2|LiNO3}} and bromine pentafluoride {{chem2|BrF5}}, in the ratio exceeding 3:1. The reaction first forms nitryl fluoride {{chem2|FNO2}} that reacts further with the lithium nitrate:<ref name=wils1982/> : {{chem2|BrF5 + 3 LiNO3 → 3 LiF + BrONO2 + O2 + 2 FNO2}} : {{chem2|FNO2 + LiNO3 → LiF + N2O5}}

The compound can also be created in the gas phase by reacting nitrogen dioxide {{chem2|NO2}} or {{chem2|N2O4}} with ozone:<ref name=yao1982>{{cite journal|doi=10.1021/j100215a023|title=Temperature-dependent ultraviolet absorption spectrum for dinitrogen pentoxide |year=1982 |last1=Yao |first1=Francis |last2=Wilson |first2=Ivan |last3=Johnston |first3=Harold |journal=The Journal of Physical Chemistry |volume=86 |issue=18 |pages=3611–3615 }}</ref> : {{chem2|2 NO2 + O3 → N2O5 + O2}} However, the product catalyzes the rapid decomposition of ozone:<ref name=yao1982/> : {{chem2|2 O3 + N2O5 → 3 O2 + N2O5}}

Dinitrogen pentoxide is also formed when a mixture of oxygen and nitrogen is passed through an electric discharge.<ref name=wils1982>{{cite journal|doi=10.1021/ic00257a033|title=Dinitrogen pentoxide. New synthesis and laser Raman spectrum |year=1987 |last1=Wilson |first1=William W. |last2=Christe |first2=Karl O. |journal=Inorganic Chemistry |volume=26 |issue=10 |pages=1631–1633 }}</ref> Another route is the reactions of Phosphoryl chloride {{chem2|POCl3}} or nitryl chloride {{chem2|NO2Cl}} with silver nitrate {{chem2|AgNO3}}<ref name=wils1982/><ref>{{cite journal|doi=10.1021/ja01541a019|title=Shock Waves in Chemical Kinetics: The Decomposition of N2O5 at High Temperatures|year=1958 |last1=Schott |first1=Garry |last2=Davidson |first2=Norman |journal=Journal of the American Chemical Society |volume=80 |issue=8 |pages=1841–1853 |bibcode=1958JAChS..80.1841S }}</ref>

==Reactions== Dinitrogen pentoxide reacts with water (hydrolyses) to produce nitric acid {{chem2|HNO3}}. Thus, dinitrogen pentoxide is the anhydride of nitric acid:<ref name=Holl/> :{{chem2|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 {{chem2|H2O}}−{{chem2|N2O5}} shows the well-known negative azeotrope at 60% {{chem2|N2O5}} (that is, 70% {{chem2|HNO3}}), a positive azeotrope at 85.7% {{chem2|N2O5}} (100% {{chem2|HNO3}}), and another negative one at 87.5% {{chem2|N2O5}} ("102% {{chem2|HNO3}}").<ref name=lloyd1955>{{cite journal|doi=10.1039/JR9550002248|title=The vapour pressures of nitric acid solutions. Part I. New azeotropes in the water–dinitrogen pentoxide system |year=1955 |last1=Lloyd |first1=L. |last2=Wyatt |first2=P. A. H. |journal=J. Chem. Soc. |pages=2248–2252 }}</ref>

The reaction with hydrogen chloride {{chem2|HCl}} also gives nitric acid and nitryl chloride {{chem2|NO2Cl}}:<ref name=wilk1976>{{cite journal|doi=10.1021/i160060a003|title=The Reaction of Dinitrogen Pentoxide with Hydrogen Chloride |year=1976 |last1=Wilkins |first1=Robert A. |last2=Hisatsune |first2=I. C. |journal=Industrial & Engineering Chemistry Fundamentals |volume=15 |issue=4 |pages=246–248 }}</ref> : {{chem2|N2O5 + HCl → HNO3 + NO2Cl}}

Dinitrogen pentoxide eventually decomposes at room temperature into {{chem2|NO2}} and {{chem2|O2}}.<ref>{{cite book | last1=Gruenhut | first1=N. S. | last2=Goldfrank | first2=M. | last3=Cushing | first3=M. L. | last4=Caesar | first4=G. V. | last5=Caesar | first5=P. D. | last6=Shoemaker | first6=C. | title=Inorganic Syntheses | chapter=Nitrogen(V) Oxide (Nitrogen Pentoxide, Dinitrogen Pentoxide, Nitric Anhydride)| date=1950| doi=10.1002/9780470132340.ch20 | pages=78–81| isbn=9780470132340 }}</ref><ref name=yao1982/> Decomposition is negligible if the solid is kept at 0 °C, in suitably inert containers.<ref name=wils1982/>

Dinitrogen pentoxide reacts with ammonia {{chem2|NH3}} to give several products, including nitrous oxide {{chem2|N2O}}, ammonium nitrate {{chem2|NH4NO3}}, nitramide {{chem2|NH2NO2}} and ammonium dinitramide {{chem2|NH4N(NO2)2}}, depending on reaction conditions.<ref name=fren2002>{{cite journal|doi=10.1002/1521-4125(200202)25:2<123::AID-CEAT123>3.0.CO;2-W|title=Modeling the Reactions Between Ammonia and Dinitrogen Pentoxide to Synthesize Ammonium Dinitramide (ADN) |year=2002 |last1=Frenck |first1=C. |last2=Weisweiler |first2=W. |journal=Chemical Engineering & Technology |volume=25 |issue=2 |page=123 }}</ref>

===Decomposition of dinitrogen pentoxide at high temperatures===

Dinitrogen pentoxide between high temperatures of {{cvt|600|and(-)|1100|K|C}}, is decomposed in two successive stoichiometric steps: : {{chem2|N2O5 → NO2 + NO3}} : {{chem2|2 NO3 → 2 NO2 + O2}} In the shock wave, {{chem2|N2O5}} has decomposed stoichiometrically into nitrogen dioxide and oxygen. At temperatures of 600 K and higher, nitrogen dioxide is unstable with respect to nitrogen oxide {{chem|NO}} and oxygen. The thermal decomposition of 0.1 mM nitrogen dioxide at 1000 K is known to require about two seconds.<ref name=garry>{{cite journal|doi=10.1021/ja01541a019|title=Shock Waves in Chemical Kinetics: The Decomposition of N2O5 at High Temperatures|year=1958 |last1=Schott |first1=Garry |last2=Davidson |first2=Norman |journal=Journal of the American Chemical Society |volume=80 |issue=8 |pages=1841–1853 |bibcode=1958JAChS..80.1841S }}</ref>

===Decomposition of dinitrogen pentoxide in carbon tetrachloride at 30 °C===

Apart from the decomposition of {{chem2|N2O5}} at high temperatures, it can also be decomposed in carbon tetrachloride {{chem2|CCl4}} at {{cvt|30|C|K}}.<ref name=Herrera>Jaime, R. (2008). [https://www.researchgate.net/publication/343920548 Determinación de orden de reacción haciendo uso de integrales definidas]. Universidad Nacional Autónoma de Nicaragua, Managua.</ref> Both {{chem2|N2O5}} and {{chem2|NO2}} are soluble in {{chem2|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 {{chem2|O2}} that is produced over time because the only form to obtain {{chem2|O2}} is with the {{chem2|N2O5}} decomposition. The equation below refers to the decomposition of {{chem2|N2O5}} in {{chem2|CCl4}}:

: {{chem2|2 N2O5 → 4 NO2 + O2(g)}}

And this reaction follows the first order rate law that says:

:<math>-\frac{d[\mathrm{A}]}{dt} = k [\mathrm{A}]</math>

===Decomposition of nitrogen pentoxide in the presence of nitric oxide===

{{chem2|N2O5}} can also be decomposed in the presence of nitric oxide {{chem2|NO}}:

: {{chem2|N2O5 + NO → 3 NO2}}

The rate of the initial reaction between dinitrogen pentoxide and nitric oxide of the elementary unimolecular decomposition.<ref>{{cite journal |title=Decomposition of Nitrogen Pentoxide in the Presence of Nitric Oxide. IV. Effect of Noble Gases|journal=Journal of the American Chemical Society|year=1953 |volume=75 |issue=22 |page=5763 |doi=10.1021/ja01118a529 |last1=Wilson |first1=David J. |last2=Johnston |first2=Harold S. |bibcode=1953JAChS..75.5763W }}</ref>

==Applications==

===Nitration of organic compounds=== Dinitrogen pentoxide, for example as a solution in chloroform, has been used as a reagent to introduce the {{chem2|\sNO2}} functionality in organic compounds. This nitration reaction is represented as follows:

:{{chem2|N2O5 + Ar\sH → HNO3 + Ar\sNO2}}

where Ar represents an arene moiety.<ref>{{cite journal|doi=10.3891/acta.chem.scand.48-0181|title=Dinitrogen Pentoxide--Sulfur Dioxide, a New Nitration System |year=1994 |last1=Bakke |first1=Jan M. |last2=Hegbom |first2=Ingrid |last3=Verne |first3=Hans Peter |last4=Weidlein |first4=Johann |last5=Schnöckel |first5=Hansgeorg |last6=Paulsen |first6=Gudrun B. |last7=Nielsen |first7=Ruby I. |last8=Olsen |first8=Carl E. |last9=Pedersen |first9=Christian |last10=Stidsen |first10=Carsten E. |journal=Acta Chemica Scandinavica |volume=48 |pages=181–182 |doi-access=free }}</ref> The reactivity of the {{chem2|NO2+}} can be further enhanced with strong acids that generate the "super-electrophile" {{chem2|HNO2(2+)}}.

In this use, {{chem2|N2O5}} has been largely replaced by nitronium tetrafluoroborate {{chem2|[NO2]+[BF4]-}}. This salt retains the high reactivity of {{chem2|NO2+}}, but it is thermally stable, decomposing at about 180 °C (into {{chem2|NO2F}} and {{chem2|BF3}}).

Dinitrogen pentoxide is relevant to the preparation of explosives.<ref name=b1/><ref>{{cite journal|author=Talawar, M. B.|title=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=Journal of Hazardous Materials|year= 2005| volume =124|issue=1–3| pages =153–64|doi=10.1016/j.jhazmat.2005.04.021|pmid=15979786|bibcode=2005JHzM..124..153T }}</ref>

==Atmospheric occurrence== In the atmosphere, dinitrogen pentoxide is an important reservoir of the {{chem2|NO_{''x''}|}} species that are responsible for ozone depletion: its formation provides a null cycle with which {{chem2|NO}} and {{chem2|NO2}} are temporarily held in an unreactive state.<ref>{{Cite book|title=Chemistry of the upper and lower atmosphere : theory, experiments, and applications|last1=Finlayson-Pitts|first1=Barbara J.|last2=Pitts|first2=James N.|date=2000|publisher=Academic Press|isbn=9780080529073|location=San Diego|oclc=162128929}}</ref> Mixing ratios of several parts per billion by volume have been observed in polluted regions of the nighttime troposphere.<ref>{{cite journal|title=High N2O5 Concentrations Observed in Urban Beijing: Implications of a Large Nitrate Formation Pathway|journal=Environmental Science and Technology Letters|volume=4|issue=10|pages=416–420|year= 2017|doi=10.1021/acs.estlett.7b00341|last1=Wang |first1=Haichao |last2=Lu |first2=Keding |last3=Chen |first3=Xiaorui |last4=Zhu |first4=Qindan |last5=Chen |first5=Qi |last6=Guo |first6=Song |last7=Jiang |first7=Meiqing |last8=Li |first8=Xin |last9=Shang |first9=Dongjie |last10=Tan |first10=Zhaofeng |last11=Wu |first11=Yusheng |last12=Wu |first12=Zhijun |last13=Zou |first13=Qi |last14=Zheng |first14=Yan |last15=Zeng |first15=Limin |last16=Zhu |first16=Tong |last17=Hu |first17=Min |last18=Zhang |first18=Yuanhang |bibcode=2017EnSTL...4..416W }}</ref> Dinitrogen pentoxide has also been observed in the stratosphere<ref>{{cite journal|author=Rinsland, C.P. |title=Stratospheric N2O5 profiles at sunrise and sunset from further analysis of the ''ATMOS/Spacelab 3'' solar spectra|journal=Journal of Geophysical Research|year= 1989|volume =94| pages =18341–18349|doi=10.1029/JD094iD15p18341|bibcode=1989JGR....9418341R}}</ref> at similar levels, the reservoir formation having been postulated in considering the puzzling observations of a sudden drop in stratospheric {{chem2|NO2}} levels above 50 °N, the so-called 'Noxon cliff'.

Variations in {{chem2|N2O5}} reactivity in aerosols can result in significant losses in tropospheric ozone, hydroxyl radicals, and {{chem2|NO_{''x''}|}} concentrations.<ref>{{Cite journal|last1=Macintyre|first1=H. L.|last2=Evans|first2=M. J.|date=2010-08-09|title=Sensitivity of a global model to the uptake of N2O5 by tropospheric aerosol|journal=Atmospheric Chemistry and Physics|volume=10|issue=15|pages=7409–7414|doi=10.5194/acp-10-7409-2010|bibcode=2010ACP....10.7409M|doi-access=free}}</ref> Two important reactions of {{chem2|N2O5}} in atmospheric aerosols are hydrolysis to form nitric acid<ref>{{Cite journal|last1=Brown|first1=S. S.|last2=Dibb|first2=J. E.|last3=Stark|first3=H.|last4=Aldener|first4=M.|last5=Vozella|first5=M.|last6=Whitlow|first6=S.|last7=Williams|first7=E. J.|last8=Lerner|first8=B. M.|last9=Jakoubek|first9=R.|date=2004-04-16|title=Nighttime removal of NO''x'' in the summer marine boundary layer|journal=Geophysical Research Letters|language=en|volume=31|issue=7|pages=n/a|doi=10.1029/2004GL019412|bibcode=2004GeoRL..31.7108B|doi-access=free}}</ref> and reaction with halide ions, particularly {{chem2|Cl-}}, to form {{chem2|ClNO2}} molecules which may serve as precursors to reactive chlorine atoms in the atmosphere.<ref>{{Cite journal|last1=Gerber|first1=R. Benny|last2=Finlayson-Pitts|first2=Barbara J.|last3=Hammerich|first3=Audrey Dell|date=2015-07-15|title=Mechanism for formation of atmospheric Cl atom precursors in the reaction of dinitrogen oxides with HCl/Cl− on aqueous films|journal=Physical Chemistry Chemical Physics|language=en|volume=17|issue=29|pages=19360–19370|doi=10.1039/C5CP02664D|pmid=26140681|bibcode=2015PCCP...1719360H|s2cid=39157816 |url=https://escholarship.org/content/qt3087m4xv/qt3087m4xv.pdf?t=oubfuu}}</ref><ref>{{Cite journal|last1=Kelleher|first1=Patrick J.|last2=Menges|first2=Fabian S.|last3=DePalma|first3=Joseph W.|last4=Denton|first4=Joanna K.|last5=Johnson|first5=Mark A.|last6=Weddle|first6=Gary H.|last7=Hirshberg|first7=Barak|last8=Gerber|first8=R. Benny|date=2017-09-18|title=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|journal=The Journal of Physical Chemistry Letters|volume=8|issue=19|pages=4710–4715|doi=10.1021/acs.jpclett.7b02120|pmid=28898581}}</ref>

==Hazards== {{chem2|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.

==References== {{reflist}}

==Cited sources== *{{cite book | editor= Haynes, William M. | date = 2016| title = CRC Handbook of Chemistry and Physics | edition = 97th | publisher = CRC Press | isbn = 9781498754293}}

{{DEFAULTSORT:Dinitrogen Pentoxide}} Category:Nitrogen oxides Category:Acid anhydrides Category:Acidic oxides Category:Nitrates Category:Nitronium compounds