Dinitrogen pentoxide
Full structural formula with dimensions
Ball-and-stick model
IUPAC name
Dinitrogen pentoxide
Other names
Nitric anhydride
Nitronium nitrate
Nitryl nitrate
Anhydrous nitric acid
3D model (JSmol)
ECHA InfoCard 100.030.227 Edit this at Wikidata
EC Number
  • 233-264-2
  • InChI=1S/N2O5/c3-1(4)7-2(5)6 checkY
  • InChI=1/N2O5/c3-1(4)7-2(5)6
  • [O-][N+](=O)O[N+]([O-])=O
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−1 (aq)
1.39 D
planar, C2v (approx. D2h)
N–O–N ≈ 180°
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
Occupational safety and health (OHS/OSH):
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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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 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dangerous oxidizer that once was used as a reagent when dissolved in chloroform for nitrations but has largely been superseded by nitronium tetrafluoroborate (NO2BF4).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.[3]


N2O5 was first reported by Deville in 1840, who prepared it by treating silver nitrate (AgNO3) with chlorine.[4][5]

Structure and physical properties

Pure solid N2O5 is a salt, consisting of separated linear nitronium ions NO+2 and planar trigonal nitrate anions NO3. Both nitrogen centers have oxidation state +5. It crystallizes in the space group D4
(C6/mmc) with Z = 2, with the NO3 anions in the D3h sites and the NO+2 cations in D3d sites.[6]

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

being about 48 torr at 0 °C, 424 torr at 25 °C, and 760 torr at 32 °C (9 °C below the melting point).[7]

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) [8]

When gaseous N2O5 is cooled rapidly ("quenched"), one can obtain the metastable molecular form, which exothermically converts to the ionic form above −70 °C.[9]

Gaseous N2O5 absorbs ultraviolet light with dissociation into the free radicals nitrogen dioxide NO2 and nitrogen trioxide NO3 (uncharged nitrate). The absorption spectrum has a broad band with maximum at wavelength 160 nm.[10]


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 LiNO3 and bromine pentafluoride BrF5, in the ratio exceeding 3:1. The reaction first forms nitryl fluoride FNO2 that reacts further with the lithium nitrate:[6]

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 dioxide NO2 or N2O4 with ozone:[11]

2 NO2 + O3 → N2O5 + O2

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

2 O3 + N2O5 → 3 O2 + N2O5

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 Phosphoryl chloride POCl3 or nitryl chloride NO2Cl with silver nitrate AgNO3[6][12]


Dinitrogen pentoxide reacts with water (hydrolyses) to produce nitric acid HNO3. Thus, dinitrogen pentoxide is the anhydride of nitric acid:[9]

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 H2ON2O5 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").[13]

The reaction with hydrogen chloride HCl also gives nitric acid and nitryl chloride NO2Cl:[14]

N2O5 + HCl → HNO3 + NO2Cl

Dinitrogen pentoxide eventually decomposes at room temperature into NO2 and O2.[15][11] Decomposition is negligible if the solid is kept at 0 °C, in suitably inert containers.[6]

Dinitrogen pentoxide reacts with ammonia NH3 to give several products, including nitrous oxide N2O, ammonium nitrate NH4NO3, nitramide NH2NO2 and ammonium dinitramide NH4N(NO2)2, depending on reaction conditions.[16]

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:

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 oxide NO and oxygen. The thermal decomposition of 0.1 mM nitrogen dioxide at 1000 K is known to require about two seconds.[17]

Decomposition of dinitrogen pentoxide in carbon tetrachloride at 30 °C

Apart from the decomposition of N2O5 at high temperatures, it can also be decomposed in carbon tetrachloride CCl4 at 30 °C (303 K).[18] 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 N2O5 in CCl4:

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

N2O5 can also be decomposed in the presence of nitric oxide NO:

N2O5 + NO → 3 NO2

The rate of the initial reaction between dinitrogen pentoxide and nitric oxide of the elementary unimolecular decomposition.[19]


Nitration of organic compounds

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.[20] 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.[5][21]

Atmospheric occurrence

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.[22] Mixing ratios of several parts per billion by volume have been observed in polluted regions of the nighttime troposphere.[23] Dinitrogen pentoxide has also been observed in the stratosphere[24] 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.[25] Two important reactions of N2O5 in atmospheric aerosols are hydrolysis to form nitric acid[26] and reaction with halide ions, particularly Cl, to form ClNO2 molecules which may serve as precursors to reactive chlorine atoms in the atmosphere.[27][28]


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.


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  2. ^ Peter Steele Connell The Photochemistry of Dinitrogen Pentoxide. Ph. D. thesis, Lawrence Berkeley National Laboratory.
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