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dangerous oxidizer that once was used as a reagent when dissolved in chloroform for nitrations but has largely been superseded by nitronium tetrafluoroborate (NO2BF4).

[6][7] Pure solid N2O5 is a salt, consisting of separated linear nitronium ions NO+2 and planar trigonal nitrate anions NO−3.

[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.

The reaction first forms nitryl fluoride FNO2 that reacts further with the lithium nitrate:[8] The compound can also be created in the gas phase by reacting nitrogen dioxide NO2 or N2O4 with ozone:[13] However, the product catalyzes the rapid decomposition of ozone:[13] Dinitrogen pentoxide is also formed when a mixture of oxygen and nitrogen is passed through an electric discharge.

The equation below refers to the decomposition of N2O5 in CCl4: And this reaction follows the first order rate law that says: N2O5 can also be decomposed in the presence of nitric oxide NO: The rate of the initial reaction between dinitrogen pentoxide and nitric oxide of the elementary unimolecular decomposition.

[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.

Variations in N2O5 reactivity in aerosols can result in significant losses in tropospheric ozone, hydroxyl radicals, and NOx concentrations.