Nitrogen trioxide or nitrate radical is an oxide of nitrogen with formula NO 3, consisting of three oxygen atoms covalently bound to a nitrogen atom. This highly unstable blue compound has not been isolated in pure form, but can be generated and observed as a short-lived component of gas, liquid, or solid systems.[1]
Nitrogen trioxide is an important intermediate in reactions between atmospheric components, including the destruction of ozone.[1][2]
History
The existence of the NO 3 radical was postulated in 1881-1882 by Hautefeuille and Chappuis to explain the absorption spectrum of air subjected to a silent electrical discharge.[1]
The absorption spectrum of NO 3 has a broad band for light with wavelengths from about 500 to 680nm, with three maxima in the visible at 590, 662, and 623nm. Absorption in the range 640–680nm does not lead to dissociation but to fluorescence: specifically, from about 605 to 800 nm following excitation at 604.4nm, and from about 662 to 800nm following excitation at 661.8nm.[1] In water solution, another absorption band appears at about 330nm (ultraviolet). An excited state NO* 3 can be achieved by photons of wavelength less than 595nm.[1]
Preparation
Nitrogen trioxide can be prepared in the gas phase by mixing nitrogen dioxide and ozone:[1]
NO 2 + O 3 → NO 3 + O 2
This reaction can be performed also in the solid phase or water solutions, by irradiating frozen gas mixtures, flash photolysis and radiolysis of nitrate salts and nitric acid, and several other methods.[1]
In the atmosphere, the formation of nitrate radical also occurs from the reaction:[1][3]NO3 concentrations maximize during the night due to the absence of photolysis.[3] In its nightly abundance, the nitrate radical can oxidize volatile organic compounds (VOCs).[3][4]
As a dominant nighttime oxidizing agent, NO3 reacts with a variety of VOCs, in which the mechanism and the rate of the reaction will depend on the species. As an example, oxygenates such as aldehydes are oxidized by abstracting an H-atom, resulting in nitric acid HNO3 as one of the products.[3] The lifetime of NO3 in the presence of these compounds typically ranges from minutes to days.[3]
Conversely, unsaturated hydrocarbons (e.g. alkenes) are commonly oxidized by adding the nitrate radical to a carbon double bond.[3] The reaction rates with these species are higher compared to those of oxygenates, notably affecting the abundance of the unsaturated hydrocarbons.[3][5] In forested areas, for instance, NO3 can rapidly oxidize biogenic volatile organic compounds (BVOCs), such as terpenes and isoprene, which have been studied to contribute as a source of secondary organic aerosols (SOAs).[3][4][5]
Rapid daytime photolysis
The reaction that efficiently removes the nitrate radical from the troposphere during the day is its photolysis:in which the pathway followed will depend on the wavelength.[1][3] When the sun is overhead, the photodissociation reaches maximum rates ranging from J = 0.02 to 0.2 s-1.[1] Therefore, the daytime lifetime of NO3 is usually less than 5 seconds.[1][5]
References
123456789101112131415R. P. Wayne, I. Barnes, P. Biggs, J. P. Burrows, C. E. Canosa-Mas, J. Hjorth, G. Le Bras. G. K. Moortgat, D. Perner, G. Poulet, G. Restelli, and H. Sidebottom (1991): "The nitrate radical: Physics, chemistry, and the atmosphere". Atmospheric Environment. Part A. General Topics. volume 25, issue 1, pages 1-203. doi:10.1016/0960-1686(91)90192-A
↑Richard A. Graham and Harold S. Johnston (1978): "The photochemistry of the nitrate radical and the kinetics of the nitrogen pentoxide-ozone system". Journal of Physical Chemistry, volume 82, issue 3, pages 254-268. doi:10.1021/j100492a002
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