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Formation and occurrence Nitric oxide (NO), also a common pollutant, oxidizes in air to the dioxide: 2NO + O2 → 2NO2 NO2 is generated by the action of nitric acid on a variety of metals, such as copper or silver. 2HNO3 + Ag → AgNO3 + NO2 + H2O "Red fuming nitric acid" owes its red color to the presence of NO2. NO2 is generated in biological settings from decomposition of peroxynitrite (ONOO−), a potent oxidizing and nitrating agent formed from the reaction of nitric oxide with superoxide. Reactions Nitrogen dioxide exists in equilibrium with its dimer, dinitrogen tetroxide. 2 NO2 ↔ N2O4 ΔG = 45.53 kJ/mol The equilibrium favors NO2 at higher temperatures. Solid NO2 can be obtained from NO2 by very rapid cooling (for example with liquid nitrogen), although it is commonly contains N2O4. At −50 °C the crystals of N2O4, which is diamagnetic, are colorless, but they become honey-yellow at the melting point. The vapour at −10 °C is pale yellow and deepens as the temperature rises. Structure and bonding NO2 is a radical, having one unpaired electron, which renders this molecule paramagnetic. Low energy electronic transitions give rise to the visible color of this molecule. The molecule is nonlinear with bond distances and angles intermediate between those for the corresponding anion, nitrite, and the cation, nitronium.!! N-O distance (Å) |- | NO2+|| |- | NO2 || |- | NO2− || |- |} Safety and pollution considerations
See also More esoteric nitrogen oxides include N2O5 and the blue species N2O3. Oxidized (cationic) and reduced (anionic) derivatives of many of these oxides exist: nitrite (NO2−), nitrate (NO3−), nitronium or NO2+, and nitrosonium or NO+. NO2 is intermediate between nitrite and nitronium: NO2+ + e− → NO2 NO2 + e− → NO2− Reference | ||||||||||||
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