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|- | align="center" colspan="2" bgcolor="#ffffff" | |- | Density and phase | 1443 kg/m³, liquid <br/> 3.4 kg/m³, gas at 294.25 K |- | NFPA 704 |
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| R-phrases | , |- | S-phrases | , , , ,<br />, |- Nitrogen dioxide is the chemical compound NO<sub>2</sub>. It is one of the several nitrogen oxides. This reddish-brown gas has a characteristic sharp, biting odor. NO<sub>2</sub> is one of the most prominent air pollutants and a poison by inhalation.
Nitric oxide (NO), a common pollutant, oxidizes in air to the dioxide: 2NO + O<sub>2</sub> → 2NO<sub>2</sub>
Ammonia oxidizing in air at 850°C, with a platinum gauze catalyst, produces NO<sub>2</sub>: 4NH<sub>3</sub> + 7O<sub>2</sub> → 4NO<sub>2</sub> + 6H<sub>2</sub>O
NO<sub>2</sub> is generated by the action of nitric acid on a variety of metals, such as copper or silver. 2HNO<sub>3</sub> + Ag → AgNO<sub>3</sub> + NO<sub>2</sub> + H<sub>2</sub>O
"Red fuming nitric acid" owes its red color to the presence of NO<sub>2</sub>.
NO<sub>2</sub> is generated in biological settings from decomposition of peroxynitrite (ONOO<sup>−</sup>), a potent oxidizing and nitrating agent formed from the reaction of nitric oxide with superoxide.
Nitrogen dioxide exists in equilibrium with its dimer, dinitrogen tetroxide. 2 NO<sub>2</sub> N<sub>2</sub>O<sub>4</sub> ΔG = 45.53 kJ/mol The equilibrium favors NO<sub>2</sub> at higher temperatures. Solid NO<sub>2</sub> can be obtained from NO<sub>2</sub> by very rapid cooling (for example with liquid nitrogen), although it commonly contains N<sub>2</sub>O<sub>4</sub>.
At −50 °C the crystals of N<sub>2</sub>O<sub>4</sub>, 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.
NO<sub>2</sub> 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.[1]
Nitrogen dioxide is toxic by inhalation. Symptoms of poisoning (lung edema) tend to appear several hours after one has inhaled a low but potentially fatal dose. Also, low concentrations (4 ppm) will anesthetize the nose, thus creating a potential for overexposure.
Long-term exposure to NO<sub>2</sub> at concentrations above 40–100 µg/m³ causes adverse health effects [1]. The most important source of NO<sub>2</sub> is internal combustion engines, which emit nitrogen oxides near people. A major industrial source is pulp mills.
This map, depicting results of satellite measurements, illustrates nitrogen dioxide as large scale pollutant, with rural background ground level concentrations in some areas around 30 µg/m³, not far below unhealthful levels. Nitrogen dioxide plays a role in atmospheric chemistry, including the formation of tropospheric ozone.
A recent study by researchers at the University of California, San Diego, suggests a link between NO<sub>2</sub> levels and Sudden Infant Death Syndrome [2].
More esoteric nitrogen oxides include N2O5 and the blue species N2O3.
Oxidized (cationic) and reduced (anionic) derivatives of many of these oxides exist: nitrite (NO<sub>2</sub><sup>−</sup>), nitrate (NO<sub>3</sub><sup>−</sup>), nitronium or NO<sub>2</sub><sup>+</sup>, and nitrosonium or NO<sup>+</sup>. NO<sub>2</sub> is intermediate between nitrite and nitronium: NO<sub>2</sub><sup>+</sup> + e<sup>−</sup> → NO<sub>2</sub> NO<sub>2</sub> + e<sup>−</sup> → NO<sub>2</sub><sup>−</sup>
