Nitrogen dioxide is a chemical compound with the formula NO2. One of several nitrogen oxides, nitrogen dioxide is a reddish-brown gas. It is a paramagnetic, bent molecule with C2v point group symmetry. Industrially, NO2 is an intermediate in the synthesis of nitric acid, millions of tons of which are produced each year, primarily for the production of fertilizers. It is a free radical. Nitrogen dioxide is poisonous and can be fatal if inhaled in large quantities. Cooking with a gas stove produces nitrogen dioxide which causes poorer indoor air quality. Combustion of gas can lead to increased concentrations of nitrogen dioxide throughout the home environment which is linked to respiratory issues and diseases. The LC50 (median lethal dose) for humans has been estimated to be 174 ppm for a 1-hour exposure. It is also included in the NOx family of atmospheric pollutants.
Properties Nitrogen dioxide is a reddish-brown gas with a pungent, acrid odor above 21.2 °C (70.2 °F; 294.3 K) and becomes a yellowish-brown liquid below 21.2 °C (70.2 °F; 294.3 K). It forms an equilibrium with its dimer, dinitrogen tetroxide (N2O4), and converts almost entirely to N2O4 below −11.2 °C (11.8 °F; 261.9 K). The bond length between the nitrogen atom and the oxygen atom is 119.7 pm. This bond length is consistent with a bond order between one and two. Nitrogen dioxide is a doublet state.
Preparation
Industrially, nitrogen dioxide is produced and transported as its cryogenic liquid dimer, dinitrogen tetroxide. It is produced industrially by the oxidation of ammonia, the Ostwald Process. This reaction is the first step in the production of nitric acid:
4 NH3 + 7 O2 → 4 NO2 + 6 H2O It can also be produced by the oxidation of nitrosyl chloride:
2 NOCl + O2 → 2NO2 + Cl2 Instead, most laboratory syntheses stabilize and then heat the nitric acid to accelerate the decomposition. For example, the thermal decomposition of some metal nitrates generates NO2:
Pb(NO3)2 → PbO + 2 NO2 + 1⁄2 O2 Alternatively, dehydration of nitric acid produces nitronium nitrate...
2 HNO3 → N2O5 + H2O 6 HNO3 + 1⁄2 P4O10 → 3 N2O5 + 2 H3PO4 ...which subsequently undergoes thermal decomposition:
N2O5 → 2 NO2 + 1⁄2 O2 NO2 is generated by the reduction of concentrated nitric acid with a metal (such as copper):
4 HNO3 + Cu → Cu(NO3)2 + 2 NO2 + 2 H2O Nitric acid decomposes slowly to nitrogen dioxide by the overall reaction:
4 HNO3 → 4 NO2 + 2 H2O + O2 The nitrogen dioxide so formed confers the characteristic yellow color often exhibited by this acid. However, the reaction is too slow to be a practical source of NO2.
Selected reactions Nitrogen dioxide interconverts to other N-O compounds. At low temperatures, NO2 reversibly converts to the colourless gas dinitrogen tetroxide (N2O4):
2 NO2 ⇌ N2O4 The exothermic equilibrium has enthalpy change ΔH = −57.23 kJ/mol. At 150 °C (302 °F; 423 K), NO2 decomposes with release of oxygen via an endothermic process (ΔH = 14 kJ/mol):
2 NO2 →2 NO + O2 Nitric dioxide combines with nitric oxide to reversibly give dinitrogen trioxide:
NO + NO2 ⇌ N2O3 This reaction is a step in the lead chamber process for producing sulfuric acid from sulfur dioxide. Absorption of light at wavelengths shorter than about 400 nm results in photolysis to form NO + O (atomic oxygen). In the atmosphere the addition of the oxygen atom so formed to O2 results in ozone.
Hydrolysis NO2 reacts with water to give nitric acid and nitrous acid:
2NO2 + H2O → HNO3 + HNO2 The nitrous acid further degrades to nitric acid and nitric oxide. These reactions underpin the Ostwald process for the industrial production of nitric acid from ammonia. This reaction is negligibly slow at low concentrations of NO2 characteristic of the ambient atmosphere, although it does proceed upon NO2 uptake to surfaces. Such surface reaction is thought to produce gaseous HNO2 (often written as HONO) in outdoor and indoor environments.
Conversion to nitrates NO2 is used to generate anhydrous metal nitrates from the oxides:
MO + 3 NO2 → M(NO3)2 + NO Alkyl and metal iodides give the corresponding nitrates:
TiI4 + 8 NO2 → Ti(NO3)4 + 4 NO + 2 I2
With organic compounds The reactivity of nitrogen dioxide toward organic compounds has long been known. For example, it reacts with amides to give N-nitroso derivatives. It is used for nitrations under anhydrous conditions.
Uses NO2 is used as an intermediate in the manufacturing of nitric acid, as a nitrating agent in the manufacturing of chemical explosives, as a polymerization inhibitor for acrylates, as a flour bleaching agent, and as a room temperature sterilization agent. It is also used as an oxidizer in rocket fuel, for example in red fuming nitric acid; it was used in the Titan rockets, to launch Project Gemini, in the maneuvering thrusters of the Space Shuttle, and in uncrewed space probes sent to various planets.
Environmental presence
Nitrogen dioxide typically arises via the oxidation of nitric oxide by oxygen in air (e.g. as result of corona discharge):
2 NO + O2 → 2 NO2 NO2 is introduced into the environment by natural causes, including entry from the stratosphere, bacterial respiration, volcanos, and lightning. These sources make NO2 a trace gas in the atmosphere of Earth, where it plays a role in absorbing sunlight and regulating the chemistry of the troposphere, especially in determining ozone concentrations.
Anthropogenic sources
Nitrogen dioxide also forms in most combustion processes. At elevated temperatures nitrogen combines with oxygen to form nitrogen dioxide:
N2 + 2 O2 → 2 NO2 For the general public, the most prominent sources of NO2 are internal combustion engines, as combustion temperatures are high enough to thermally combine some of the nitrogen and oxygen in the air to form NO2. Nitrogen dioxide accounts for a small fraction (generally well under 0.1) of NOx auto emissions.
… excerpt ends here. Continue reading the full article.
