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Chemical compound with formula NO₂ From Wikipedia, the free encyclopedia
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.
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NO 2 converts to the colorless dinitrogen tetroxide (N 2O 4) at low temperatures and reverts to NO 2 at higher temperatures. | |||
| Names | |||
|---|---|---|---|
| IUPAC name
Nitrogen dioxide | |||
| Other names
Nitrogen(IV) oxide,[1] deutoxide of nitrogen | |||
| Identifiers | |||
3D model (JSmol) |
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| ChEBI | |||
| ChemSpider | |||
| ECHA InfoCard | 100.030.234 | ||
| EC Number |
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| 976 | |||
PubChem CID |
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| RTECS number |
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| UNII | |||
| UN number | 1067 | ||
CompTox Dashboard (EPA) |
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| Properties | |||
| NO• 2 | |||
| Molar mass | 46.005 g·mol−1 | ||
| Appearance | Brown gas[2] | ||
| Odor | Chlorine-like | ||
| Density | 1.880 g/L[2] | ||
| Melting point | −9.3 °C (15.3 °F; 263.8 K)[2] | ||
| Boiling point | 21.15 °C (70.07 °F; 294.30 K)[2] | ||
| Hydrolyses | |||
| Solubility | Soluble in CCl 4, nitric acid,[3] chloroform | ||
| Vapor pressure | 98.80 kPa (at 20 °C) | ||
| +150.0·10−6 cm3/mol[4] | |||
Refractive index (nD) |
1.449 (at 20 °C) | ||
| Structure | |||
| C2v | |||
| Bent | |||
| Thermochemistry[5] | |||
Heat capacity (C) |
37.2 J/(mol·K) | ||
Std molar entropy (S⦵298) |
240.1 J/(mol·K) | ||
Std enthalpy of formation (ΔfH⦵298) |
+33.2 kJ/mol | ||
| Hazards | |||
| Occupational safety and health (OHS/OSH): | |||
Main hazards |
Poison, oxidizer | ||
| GHS labelling: | |||
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| Danger | |||
| H270, H314, H330 | |||
| P220, P260, P280, P284, P305+P351+P338, P310 | |||
| NFPA 704 (fire diamond) | |||
| Lethal dose or concentration (LD, LC): | |||
LC50 (median concentration) |
30 ppm (guinea pig, 1 h) 315 ppm (rabbit, 15 min) 68 ppm (rat, 4 h) 138 ppm (rat, 30 min) 1000 ppm (mouse, 10 min)[6] | ||
LCLo (lowest published) |
64 ppm (dog, 8 h) 64 ppm (monkey, 8 h)[6] | ||
| NIOSH (US health exposure limits): | |||
PEL (Permissible) |
C 5 ppm (9 mg/m3)[7] | ||
REL (Recommended) |
ST 1 ppm (1.8 mg/m3)[7] | ||
IDLH (Immediate danger) |
13 ppm[7] | ||
| Safety data sheet (SDS) | ICSC 0930 | ||
| Related compounds | |||
Related nitrogen oxides |
Dinitrogen pentoxide Dinitrogen tetroxide | ||
Related compounds |
Chlorine dioxide Carbon dioxide | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Nitrogen dioxide is poisonous and can be fatal if inhaled in large quantities.[8] 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.[9][10] The LC50 (median lethal dose) for humans has been estimated to be 174 ppm for a 1-hour exposure.[11] It is also included in the NOx family of atmospheric pollutants.
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).[7]
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.
Unlike ozone (O3) the ground electronic state of nitrogen dioxide is a doublet state, since nitrogen has one unpaired electron,[12] which decreases the alpha effect compared with nitrite and creates a weak bonding interaction with the oxygen lone pairs. The lone electron in NO2 also means that this compound is a free radical, so the formula for nitrogen dioxide is often written as •NO2.
The reddish-brown color is a consequence of preferential absorption of light in the blue region of the spectrum (400–500 nm), although the absorption extends throughout the visible (at shorter wavelengths) and into the infrared (at longer wavelengths). 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.
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:[13]
It can also be produced by the oxidation of nitrosyl chloride:
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:[14]
Alternatively, dehydration of nitric acid produces nitronium nitrate...
...which subsequently undergoes thermal decomposition:
NO2 is generated by the reduction of concentrated nitric acid with a metal (such as copper):
Nitric acid decomposes slowly to nitrogen dioxide by the overall reaction:
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.
At low temperatures, NO2 reversibly converts to the colourless gas dinitrogen tetroxide (N2O4):
The exothermic equilibrium has enthalpy change ΔH = −57.23 kJ/mol.[15]
At 150 °C (302 °F; 423 K), NO2 decomposes with release of oxygen via an endothermic process (ΔH = 14 kJ/mol):
As suggested by the weakness of the N–O bond, NO2 is a good oxidizer. Consequently, it will combust, sometimes explosively, in the presence of hydrocarbons.[16]
NO2 reacts with water to give nitric acid and nitrous acid:
This reaction is one of the steps in the Ostwald process for the industrial production of nitric acid from ammonia.[13] 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.[17]
NO2 is used to generate anhydrous metal nitrates from the oxides:[15]
Alkyl and metal iodides give the corresponding nitrates:[12]
The reactivity of nitrogen dioxide toward organic compounds has long been known.[18] For example, it reacts with amides to give N-nitroso derivatives.[19] It is used for nitrations under anhydrous conditions.[20]
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,[21]: 223 and as a room temperature sterilization agent.[22] 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.[23]
Nitrogen dioxide typically arises via the oxidation of nitric oxide by oxygen in air (e.g. as result of corona discharge):[15]
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.[24]
Nitrogen dioxide also forms in most combustion processes. At elevated temperatures nitrogen combines with oxygen to form nitrogen dioxide:
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.[8]
Outdoors, NO2 can be a result of traffic from motor vehicles.[25] Indoors, exposure arises from cigarette smoke,[26] and butane and kerosene heaters and stoves.[27] Indoor exposure levels of NO2 are, on average, at least three times higher in homes with gas stoves compared to electric stove.[28][29]
Workers in industries where NO2 is used are also exposed and are at risk for occupational lung diseases, and NIOSH has set exposure limits and safety standards.[7] Workers in high voltage areas especially those with spark or plasma creation are at risk.[citation needed] Agricultural workers can be exposed to NO2 arising from grain decomposing in silos; chronic exposure can lead to lung damage in a condition called "silo-filler's disease".[30][31]
NO2 diffuses into the epithelial lining fluid (ELF) of the respiratory epithelium and dissolves. There, it chemically reacts with antioxidant and lipid molecules in the ELF. The health effects of NO2 are caused by the reaction products or their metabolites, which are reactive nitrogen species and reactive oxygen species that can drive bronchoconstriction, inflammation, reduced immune response, and may have effects on the heart.[32]
Acute harm due to NO2 exposure is rare. 100–200 ppm can cause mild irritation of the nose and throat, 250–500 ppm can cause edema, leading to bronchitis or pneumonia, and levels above 1000 ppm can cause death due to asphyxiation from fluid in the lungs. There are often no symptoms at the time of exposure other than transient cough, fatigue or nausea, but over hours inflammation in the lungs causes edema.[33][34]
For skin or eye exposure, the affected area is flushed with saline. For inhalation, oxygen is administered, bronchodilators may be administered, and if there are signs of methemoglobinemia, a condition that arises when nitrogen-based compounds affect the hemoglobin in red blood cells, methylene blue may be administered.[35][36]
It is classified as an extremely hazardous substance in the United States as defined in Section 302 of the U.S. Emergency Planning and Community Right-to-Know Act (42 U.S.C. 11002), and it is subject to strict reporting requirements by facilities which produce, store, or use it in significant quantities.[37]
Exposure to low levels of NO2 over time can cause changes in lung function.[38] Cooking with a gas stove is associated with 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.[9][10] Children exposed to NO2 are more likely to be admitted to hospital with asthma.[39]
Interaction of NO2 and other NOx with water, oxygen and other chemicals in the atmosphere can form acid rain which harms sensitive ecosystems such as lakes and forests.[40] Elevated levels of NO
2 can also harm vegetation, decreasing growth, and reduce crop yields.[41]
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