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Nitrogen dioxide

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 C_2v 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.

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 NO_x 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.

Unlike ozone (O₃) the ground electronic state of nitrogen dioxide is a doublet state, since nitrogen has one unpaired electron, 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.

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

Selected reactions

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.

Thermal properties

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

As an oxidizer

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.

Hydrolysis

NO₂ reacts with water to give nitric acid and nitrous acid:

2 NO₂ + H₂O → HNO₃ + HNO₂

This reaction is one of the steps in 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 reactiivity 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,^: 223 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.

Outdoors, NO₂ can be a result of traffic from motor vehicles. Indoors, exposure arises from cigarette smoke, and butane and kerosene heaters and stoves. Indoor exposure levels of NO₂ are, on average, at least three times higher in homes with gas stoves compared to electric stove.

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. Workers in high voltage areas especially those with spark or plasma creation are at risk. 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".

Toxicity

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.

Acute exposure

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.

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.

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.

Long-term

Exposure to low levels of NO₂ over time can cause changes in lung function. 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. Children exposed to NO₂ are more likely to be admitted to hospital with asthma.

Environmental effects

Interaction of NO₂ and other NO_x with water, oxygen and other chemicals in the atmosphere can form acid rain which harms sensitive ecosystems such as lakes and forests. Elevated levels of NO
₂ can also harm vegetation, decreasing growth, and reduce crop yields.

References

Cited sources

Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). CRC Press. ISBN 978-1-4398-5511-9.

External links

International Chemical Safety Card 0930
National Pollutant Inventory – Oxides of nitrogen fact sheet
NIOSH Pocket Guide to Chemical Hazards
WHO-Europe reports: Health Aspects of Air Pollution (2003) (PDF) and "Answer to follow-up questions from CAFE (2004) (PDF)
Nitrogen Dioxide Air Pollution
Current global map of nitrogen dioxide distribution
A review of the acute and long term impacts of exposure to nitrogen dioxide in the United Kingdom IOM Research Report TM/04/03
Reaction of nitrogen dioxide with hydrocarbons and its influence on spontaneous ignition