Nitrogen dioxide
Systematic name Dioxidonitrogen(•)[1] (additive)
Other names Nitrogen(IV) oxide[1] Nitrogen(II) oxide(-I)[citation needed]
Identifiers
CAS number	10102-44-0 Y
PubChem	3032552
ChemSpider	2297499 Y
EC number	233-272-6
UN number	1067
ChEBI	CHEBI:33101 Y
RTECS number	QW9800000
Gmelin Reference	976
Jmol-3D images	Image 1 Image 2 Image 3
SMILES O=[N]=O o:n:o [O-][N++][O-]
InChI InChI=1S/NO2/c2-1-3 Y Key: JCXJVPUVTGWSNB-UHFFFAOYSA-N Y InChI=1/NO2/c2-1-3 Key: JCXJVPUVTGWSNB-UHFFFAOYAA
Properties
Molecular formula	NO2•
Molar mass	46.0055 g mol-1
Appearance	Vivid orange gas
Density	2.62 g dm-3
Boiling point	21 °C, 294 K, 70 °F
Solubility in water	Reacts
Vapor pressure	98.80 kPa (at 20 °C)
Refractive index (nD)	1.449 (at 20 °C)
Structure
Molecular shape	Dihedral digonal
Thermochemistry
Std enthalpy of formation ΔfHo298	−34 kJ·mol−1[2]
Standard molar entropy So298	240 J·mol−1·K−1[2]
Hazards
MSDS	ICSC 0930
GHS pictograms
GHS signal word	Danger
GHS hazard statements	H270, H314, H330
GHS precautionary statements	P220, P260, P280, P284, P305+351+338, P310
EU Index	007-002-00-0
EU classification	T+
R-phrases	R26, R34, R8
S-phrases	(S1/2), S9, S26, S28, S36/37/39, S45
NFPA 704	0 3 0 OX
Related compounds
Related Nitrogen oxides	Dinitrogen pentoxide Dinitrogen tetroxide Dinitrogen trioxide Nitric oxide Nitrous oxide
Y (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Nitrogen dioxide is the chemical compound with the formula NO₂. It is one of several nitrogen oxides. NO₂ is an intermediate in the industrial synthesis of nitric acid, millions of tons of which are produced each year. This reddish-brown toxic gas has a characteristic sharp, biting odor and is a prominent air pollutant.^[3] Nitrogen dioxide is a paramagnetic, bent molecule with C_2v point group symmetry.

  Molecular properties

Nitrogen dioxide has a molar mass of 46.0055, which makes it heavier than air, whose average molar mass is 28.8. According to the ideal gas law, NO₂ is therefore more dense than air.

The bond length between the nitrogen atom and the oxygen atom is 119.7 pm. This bond length is consistent with a bond order of one and a quarter.

Unlike ozone, O₃, the ground electronic state of nitrogen dioxide is a doublet state, since there is one unpaired bonding electron delocalised over both bonds, which decreases the bond length; hence the bond order of one and a quarter.

  Occurrence

NO₂ exists in equilibrium with the colourless gas dinitrogen tetroxide (N₂O₄):

2 NO₂ N₂O₄

The equilibrium is characterized by ΔH = −57.23 kJ/mol, which is exothermic. Resulting from an endergonic reaction at higher temperatures, the paramagnetic monomer is favored. Colourless diamagnetic N₂O₄ can be obtained as a solid melting at m.p. −11.2 °C.^[4]

  Preparation and reactions

Nitrogen dioxide typically arises via the oxidation of nitric oxide by oxygen in air:^[4]

2 NO + O₂ → 2 NO₂

In the laboratory, NO₂ can be prepared in a two step procedure by thermal decomposition of dinitrogen pentoxide, which is obtained by dehydration of nitric acid:

2 HNO₃ → N₂O₅ + H₂O

2 N₂O₅ → 4 NO₂ + O₂

The thermal decomposition of some metal nitrates also affords NO₂:

2 Pb(NO₃)₂ → 2 PbO + 4 NO₂ + O₂

Alternatively, reduction of concentrated nitric acid by metal (such as copper).

4 HNO₃ + Cu → Cu(NO₃)₂ + 2 NO₂ +2 H₂O

Or finally by adding concentrated nitric acid over tin. Stannic acid is produced as byproduct.

4HNO₃ + Sn → H₂O + H₂SnO₃ + 4 NO₂

  Main reactions

The chemistry of nitrogen dioxide has been investigated extensively. At 150 °C, NO₂ decomposes with release of oxygen via an endothermic process (ΔH = 114 kJ/mol):

2 NO₂ → 2 NO + O₂

As suggested by the weakness of the N–O bond, NO₂ is a good oxidizer. Consequently, it will combust, sometimes explosively, with many compounds, such as hydrocarbons.

It hydrolyzes to give nitric acid and nitrous acid:

2 NO₂/N₂O₄ + H₂O →HNO₂ + HNO₃

This reaction is one step in the Ostwald process for the industrial production of nitric acid from ammonia.^[5] Nitric acid decomposes slowly to nitrogen dioxide, which confers the characteristic yellow color of most samples of this acid:

4 HNO₃ → 4 NO₂ + 2 H₂O + O₂

NO₂ is used to generate anhydrous metal nitrates from the oxides:^[4]

MO + 3 NO₂ → 2 M(NO₃)₂ + NO

Alkyl and metal iodides give the corresponding nitrites:

2 CH₃I + 2 NO₂ → 2 CH₃NO₂ + I₂

TiI₄ + 4 NO₂ → Ti(NO₂)₄ + 2 I₂

  Safety and pollution considerations

Nitrogen dioxide is toxic by inhalation. However, as the compound is acrid and easily detectable by smell at low concentrations, inhalation exposure can generally be avoided. One potential source of exposure is fuming nitric acid, which spontaneously produces NO₂ above 0 °C. Symptoms of poisoning (lung edema) tend to appear several hours after inhalation of a low but potentially fatal dose. Also, low concentrations (4 ppm) will anesthetize the nose, thus creating a potential for overexposure.

There is some evidence that long-term exposure to NO₂ at concentrations above 40–100 µg/m³ may decrease lung function and increase the risk of respiratory symptoms.^[6]

Nitrogen dioxide is formed in most combustion processes using air as the oxidant. At elevated temperatures nitrogen combines with oxygen to form nitric oxide:

O₂ + N₂ → 2 NO

Nitric oxide can be oxidized in air to form nitrogen dioxide. At normal atmospheric concentrations this is a very slow process.

2 NO + O₂ → 2 NO₂

The most prominent sources of NO₂ are internal combustion engines,^[7] thermal power stations and, to a lesser extent, pulp mills. Butane gas heaters and stoves are also sources. The excess air required for complete combustion of fuels in these processes introduces nitrogen into the combustion reactions at high temperatures and produces nitrogen oxides (NO_x). Limiting NO_x production demands the precise control of the amount of air used in combustion. In households, kerosene heaters and gas heaters are sources of nitrogen dioxide.

  Nitrogen Dioxide 2011 tropospheric column density.

Nitrogen dioxide is also produced by atmospheric nuclear tests, and is responsible for the reddish colour of mushroom clouds.^{[citation needed]}

Nitrogen dioxide is a large scale pollutant, with rural background ground level concentrations in some areas around 30 µg/m³, not far below unhealthy levels. Nitrogen dioxide plays a role in atmospheric chemistry, including the formation of tropospheric ozone. A 2005 study by researchers at the University of California, San Diego, suggests a link between NO₂ levels and Sudden Infant Death Syndrome.^[8]

Nitrogen dioxide is also produced naturally during electrical storms. The term for this process is "atmospheric fixation of nitrogen". The rain produced during such storms is especially good for the garden as it contains trace amounts of fertilizer. (Henry Cavendish 1784, Birkland -Eyde Process 1903, et-al)

  See also

  Nitrogen dioxide (NO₂) gas converts to the colorless gas dinitrogen tetroxide (N₂O₄) at low temperatures, and converts back to NO₂ at higher temperatures. The bottles in this photograph contain equal amounts of gas at different temperatures.

• Nitrite
• Dinitrogen tetroxide
• Nitryl
• Nitrous oxide (N₂O), "laughing gas", a linear molecule, isoelectronic with CO₂ but with a nonsymmetric arrangement of atoms (NNO)
• Nitric oxide (NO), a problematic pollutant that is short lived because it converts to NO₂ in the presence of free oxygen.

  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
• Nitrogen dioxide pollution in the world (image)
• A review of the acute and long term impacts of exposure to nitrogen dioxide in the United Kingdom IOM Research Report TM/04/03

  References