nitrogen dioxide

Nitrogen dioxide
Identifiers
CAS number	10102-44-0 Y
PubChem	3032552
EC number	233-272-6
UN number	1067
ChEBI	33101
RTECS number	QW9800000
Properties
Molecular formula	NO2
Molar mass	46.0055(5) g/mol
Appearance	brown gas
Density	1449 kg/m3 (liquid, 20 °C) 3.4 kg/m3 (gas, 22 °C)
Melting point	-11.2 °C, 262 K, 12 °F
Boiling point	21.1 °C, 294 K, 70 °F
Solubility in water	reacts
Refractive index (nD)	1.449 (20 °C)
Structure
Molecular shape	bent, C2v
Hazards
MSDS	ICSC 0930
EU Index	007-002-00-0
EU classification	Highly toxic (T+) Corrosive (C)
R-phrases	R26, R34
S-phrases	(S1/2), S9, S26, S28,S36/37/39, S45
NFPA 704	0 3 0 OX
Flash point	Non-flammable
Related compounds
Related nitrogen oxides	Nitrous oxide Nitric oxide Dinitrogen trioxide Dinitrogen tetroxide Dinitrogen pentoxide
Y (what is this?)  (verify) 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 exists as a radical in nature.^[1] 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. Nitrogen dioxide is a paramagnetic bent molecule with C_2v point group symmetry.

Preparation and reactions

Nitrogen dioxide typically arises via the oxidation of nitric oxide by oxygen in air:^[2]

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

Monomer-dimer equilibrium

NO₂ exists in equilibrium with N₂O₄:

2 NO₂ N₂O₄

The equilibrium is characterized by ΔH = −57.23 kJ/mol, which is exothermic. Resulting from an endothermic 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.^[2]

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 and will sustain the combustion, sometimes explosively, with many compounds, such as hydrocarbons.

It hydrolyzes with disproportionation to give nitric acid:

3 NO₂ + H₂O → NO + 2 HNO₃

This reaction is one step in the Ostwald process for the industrial production of nitric acid from ammonia.^[3] 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:^[2]

MO + 3 NO₂ → 2 M(NO₃)₂ + NO

Similarly, alkyl and metal iodides give the corresponding nitrates:

2 CH₃I + 3 NO₂ → 2 CH₃NO₃ + NO + I₂

TiI₄ + 8 NO₂ → Ti(NO₃)₄ + 4 NO + 2 I₂

Safety and pollution considerations

Nitrogen dioxide is toxic by inhalation, but this could be avoided as the material is acrid and easily detected by our sense of smell. 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 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 NO₂ at concentrations above 40–100 µg/m³ causes adverse health effects.^[4]

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

2 O₂ + N₂ → 2 NO₂

The most important sources of NO₂ are internal combustion engines,^[5] thermal power stations and, to a lesser extent, pulp mills. Butane gas heaters and stoves is also a source. 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.

Nitrogen dioxide is also produced by atmospheric nuclear tests, and is responsible for the reddish colour of mushroom clouds.^[6]

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.^[7]

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.

• 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.

References

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)
• Computational Chemistry Wiki