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Nitrogen dioxide

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Nitrogen dioxide
Stick model of nitrogen dioxide
Stick model of nitrogen dioxide
Spacefill model of nitrogen dioxide
Spacefill model of nitrogen dioxide
Nitrogen dioxide in a test tube
Names
Preferred IUPAC name
Nitrogen dioxide
Systematic IUPAC name
Nitrogen dioxide
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.030.234 Edit this at Wikidata
EC Number
  • 233-272-6
976
RTECS number
  • QW9800000
UN number UN 1067
  • InChI=1S/NO2/c2-1-3
    Key: JCXJVPUVTGWSNB-UHFFFAOYSA-N
  • InChI=1/NO2/c2-1-3
    Key: JCXJVPUVTGWSNB-UHFFFAOYAA
  • N(=O)[O]
Properties
NO
2
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 (11.8 °F; 261.9 K)
Boiling point 21.1 °C (70.0 °F; 294.2 K)
reacts
1.449 (20 °C)
Structure
bent, C2v
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazard OX: Oxidizer. E.g. potassium perchlorate
3
0
0
OX
Flash point Non-flammable
Related compounds
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Nitrogen dioxide is the chemical compound with the formula NO
2. One of several nitrogen oxides, NO
2 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 C2v point group symmetry.

Occurrence

NO
2 exists in equilibrium with dinitrogen tetroxide (N
2O
4):

2 NO
2 ⇌ N
2O
4

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 N
2O
4 can be obtained as a solid melting at m.p. −11.2 °C.[1]

Preparation and reactions

Nitrogen dioxide typically arises via the oxidation of nitric oxide by oxygen in air:[1]

2 NO + O
2 → 2 NO
2

In the laboratory, NO
2 can be prepared in a two step procedure by thermal decomposition of dinitrogen pentoxide, which is obtained by dehydration of nitric acid:

2 HNO
3 → N
2O
5 + H
2O
2 N
2O
5 → 4 NO
2 + O
2

The thermal decomposition of some metal nitrates also affords NO
2:

2 Pb(NO
3)
2 → 2 PbO + 4 NO
2 + O
2

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

4 HNO
3 + Cu → Cu(NO3)2 + 2 NO
2 +2 H2O


Main reactions

The chemistry of nitrogen dioxide has been investigated extensively. At 150 °C, NO
2 decomposes with release of oxygen via an endothermic process (ΔH = 114 kJ/mol):

2 NO
2 → 2 NO + O
2

As suggested by the weakness of the N–O bond, NO
2 is a good oxidizer. Consequently, it will combust, sometimes explosively, with many compounds, such as hydrocarbons.

It hydrolyzes with disproportionation to give nitric acid:

3 NO
2 + H
2O → NO + 2 HNO
3

This reaction is one step in the Ostwald process for the industrial production of nitric acid from ammonia.[2] Nitric acid decomposes slowly to nitrogen dioxide, which confers the characteristic yellow color of most samples of this acid:

4 HNO
3 → 4 NO
2 + 2 H
2O + O
2

NO
2 is used to generate anhydrous metal nitrates from the oxides:[1]

MO + 3 NO
2 → 2 M(NO
3)
2 + NO

Alkyl and metal iodides give the corresponding nitrites:

2 CH
3I + 2 NO
2 → 2 CH
3NO
2 + I
2
TiI
4 + 4 NO
2 → Ti(NO
2)
4 + 2 I
2

The pure gas is produced by adding concentrated nitric acid over tin. Stannic acid is produced as byproduct.

4HNO3 + Sn → H2O + H2SnO3 + 4 NO2

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 NO
2 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.

Long-term exposure to NO
2 at concentrations above 40– 100 µg/m3 causes adverse health effects.[3]

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

O
2 + N
2 → 2 NO

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

2 NO + O
2 → 2 NO
2

The most important sources of NO
2 are internal combustion engines,[4] 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.

Nitrogen Dioxide 2009 tropospheric column density.

Nitrogen dioxide is also produced by atmospheric nuclear tests, and is responsible for the reddish colour of mushroom clouds.[5]

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 NO
2 levels and Sudden Infant Death Syndrome.[6]

See also

Nitrogen dioxide (NO
2) gas converts to the colorless gas dinitrogen tetroxide (N
2O
4) at low temperatures, and converts back to NO
2 at higher temperatures. The bottles in this photograph contain equal amounts of gas at different temperatures.

References

  1. ^ a b c Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  2. ^ Michael Thiemann, Erich Scheibler, Karl Wilhelm Wiegand “Nitric Acid, Nitrous Acid, and Nitrogen Oxides” in Ullmann’s Encyclopedia of Industrial Chemistry, Wiley-VCH, 2005, Weinheim.
  3. ^ "Health Aspects of Air Pollution with Particulate Matter,Ozone and Nitrogen Dioxide" (PDF). Retrieved 2008-02-25.
  4. ^ Son, Busoon (2004). "Estimation of occupational and nonoccupational nitrogen dioxide exposure for Korean taxi drivers using a microenvironmental model". Environmental Research. 94 (3): 291–296. doi:10.1016/j.envres.2003.08.004. PMID 15016597. Retrieved 2008-02-25. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help); Unknown parameter |month= ignored (help)
  5. ^ "Air emissions". Botnia. Retrieved 2008-02-25.
  6. ^ "Sids Linked to Nitrogen Dioxide Pollution". Retrieved 2008-02-25.
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