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Nitrogen dioxide is the chemical compound with the formula NO2. It is 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.[1] Nitrogen dioxide is a paramagnetic, bent molecule with C2v point group symmetry.

Molecular properties Edit

Nitrogen dioxide has a molar mass of 46.0055, which makes it heavier than air, whose average molar mass is 28.8.

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,[2] which decreases the alpha effect compared to nitrite and creates a weak bonding interaction with the oxygen lone pairs. The lone electron in NO
2
also means that this compound is a free radical, so the formula for nitrogen dioxide is often written as ·NO2.

Preparation and reactionsEdit

Nitrogen dioxide typically arises via the oxidation of nitric oxide by oxygen in air:[3]

2 NO + O
2
→ 2 NO
2

In the laboratory, NO
2
can be prepared in a two-step procedure where dehydration of nitric acid produces dinitrogen pentoxide, which subsequently undergoes thermal decomposition:

2 HNO
3
N
2
O
5
+ H2O
2 N
2
O
5
→ 4 NO
2
+ O
2

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

2 Pb(NO3)2 → 2 PbO + 4 NO
2
+ O
2

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

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

Or finally by adding concentrated nitric acid over tin; hydrated tin dioxide is produced as byproduct.

4HNO3 + Sn → H2O + H2SnO3 + 4 NO2

Main reactionsEdit

Basic thermal propertiesEdit

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

2 NO
2
15px N
2
O
4

The equilibrium is characterized by ΔH = −57.23 kJ/mol, which is exothermic. NO2 is favored at higher temperatures, while at lower temperatures, dinitrogen tetroxide (N2O4) predominates. Dinitrogen tetroxide (N
2
O
4
) can be obtained as a white solid with melting point −11.2 °C.[3] NO2 is paramagnetic due to its unpaired electron, while N2O4 is diamagnetic.

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 an oxidizer Edit

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.

HydrolysisEdit

It hydrolyses to give nitric acid and nitrous acid:

2 NO
2
/N
2
O
4
+ H
2
O
HNO
2
+ HNO
3

This reaction is one step in the Ostwald process for the industrial production of nitric acid from ammonia.[4] 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
2
O
+ O
2

Conversion to nitratesEdit

NO
2
is used to generate anhydrous metal nitrates from the oxides:[3]

MO + 3 NO
2
M(NO
3
)
2
+ NO

Alkyl and metal iodides give the corresponding nitrites:

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

Safety and pollution considerationsEdit

File:Aura OMI Nitrogen dioxide troposphere column.png
File:NO2-N2O4.jpg

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 red 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.

There is some evidence that long-term exposure to NO
2
at concentrations above 40–100 µg/m3 may decrease lung function and increase the risk of respiratory symptoms.[5]

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. At normal atmospheric concentrations, this is a very slow process.

2 NO + O
2
→ 2 NO
2

The most prominent sources of NO
2
are internal combustion engines,[6] 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 (NOx). Limiting NOx production demands the precise control of the amount of air used in combustion. In households, kerosene heaters and gas heaters[7] are sources of nitrogen dioxide.

Nitrogen dioxide is also produced by atmospheric nuclear tests, and is responsible for the reddish colour of mushroom clouds.[8]

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.[9]

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.Script error (Henry Cavendish 1784, Birkland -Eyde Process 1903, et-al)

See alsoEdit

ReferencesEdit

External linksEdit

Template:Oxides

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