Nitrogen dioxide is the chemical compound with the formula NO
2. 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. Nitrogen
dioxide is a paramagnetic, bent molecule with C2v 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.
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 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. Preparation and reactions
Nitrogen dioxide typically arises via the oxidation of nitric oxide by oxygen in air:
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
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(NO3)2 → 2 PbO + 4 NO
2 + O 2
Alternatively, reduction of concentrated nitric acid by metal.
4 HNO 3 + Cu → Cu(NO3)2 + 2 NO
2 +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 Basic thermal properties
NO 2 exists in equilibrium with the colourless
gas dinitrogen tetroxide (N 2O
4): 2 NO
2 N 2O
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 predominates. Dinitrogen tetroxide (N
2O 4) can be obtained as a white solid with melting
point −11.2 °C. 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:
2 NO 2 → 2 NO + O
2 As an oxidizer
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. Hydrolysis
It hydrolyses to give nitric acid and nitrous acid:
2 NO 2/N
2O 4 + H
2O →HNO 2 + HNO
3 This reaction is one step in the Ostwald process
for the industrial production of nitric acid from ammonia. 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 Conversion to nitrates
NO 2 is used to generate anhydrous metal nitrates
from the oxides: MO + 3 NO
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 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 white fuming nitric acid, which spontaneously produces NO
2 above 0 °C. Symptoms of poisoning tend to appear several hours after inhalation of
a low but potentially fatal dose. Also, low concentrations 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.
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, 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. Limiting NOx 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 is also produced by atmospheric nuclear tests, and is responsible for the
reddish colour of mushroom clouds. 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.
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.
See also
Nitrite Dinitrogen tetroxide
Nitryl Nitrous oxide (N
2O), "laughing gas", a linear molecule, isoelectronic with CO
2 but with a nonsymmetric arrangement of atoms Nitric oxide, a problematic pollutant that
is short lived because it converts to NO 2 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 and "Answer to follow-up questions from CAFE
Nitrogen Dioxide Air Pollution Nitrogen dioxide pollution in the world
A review of the acute and long term impacts of exposure to nitrogen dioxide in the United
Kingdom IOM Research Report TM03 References