- Nitric oxide, also known as nitrogen monoxide (NO), nitrogen(II) oxide
- Nitrogen dioxide (NO2), nitrogen(IV) oxide
- Nitrous oxide (N2O), nitrogen(−I,III) oxide
- Nitrosylazide (N4O), nitrogen(−I,0,I,II) oxide
- Oxatetrazole (N4O)
- Dinitrogen trioxide (N2O3), nitrogen(II,IV) oxide
- Dinitrogen tetroxide (N2O4), nitrogen(IV) oxide
- Dinitrogen pentoxide (N2O5), nitrogen(V) oxide
- Trinitramide (N(NO2)3 or N4O6), nitrogen(0,IV) oxides
- Dinitramide (N(NO2)3−)
- Nitrite (NO−
- Nitrate (NO−
- Nitronium (NO+
- Nitrosonium (NO+
- peroxonitrite (ONO−
- Nitrate (NO3−), trioxonitrate(V) ion
Only the first three of these compounds can be isolated at room temperature. N2O3, N2O4, and N2O5 all decompose rapidly at room temperature. NO3, N4O, and N(NO2)3 are very reactive.
N2O is stable and rather unreactive at room temperature, while NO and NO2 are quite reactive but nevertheless quite stable when isolated.
NOx (often written NOx) refers to NO and NO2. They are produced during combustion, especially at high temperature. These two chemicals are important trace species in Earth's atmosphere. In the troposphere, during daylight, NO reacts with partly oxidised organic species (or the peroxy radical) to form NO2, which is then photolysed by sunlight to reform NO:
- NO + CH3O2 → NO2 + CH3O
- NO2 + sunlight → NO + O
The oxygen atom formed in the second reaction then goes on to form ozone; this series of reactions is the main source of tropospheric ozone. CH3O2 is just one example of many partly oxidised organic molecules that can react with NO to form NO2.
These reactions are rather fast, so NO and NO2 tend to cycle, but the sum of their concentration ([NO] + [NO2]) tends to remain fairly constant. Because of this cycling, it is convenient to think of the two chemicals as a group; hence the term NOx.
In addition to acting as a main precursor for tropospheric ozone, NOx is also harmful to human health in its own right.
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