The nitrogen oxides are inorganic compounds essentially gaseous that link between nitrogen atoms and oxygen. Its group chemical formula is NO x , indicating that oxides have different ratios of oxygen and nitrogen.
Nitrogen heads group 15 of the periodic table, while oxygen heads group 16; both elements are members of period 2. This closeness is the cause that in oxides the N – O bonds are covalent. Thus, the bonds in nitrogen oxides are covalent.
All these bonds can be explained using the molecular orbital theory, which reveals the paramagnetism (an unpaired electron in the last molecular orbital) of some of these compounds. Of these, the most common compounds are nitric oxide and nitrogen dioxide.
The molecule in the upper image corresponds to the angular structure in the gas phase of nitrogen dioxide (NO 2 ). In contrast, nitric oxide (NO) has a linear structure (considering sp hybridization for both atoms).
The high temperatures of the rays break the energy barrier that prevents this reaction from occurring under normal conditions. What energy barrier? That formed by the triple bond N≡N, making the N 2 molecule an inert gas in the atmosphere .
Oxidation numbers for nitrogen and oxygen in their oxides
The electron configuration for oxygen is [He] 2s 2 2p 4 , needing only two electrons to complete the octet of its valence shell; that is, it can gain two electrons and have an oxidation number equal to -2.
On the other hand, the electron configuration for nitrogen is [He] 2s 2 2p 3 , being able to gain up to three electrons to fill its valence octet; for example, in the case of ammonia (NH 3 ) it has an oxidation number equal to -3. But oxygen is much more electronegative than hydrogen and “forces” nitrogen to share its electrons.
How many electrons can nitrogen share with oxygen? If you share the electrons of your valence shell one by one, you will reach the limit of five electrons, corresponding to an oxidation number of +5.
Different formulations and nomenclatures
Nitrogen oxides, in increasing order of nitrogen oxidation numbers, are:
– N 2 O, nitrous oxide (+1)
– NO, nitric oxide (+2)
– N 2 O 3 , dinitrogen trioxide (+3)
– NO 2 , nitrogen dioxide (+4)
– N 2 O 5 , dinitrogen pentoxide (+5)
Nitrous oxide (N 2 O)
Nitrous oxide (or popularly known as laughing gas) is a colorless gas, with a slightly sweet odor and little reactive. It can be visualized as a N 2 molecule (blue spheres) that has added an oxygen atom to one of its ends. It is prepared by the thermal decomposition of nitrate salts and is used as an anesthetic and analgesic.
Nitrogen has an oxidation number of +1 in this oxide, which means that it is not very oxidized and its demand for electrons is not pressing; however, it only needs to gain two electrons (one for each nitrogen) to become the stable molecular nitrogen.
In basic and acid solutions the reactions are:
N 2 O (g) + 2H + (aq) + 2e – => N 2 (g) + H 2 O (l)
N 2 O (g) + H 2 O (l) + 2e – => N 2 (g) + 2OH – (aq)
These reactions, although thermodynamically favored by the formation of the stable molecule N 2 , occur slowly and the reagents that donate the electron pair must be very strong reducing agents.
Nitric oxide (NO)
This oxide consists of a colorless, reactive and paramagnetic gas. Like nitrous oxide, it has a linear molecular structure, but with the great difference that the N = O bond also has the character of a triple bond.
NO is rapidly oxidized in air to produce NO 2 , thus generating more stable molecular orbitals with a more oxidized nitrogen atom (+4).
2NO (g) + O 2 (g) => 2NO 2 (g)
Biochemical and physiological studies are behind the benign role that this oxide has in living organisms.
It cannot form NN bonds with another NO molecule due to the delocalization of the unpaired electron in the molecular orbital, which is directed more towards the oxygen atom (due to its high electronegativity). The opposite occurs with NO 2 , which can form gaseous dimers.
Nitrogen trioxide (N 2 O 3 )
Dotted lines in the structure indicate double bond resonance. Like all atoms, they have sp 2 hybridization , the molecule is flat, and the molecular interactions are effective enough for nitrogen trioxide to exist as a blue solid below -101ºC. At higher temperatures it melts and dissociates into NO and NO 2 .
Why is it dissociated? Because the oxidation numbers +2 and +4 are more stable than +3, present the latter in the oxide for each of the two nitrogen atoms. This, again, can be explained by the stability of the molecular orbitals resulting from the disproportion.
In the image, the left side of N 2 O 3 corresponds to NO, while the right side to NO 2 . Logically, it is produced by the coalescence of the previous oxides at very cold temperatures (-20ºC). N 2 O 3 is nitrous acid anhydride (HNO 2 ).
Nitrogen dioxide and tetroxide (NO 2 , N 2 O 4 )
NO 2 is a reactive, paramagnetic, brown or brown gas. As it has an unpaired electron, it dimerizes (binds) with another gas molecule of NO 2 to form nitrogen tetroxide, a colorless gas, establishing an equilibrium between both chemical species:
2NO 2 (g) <=> N 2 O 4 (g)
It is a poisonous and versatile oxidizing agent, capable of disproportionate in its redox reactions in the ions (oxoanions) NO 2 – and NO 3 – (generating acid rain), or in NO.
Likewise, NO 2 is involved in complex atmospheric reactions causing variations in ozone (O 3 ) concentrations at terrestrial levels and in the stratosphere.
Dinitrogen pentoxide (N 2 O 5 )
When it is hydrated, it generates HNO 3 , and at higher concentrations of the acid the oxygen is mainly protonated with a positive partial charge —O + —H, accelerating the redox reactions.