Hybridization of NO2

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Nitrogen Dioxide (NO2) involves an sp2 hybridization type. The simple way to determine the hybridization of NO2 is by counting the bonds and lone electron pairs around the nitrogen atom and by drawing the Lewis structure. We will also find that in nitrogen dioxide, there are two sigma bonds and one lone electron pair.

If we apply the hybridization rule now, then it states that if the sum of the number of sigma bonds, the electrons’ lone pair, and odd electrons is equal to three, then the hybridization is sp2. The molecular name, formula, and other related properties are tabulated below.

Name of the Molecule

Nitrogen Dioxide

Molecular Formula

NO2

Hybridization Type

sp2

Bond Angle

134°

Geometry

Bent


What is the Hybridization of Nitrogen Dioxide (NO2)?

In the hybridization of nitrogen dioxide i.e. NO2 (during the NO2 formation), let us take a look at the Nitrogen atom first. Here we will notice that the nitrogen atom is the centre atom and has only one lone electron. The atom, however, does not have an octet because it is short on electrons. Since there is an electron deficit in the nitrogen molecule, usually it tends to react with some other molecule (oxygen, in this case) for its octet completion. On the other hand, the two oxygen atoms have an octet of electrons each. When the bonding occurs, the two oxygen atoms will form a single and a double bond with the nitrogen atom.

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At the same time, nitrogen must have three hybridized orbitals that are used to harbour the two sigma bonds and including one electron resulting in sp2 hybridization. The three sp2 hybrid orbitals present in nitrogen will have one electron, and the p orbital will also have one electron. However, when it forms two sigma bonds, only one p orbital and sp2 hybrid orbital will contain one electron each. The p orbital will then form a pi bond with the oxygen atom.


Lewis Structure of NO2

The Lewis structure of NO2 has 17 valence electrons. In a Lewis structure, it's not common to have an odd number of valence electrons. For this reason, we'll try to get closer to an octet as we can be on the central Nitrogen (N) atom. It means it will only have 7 valence electrons.

The Nitrogen atom in the Lewis structure for NO2 is the least electronegative atom and passes at the centre of the structure.

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Molecular Geometry and Bond Angles of NO2

Since the Nitrogen Dioxide (NO2) has an extra electron in a nitrogen atom’s orbital,  it will result in a higher degree of repulsions. However, if we consider one lone electron or the single-electron region, there is less repulsion on bonding two oxygen atoms. So, the repulsions are unidentical. Resultantly, the oxygen atoms spread widely. In the classical sense, NO2 is sp2 hybridized.

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In general, the Single-electron orbitals are unhybridized, and pure p-orbitals, like in methyl free radicals.

However, a Single-electron orbital is known to be hybridized, when the central atom is bonded to the highly electronegative groups or atoms. It is because of the decreased electron density at the central atom, which attracts the odd electron-orbital density then, and thereby the size of the odd-electron orbital is decreased. Hence, the odd-electron orbital acquires some “s” character and, as a result, becomes hybridized.

For the same reason, ClO2, ClO3 & CF3 are sp3 hybridized.

The bond angle in NO2 is 115°, and here both N-O bonds are equivalent because of the resonance.

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The nitrogen dioxide (NO2), [nitrite or nitro ion] is the sp2 hybridized as well, and here, the lone pair orbital is hybridized.

It is to note that both the N-O bonds are equivalent because of the resonance. Even the double bonds behave similarly to lone pairs for the repulsions, in effect (also, note that SO2 has a bond angle of approx 120°).

Here, the bond angle in NO2- is 115°.

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Considering NO2+(nitronium ion), it is an sp hybridized with a bond angle of 180°.

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The hybridization of NO2+ has a non-equivalent resonating structure.

For example, the Acylium cations (RCO+) are linear, sp hybridized, and the triply bonded resonating structure is more stable because of the complete octet of all atoms.

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Bond Order for NO2+

NO2+ has a total of 16 valence electrons (2 x 6, where 12 from the oxygen atoms, 5 from the nitrogen atoms, and -1 due to the +1 charge); the Lewis structure can be given below.

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Therefore, the bond order of both N and O bonds is 2.


Some Key Points to Remember

A few basic key points that are to remember are listed below.

  • The two oxygen atoms have each electrons octet

  • In nitrogen dioxide (NO2), there are 1 lone electron pair and 2 sigma bonds

  • The p orbital of nitrogen atom forms a pi bond with the oxygen atom

FAQ (Frequently Asked Questions)

1. Which Bond Angle is Larger Among NO2+ and NO2?

In NO2(+), that is, in the nitronium ion, the N-atom has sp-hybridization; thus, it adopts the linear geometry, and the O-N-O bond angle is 180°.


While coming to NO2(-), that is, nitrite ion, the N-atom has sp2 hybridization; thus, it adopts the bent geometry, for NO2(-), and the actual O-N-O bond angle is 115 ° (slightly deviated from the expected 120° due to the repulsion between lone pair of electrons and interacting bond pairs).


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This is the 'nitronium ion,' and it clearly shows that the O–N–O bond angle in NO2(+) is 180°.


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And, this is the geometry or structure of NO2(–), having 115 N–O–N bond angle.

So, the NO2- has a bond angle less than NO2+.

2. Why Does the NO2 Carry a Negative Charge?

As that set of atoms are more stable in that configuration while an additional electron is available, the “octet rule” clarifies that an atom is more stable if it has a filled outermost orbital of s and p. There are one s, and 3 p’s, and each orbital carries 2 electrons, so, as a result, eight is the good number.


Since the nitrogen starts with 5 and oxygen with 6, by sharing the pairs, they form bonds and almost get up to 8, but there is an odd number. We call it as a radical - fairly unstable, prone to stealing a single electron from the other molecule (creating another radical and continuing - this kind of chain reaction is how we destroyed the ozone layer almost with a small number of chemicals.) If there happens to be a surplus electron around, NO2 is stabilized by adding it on. That’s why nitrite exists and why its charge is -1.