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The nitrogen dioxide Lewis structure is an odd-electron (free-radical) species best drawn as O–N=O with one lone pair and one unpaired electron on nitrogen (resonance between the two N–O bonds); the molecule is bent, nitrogen is sp2-hybridized, the O–N–O angle is about 134°, and because the bent shape prevents dipole cancellation, NO2 is polar.
Figure 1: Lewis Structure of Nitrogen dioxide (NO2)
| Chemical Formula | NO2 |
| CAS Number | 10102-44-0 |
| Total Valence Electrons | 17 (odd-electron radical) |
| Lewis Structure | O–N=O, 1 lone pair + 1 unpaired e⁻ on N (resonance) |
| Electron-Domain Geometry | Trigonal planar (approx.) |
| Molecular Geometry | Bent / angular |
| Hybridization Model | sp2 (nitrogen) |
| O–N–O Bond Angle | ≈ 134° |
| Polarity | Polar molecule |
Nitrogen dioxide is a key intermediate in the industrial production of nitric acid via the Ostwald process, where ammonia is oxidized to NO and then to NO2, which is absorbed in water to give HNO3. It is therefore foundational to fertilizers, nitrates and explosives.
NO2 is also a precursor/oxidizer in the formation of nitrating mixtures for nitroaromatics and energetic materials, and it dimerizes to N2O4 (dinitrogen tetroxide), an important rocket oxidizer and a convenient liquid source of NO2. In the environment it is a significant urban air pollutant formed from combustion and a precursor to photochemical smog and acid rain.
Count valence electrons for NO2: nitrogen (Group 15) contributes 5 valence electrons; each oxygen (Group 16) contributes 6 valence electrons.
Total valence electrons = 5 from N + 6 from O + 6 from O = 17 electrons — an odd number, so NO2 is a free radical with one unpaired electron.
With 17 electrons, a perfect octet on all three atoms is impossible. Start with O–N–O single bonds (4 e), then build toward O=N–O with one double bond. The remaining electrons place two lone pairs on one oxygen, one lone pair plus the unpaired electron on nitrogen, and two lone pairs on the other oxygen.
The two N–O bonds are equivalent by resonance: the unpaired electron and the double-bond character delocalize over the two oxygens. The best simple Lewis picture is O–N=O with one lone pair and one unpaired electron on nitrogen, recognizing that the true structure is a resonance hybrid.
The Lewis picture gives roughly three electron domains around nitrogen: two N–O bonding domains and one domain associated with the lone pair/unpaired electron. Three domains adopt an approximately trigonal-planar electron arrangement.
Because the central atom does not have a perfect octet and carries a lone-pair/unpaired domain, the molecular geometry is bent. The O–N–O angle is wider than in a typical AX2E molecule — about 134° — because the unpaired electron exerts less repulsion than a full lone pair would.
Repulsion ordering still applies qualitatively: a lone-pair/unpaired domain repels the N–O bonding domains. However, a single unpaired electron occupies less space than a full lone pair, so it compresses the angle less — the O–N–O angle opens to ~134° rather than the ~120° seen in SO2.
Nitrogen in NO2 is commonly described as sp2-hybridized: three domains correspond to three trigonal-planar hybrids (two for σ bonds to oxygen, one holding the lone-pair/unpaired-electron density). The remaining unhybridized p orbital participates in the π system that delocalizes the unpaired electron and gives the N–O bonds partial double-bond character.
Advanced MO descriptions treat the unpaired electron as residing in a π* (antibonding) orbital delocalized over the molecule.
Formal charge = valence − nonbonding − ½ × bonding. In the O=N–O contributor: nitrogen 5 − 2 − ½×6 = +2 (counting the unpaired e⁻ in the nonbonding count); double-bonded O: 6 − 4 − ½×4 = 0; single-bonded O: 6 − 6 − ½×2 = −1. The resonance average distributes charge, but the key point is that NO2 is an electron-deficient radical, not a closed-shell octet species.
Nitrogen dioxide is polar. The N–O bonds are polar (oxygen more electronegative than nitrogen), and the molecule is bent, so the bond dipoles do not cancel.
The resonance-delocalized odd electron gives the two N–O bonds nearly equal length and polarity, but the bent geometry ensures a net molecular dipole. NO2 therefore has a measurable permanent dipole, consistent with its polar, reactive character.
"NO2 is the instructive exception: 17 electrons means it cannot obey the octet rule, so it is a radical. The unpaired electron sits in a delocalized π* orbital, the geometry stays bent, and that bent shape is why the polar N–O bonds sum to a net dipole."
Interpretation consistent with radical-molecule VSEPR/MO treatments of NO2.This polarity and radical character make NO2 a strong oxidizer and a reactive air pollutant. Its tendency to dimerize to nonpolar N2O4 at lower temperatures is a classic equilibrium that interconverts a bent polar monomer with a planar nonpolar dimer.
| Species | Formula | Central Atom | Electron Count | Molecular Geometry | Polarity |
|---|---|---|---|---|---|
| Nitrogen dioxide | NO2 | N | 17 (radical) | Bent | Polar |
| Nitrous oxide | N2O | central N | 16 | Linear | Polar |
| Dinitrogen tetroxide | N2O4 | N | 34 | Planar (dimer) | Nonpolar |
| Sulfur dioxide | SO2 | S | 18 | Bent | Polar |
| Nitrate ion | NO3− | N | 24 | Trigonal planar | Ionic species |
O–N=O with one lone pair and one unpaired electron on N (resonance). 17 valence electrons make it an odd-electron radical.
NO2 is polar. Its bent geometry prevents the polar N–O bond dipoles from canceling.
The O–N–O angle is about 134°. The unpaired electron repels less than a full lone pair, widening the angle beyond the SO2 value.
Nitrogen is sp2-hybridized. Three domains (2 bonds + lone-pair/unpaired domain) correspond to trigonal-planar hybrids.
The CAS number is 10102-44-0. PubChem and NIST identify this with NO2 (dimer N2O4 is 10544-72-6).
Technical scope: Lewis structures, VSEPR and hybridization are model-based descriptions for chemical communication; NO2 is a genuine free radical and is best described by molecular-orbital theory (unpaired electron in a delocalized π* orbital). CAS 10102-44-0 identifies NO2; it exists in temperature-dependent equilibrium with N2O4 (10544-72-6) and is toxic, corrosive and a regulated pollutant.
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