NO2 Lewis Structure
Nitrogen dioxide · The brown gas in photochemical smog, produced by high-temperature combustion.
The NO2 Lewis structure has 17 valence electrons, drawn as one N–O bond and one N=O bond, with five lone pairs in total. The central N is sp2 hybridized, which makes it a bent molecule with a bond angle of about 134 degrees. NO2 is polar: the N-O bond dipoles do not cancel in this shape, leaving a net dipole. There are two equivalent resonance structures.
- 17 valence e−
- bent
- about 134° bond angle
- sp2
- polar
Odd number of valence electrons: this is a free radical, so one electron stays unpaired.
Calculated properties
| Total valence electrons | 17 |
|---|---|
| Bonding electrons | 6 (3 shared pairs) |
| Nonbonding electrons | 11 (5 lone pairs + 1 unpaired electron) |
| Lone pairs on N | 0 |
| Electron domains (steric number) | 3 |
| VSEPR notation | AX2E1 |
| Electron geometry | trigonal planar |
| Molecular geometry | bent |
| Bond angle | about 134° |
| Hybridization | sp2 |
| Polarity | polar |
| Formal charges | N: +1, O: -1 |
| Resonance structures | 2 |
| Molar mass | 46.005 g/mol |
How to draw the NO2 Lewis structure
Step through it. The bar shows how much of the 17-electron budget is spent at each stage - it never grows, which is the whole constraint.
-
Skeleton. Join every connected pair with one bond. 2 bonds spends 4 electrons.
-
Outer octets. Spend electrons as lone pairs on the outer atoms until each has eight. 1 are still unplaced.
-
Centre. Everything still unplaced goes on N, which now has 5 electrons around it.
-
Multiple bonds. The central atom is still short of an octet, so one lone pair on the outer atoms moves in to be shared. The electron count does not change - the same electrons are just counted by both atoms now.
The reasoning at each step
-
Count the valence electrons
Add up the outer-shell electrons every atom brings. That total is the budget for the whole structure - every line and every dot has to come out of it, and nothing may be added.
5 (N) + 2 x 6 (O) = 17 valence electrons -
Choose the central atom
N is the least electronegative atom here, so it takes the centre. Hydrogen is never central: it can only form one bond.
-
Join everything with single bonds first
Every connection starts as one shared pair. Draw them all before worrying about double bonds - the arithmetic in the next step is what tells you where the double bonds have to go.
2 bonds x 2 = 4 electrons used, 13 left -
Work out how far short you are
Add up what every atom still needs to fill its shell - eight electrons for most atoms, two for hydrogen - minus what its single bonds already give it. Compare that with the electrons you have left. The difference decides everything that follows. This molecule has an odd electron count, so one atom has to settle for seven electrons rather than eight.
needs 15, has 13, short by 2 -
Turn the shortfall into multiple bonds
Being short means atoms have to share more. Every pair that moves from a lone pair into a bond counts twice - once for each atom - so a shortfall of 2 is covered by 1 extra shared pair. That is where the double bond comes from.
N=O2 -
Fill in the lone pairs
Every electron not in a bond sits as a lone pair on an atom. Each atom takes exactly what it needs to finish its shell - there is no choice left at this point.
O1: 3, O2: 2 (10 electrons) -
Check the central atom
N ends up with 7 electrons, short of an octet, which is correct for this molecule.
N: 6 bonding + 0 nonbonding + 1 unpaired = 7 -
Check the formal charges
Formal charge is valence electrons, minus lone-pair electrons, minus the number of bonds. They are not real charges, but they have to add up to the overall charge on the species, and a structure that keeps them small is the better one.
N: 5 - 1 - 3 = +1 ; O1: 6 - 6 - 1 = -1 -
Work out the shape
Count the electron domains on N: 2 bonded groups plus one unpaired electron. A double or triple bond still counts as one domain, because it points in one direction. That gives trigonal planar electron geometry; ignore the lone pairs and the atoms themselves sit in a bent arrangement. A single unpaired electron repels less than a full lone pair, so the angle opens up past the value a lone pair would give.
steric number 3 -> sp2 -> bent, bond angle about 134 -
Decide whether it is polar
The N-O bond dipoles do not cancel in this shape, leaving a net dipole.
dipoles do not cancel -> polar
Resonance structures
More than one drawing gives the same electron count and the same formal charges, and no single one of them is the real molecule. NO2 is an average of the 2 structures below, which is why bonds that look different here are actually identical in the real molecule.
Is NO2 polar or nonpolar?
NO2 is polar. The N-O bond dipoles do not cancel in this shape, leaving a net dipole.
| Bond | Electronegativity difference | Character |
|---|---|---|
| N–O | 0.40 | polar covalent |
Common questions
How many valence electrons does NO2 have?
NO2 has 17 valence electrons. 6 of them are in bonds and 11 sit in lone pairs, with one electron left unpaired.
What is the molecular geometry of NO2?
NO2 is a bent molecule. The central N has 2 bonded groups, a steric number of 3, which gives trigonal planar electron geometry and a bent molecule.
What is the bond angle in NO2?
The bond angle in NO2 is about 134 degrees. A single unpaired electron repels less than a full lone pair, so the angle opens up past the value a lone pair would give.
What is the hybridization of NO2?
The central N in NO2 is sp2 hybridized. Steric number 3 means 3 orbitals have to be mixed, which is exactly what sp2 gives you.
Is NO2 polar or nonpolar?
NO2 is polar. The N-O bond dipoles do not cancel in this shape, leaving a net dipole.
Does NO2 have resonance structures?
Yes. NO2 has two equivalent resonance structures. The real molecule is not any one of them - it is an average, so every bond that differs between the drawings is really the same length in the actual molecule.