NO₂ Lewis Structure
Nitrogen Dioxide (neutral NO₂ radical)
NO₂ is an odd-electron radical with 17 total valence electrons. Nitrogen is the central atom. The Lewis structure is drawn as two equivalent resonance forms with one N=O double bond, one N–O single bond, formal charges of +1 on N and −1 on the single-bonded oxygen, and one unpaired electron on nitrogen. The molecule is bent, polar, and has an O–N–O angle near 134°.
Looking for NO₂⁻?
Neutral NO₂ (nitrogen dioxide) is not the nitrite ion NO₂⁻. The neutral radical has 17 valence electrons; nitrite has 18 and is not a radical. NO₂⁻ (nitrite) Lewis structure →Lewis / Resonance
Two equivalent resonance contributors. The bent shape is shown beside it.
Molecular Geometry
Bent shape · ~134° · AX₂E• · 1 unpaired electron on N
| Total valence electrons | 17 |
|---|---|
| Central atom | Nitrogen (N) |
| Odd-electron species | Yes |
| Unpaired electrons | 1 |
| N=O | Double bond |
| N–O | Single bond |
| Lone pairs on N | 0 |
| Lone pairs on O | 2 |
| Electron geometry | Trigonal planar |
| Molecular geometry | Bent |
| VSEPR | AX₂E• |
| Bond angle | ~134° |
| Polarity | Polar |
| N hybridization | sp² |
| Resonance | Yes |
| Formal charges | see formal charge table |
How to Draw the NO₂ Lewis Structure
Draw the Lewis structure of nitrogen dioxide step by step. Count the 17 valence electrons, put nitrogen in the center, finish the oxygen octets, form one N=O double bond, then draw the matching resonance form. Because 17 is odd, one electron stays unpaired.
Step 1Count the valence electrons
Nitrogen brings 5 valence electrons and each oxygen brings 6. With two oxygens that is 5 + 12 = 17 electrons to place — and NO₂ is neutral, so there is no charge to add or remove. Because 17 is odd, NO₂ cannot give every atom a conventional paired-electron octet. NO₂ is an odd-electron radical.
Atom Count Valence e⁻ Total N 1 5 5 O 2 6 12 Total 17 Step 2Put nitrogen in the center
Nitrogen is less electronegative than oxygen and can form more than one bond, so it sits in the middle. Connect the atoms with single bonds first: O—N—O. Those two bonds use 4 of the 17 valence electrons, leaving 13 to assign as lone pairs, multiple bonds, and the unpaired electron.
Step 3Complete the oxygen octets
Give each oxygen enough lone pairs to reach eight electrons around it. That uses 12 of the remaining electrons (three lone pairs on each oxygen). One electron is left over — it cannot form a pair, so it stays as an unpaired electron on nitrogen.
Step 4Form one N=O double bond
Nitrogen still has only five electrons around it (two single bonds plus the unpaired electron). Move one lone pair from either oxygen into the N–O link to make a double bond: O=N–O. Nitrogen then has seven electrons — as close to an octet as an odd-electron count allows — with formal charges +1 on N, 0 on the double-bonded oxygen, and −1 on the single-bonded oxygen.
Step 5Draw the resonance form
The double bond can sit on either oxygen. Swap the N=O and N–O sides to get the second contributor: O=N–O ↔ O–N=O. The two drawings are equivalent resonance forms of the same hybrid, not two different molecules. Formal charges swap with the bonds; the unpaired electron stays on nitrogen; the total electron count stays 17.
- 17 valence electrons placed
- Nitrogen central with one unpaired electron
- One N=O double bond and one N–O single bond (or the matching resonance form)
- Formal charges sum to 0
- Both resonance forms shown
NO₂ Valence Electrons
Nitrogen contributes 5 valence electrons, and 2 × oxygens contribute 2 × 6 = 12. Added together, NO₂ starts with 17 valence electrons to place, and as a neutral molecule there is no charge to add or remove. Of those, 6 fill the bonds and 10 sit as lone pairs (O: 2, 3), and 1 remains unpaired on N.
| nitrogen: 5 valence electrons |
| 2 × oxygen: 2 × 6 = 12 |
| Total: 5 + 12 = 17 |
Why Does NO₂ Have an Odd Electron?
| N | 5 × 1 = 5 |
|---|---|
| O | 6 × 2 = 12 |
| Total | 17 |
Add the valence electrons atom by atom: nitrogen contributes 5, and the two oxygens contribute 6 each, for 5 + 12 = 17. An odd total means at least one electron cannot be paired. In the usual NO₂ Lewis drawings that unpaired electron sits on nitrogen, so nitrogen ends with seven electrons instead of eight. That is why NO₂ is called an odd-electron radical — and why it differs from NO₂⁻ (18 electrons, every electron paired) and from molecules like CO₂ or SO₂ where every electron is paired.
The finished drawing keeps 1 unpaired electron on the central atom. No rearrangement of bonds can pair every electron when the total count is odd.
NO₂ Resonance Structures
NO₂ has two equivalent resonance contributors. In one form the left oxygen is double-bonded and the right oxygen is single-bonded (with formal charge −1); in the other form the bonds and the oxygen formal charges swap. Nitrogen keeps formal charge +1 and the unpaired electron in both drawings. The real electronic structure is not a molecule flipping back and forth between the two pictures — it is a resonance hybrid in which both N–O bonds are equivalent and share the same average bond order. This is not the same pair used for nitrite: NO₂⁻ has 18 electrons, no unpaired electron, and different formal charges.
NO₂ Lewis Structure with Formal Charges
Formula
Formal charge = valence e⁻ − nonbonding e⁻ − ½ × bonding e⁻
| Atom | Calculation | Formal charge |
|---|---|---|
| N | 5 − 1 − ½(6) = +1 | +1 |
| O | 6 − 4 − ½(4) = 0 | 0 |
| O | 6 − 6 − ½(2) = −1 | −1 |
The formal charges sum to 0, matching the neutral molecule. The formal charges sum to the overall charge of the molecule; the unpaired electron is counted as nonbonding in the formula.
Formal charge is a counting tool, not a measured partial charge and not the overall molecular charge. For each atom: formal charge = valence electrons − nonbonding electrons − ½(bonding electrons). Nonbonding electrons include lone-pair electrons and the unpaired radical electron.
On one resonance contributor O=N–O•: the double-bonded oxygen has 2 lone pairs (4 nonbonding) and one double bond (4 bonding), so 6 − 4 − ½(4) = 0. Nitrogen has 1 unpaired electron (1 nonbonding) and one double bond plus one single bond (6 bonding), so 5 − 1 − ½(6) = +1. The single-bonded oxygen has 3 lone pairs (6 nonbonding) and one single bond (2 bonding), so 6 − 6 − ½(2) = −1. The other contributor only swaps which oxygen is which. The formal charges always sum to 0 for neutral NO₂ — charge separation inside a contributor is not the same as a net molecular charge of −1 (that would be NO₂⁻).
If a drawing shows brackets and an overall −1 charge, it is nitrite (NO₂⁻), not neutral nitrogen dioxide. For NO₂ the formal charges inside each contributor cancel to 0 because the molecule itself is neutral.
NO₂ Molecular Geometry
| Electron geometry | Trigonal planar |
|---|---|
| Molecular geometry / shape | Bent |
| VSEPR notation | AX₂E• |
| Bond angle | ~134° |
On the central nitrogen, count electron domains the VSEPR way: each bonded neighbor is one domain (a double or triple bond still counts as one), and each lone pair is one more; an unpaired electron also counts as one electron region. NO₂ has 2 bonding domains and 1 unpaired electron (AX₂E•), for 3 electron domains in total.
Those domains arrange to minimize repulsion around the nitrogen, so O═N─O is bent with a bond angle of ~134°. The unpaired electron on nitrogen still takes space in the electron count, so the electron geometry (trigonal planar) differs from the molecular geometry (bent).
VSEPR notation is often written AX₂E•: A is nitrogen, X₂ are the two oxygen bonding domains, and E• is the unpaired electron on nitrogen treated as one electron region. The electron geometry is trigonal planar because three regions need space around N; the molecular geometry is bent (angular) because only the two oxygens count in the shape. A simple AX₂E estimate is about 120°; the measured gas-phase O–N–O angle is about 134°, wider than SO₂ or H₂O because an unpaired electron takes less space than a full lone pair.
Why Is NO₂ Bent?
NO₂ is not linear like CO₂. Carbon dioxide is AX₂ with no lone pairs and no unpaired electron on carbon, so the two double bonds sit 180° apart. Nitrogen dioxide has two bonding regions plus one unpaired electron on nitrogen. Those three electron regions arrange in a trigonal planar pattern, and the unpaired electron occupies one corner, so the two N–O bonds bend. Name the electron geometry from all regions first; name the molecular geometry from the atoms only — that is why the shape is bent while the electron geometry stays trigonal planar.
NO₂ Bond Angle
The O–N–O bond angle in NO₂ is approximately 134°. That is wider than the ideal 120° of a simple AX₂E molecule and wider than SO₂ (~119°) or H₂O (104.5°), because the unpaired electron on nitrogen repels bonding pairs less strongly than a full lone pair. The comparison below runs from linear CO₂ down to tightly bent water.
| CO₂ | 180° |
|---|---|
| NO₂ | ~134° |
| SO₂ | ~120° (experimental ~119°) |
| H₂O | 104.5° |
Is NO₂ Polar or Nonpolar?
NO₂ is polar.
Each N–O bond is polar: oxygen (electronegativity 3.44) pulls shared electrons harder than nitrogen (3.04), so every bond dipole points toward oxygen. Bond polarity alone does not decide the molecule. If NO₂ were linear like CO₂, the two dipoles would sit farther apart and could cancel more completely. Nitrogen keeps an unpaired electron that still counts as an electron region (AX₂E•), so the shape is bent near 134° instead. In that bent geometry the bond dipoles are no longer opposite on one line — they share a component toward the oxygen side of the bend, those components add, and a net molecular dipole remains. NO₂ is therefore polar. CO₂ has polar X=O bonds but linear AX₂ geometry, so its dipoles cancel and the molecule is nonpolar. Geometry, not bond polarity alone, makes the difference.
NO₂ Hybridization
Nitrogen in NO₂ is sp²-hybridized.
Central nitrogen has three electron regions — two N–O bonding domains plus the unpaired-electron region — so it is described as sp² hybridized. Two hybrids point toward the oxygens; the third holds the unpaired electron. That arrangement matches the trigonal planar electron geometry and the bent molecular shape. It is the same three-region picture used for SO₂ (sp² on S), not the four-region sp³ picture used for H₂O.
NO₂ vs NO₂⁻: What Is the Difference?
Nitrogen dioxide (NO₂) and the nitrite ion (NO₂⁻) share a bent O–N–O framework and both can be drawn with resonance, but they are not the same species. Neutral NO₂ has 17 valence electrons and one unpaired electron; NO₂⁻ has 18 valence electrons and no unpaired electron.
Formal charges follow that count. On each NO₂ resonance contributor, nitrogen is typically +1 and the single-bonded oxygen −1, summing to 0 for the neutral molecule. NO₂⁻ places 18 electrons, so its resonance forms have no unpaired electron and formal charges that sum to −1.
| Property | NO₂ | NO₂⁻ |
|---|---|---|
| Name | Nitrogen dioxide | Nitrite ion |
| Charge | 0 | −1 |
| Valence electrons | 17 | 18 |
| Radical | Yes (1 unpaired e⁻) | No |
| Resonance | Yes | Yes |
| Shape | Bent | Bent |
| Formal charges (one contributor) | N +1, =O 0, –O −1 (sum 0) | N 0, =O 0, –O −1 (sum −1) |
| Bond angle | ~134° | ~115° |
NO₂ Lewis Structure Summary
| Chemical name | Nitrogen dioxide |
|---|---|
| Total valence electrons | 17 |
| Central atom | Nitrogen |
| Odd-electron species | Yes |
| Unpaired electrons | 1 |
| Structure | O═N─O |
| N=O | Double bond |
| N–O | Single bond |
| Lone pairs on nitrogen | 0 |
| Lone pairs on outer atoms | O: 2, 3 |
| Electron geometry | Trigonal planar |
| Molecular geometry | Bent |
| VSEPR | AX₂E• |
| Bond angle | ~134° |
| Hybridization | sp² |
| Polarity | Polar |
| Formal charge | see formal charge table |
| Resonance | Yes |
| Overall charge | 0 |
Try Another Lewis Structure
Enter another formula to draw its Lewis structure — bonds, lone pairs, formal charges, and the steps behind them.
Related Lewis Structures
Common next structures from general chemistry — each link opens that molecule so you can compare geometry, polarity, or the drawing steps.
Frequently Asked Questions
What is the Lewis structure of NO₂?
The NO₂ Lewis structure is a resonance pair with nitrogen central, one N=O double bond, one N–O single bond, an unpaired electron on nitrogen, and formal charges +1 on N and −1 on the single-bonded oxygen. The two contributors swap which oxygen is double-bonded. NO₂ has 17 valence electrons and is an odd-electron radical — not the nitrite ion NO₂⁻.
How many valence electrons does NO₂ have?
NO₂ has 17 valence electrons: 5 from nitrogen and 6 from each oxygen (2 × 6 = 12). Because 17 is odd, one electron stays unpaired and NO₂ is a radical.
Why does NO₂ have an odd electron?
The atom valence totals are odd: 5 + 6 + 6 = 17. No arrangement of pairs can use an odd number of electrons without leaving one unpaired. In the usual drawings that unpaired electron sits on nitrogen.
Does NO₂ have resonance structures?
Yes. NO₂ has two equivalent resonance forms that differ only in which oxygen carries the double bond (and the −1 formal charge). The actual molecule is a hybrid with equivalent N–O bonds, not a frozen single/double pair.
What is the molecular geometry of NO₂?
NO₂ has a bent molecular geometry. Nitrogen has two bonding domains plus an unpaired electron counted as an electron region (AX₂E•), so the electron geometry is trigonal planar and the visible shape is bent with an O–N–O angle near 134°.
Is NO₂ polar or nonpolar?
NO₂ is polar. The N–O bonds are polar, and the bent shape prevents the bond dipoles from canceling, so the molecule has a net dipole moment.
What is the difference between NO₂ and NO₂⁻?
NO₂ is neutral nitrogen dioxide with 17 valence electrons and one unpaired electron (a radical). NO₂⁻ is the nitrite ion with charge −1, 18 valence electrons, and no unpaired electron. Both are bent and both show resonance, but only NO₂ has an odd electron count.
What is the bond angle of NO₂?
The O–N–O bond angle in NO₂ is about 134°. The unpaired electron on nitrogen counts as an electron region (bent AX₂E• geometry) but repels less than a full lone pair, so the angle sits between linear CO₂ (180°) and more tightly bent molecules like SO₂ or H₂O.