NO₃⁻ Lewis Structure (Nitrate Ion)
The NO₃⁻ Lewis structure represents the nitrate ion with 24 valence electrons. Nitrogen is the central atom, bonded to three oxygen atoms. Each resonance contributor contains one N=O double bond and two N–O single bonds, with a +1 formal charge on nitrogen and −1 on each single-bonded oxygen. Three equivalent resonance structures produce three identical N–O bonds. The ion has a trigonal planar geometry with bond angles of approximately 120°.
| Valence electrons | 24 |
|---|---|
| Resonance forms | 3 |
| Molecular geometry | Trigonal planar |
| Bond angle | 120° |
| VSEPR | AX₃ |
| Average N–O bond order | 1⅓ |
| Overall charge | −1 |
How to Draw the NO₃⁻ Lewis Structure
Draw the nitrate ion step by step: count 24 valence electrons, put nitrogen in the center, complete all three oxygen octets, turn one oxygen lone pair into an N=O double bond, then draw the two other equivalent resonance contributors.
Step 1Count the valence electrons
Nitrogen contributes 5 valence electrons and three oxygens contribute 18. The −1 charge adds one electron: 5 + 3(6) + 1 = 24 valence electrons. All 24 must appear in the finished Lewis drawing.
Atom Count Valence e⁻ Total N 1 5 5 O 3 6 18 Total 24 Step 2Place nitrogen in the center
Nitrogen is less electronegative than oxygen, so it belongs in the center. Draw three N—O single bonds. The three bonds use 6 electrons, leaving 18 electrons to assign.
Step 3Complete the oxygen octets
Place three lone pairs on each outer oxygen. Those nine lone pairs use the remaining 18 electrons. Every oxygen now has an octet, but nitrogen has only six electrons from its three single bonds.
Step 4Form one N=O double bond
Move one lone pair from an oxygen into its N—O bond. The new N=O double bond gives nitrogen an octet. The finished contributor has one double bond, two single bonds, and no lone pairs on nitrogen.
Step 5Verify charges and draw all resonance forms
In this contributor nitrogen has formal charge +1, the double-bonded oxygen has 0, and each single-bonded oxygen has −1. The charges sum to −1. Put the N=O double bond on each oxygen in turn to make three equivalent contributors; the resonance arrows describe one delocalized nitrate ion, not three different molecules.
- 24 valence electrons placed
- Nitrogen octet satisfied (two single bonds and one double bond)
- Each oxygen octet satisfied
- Formal charges sum to −1
- All three equivalent resonance forms shown
NO₃⁻ Valence Electrons
Nitrogen contributes 5 valence electrons. The 3 oxygen atoms contribute 3 × 6 = 18 valence electrons. After adjusting for the ion charge, NO₃⁻ starts with 24 valence electrons to place. Of those, 8 fill the bonds and 16 sit as lone pairs (O: 2, 3, 3).
| nitrogen: 5 valence electrons |
| 3 × oxygen: 3 × 6 = 18 |
| Total: 5 + 18 = 24 |
NO₃⁻ Resonance Structures
NO₃⁻ has three equivalent resonance contributors. Every contributor has one N=O double bond and two N–O single bonds; only the position of the double bond and the two −1 oxygen formal charges changes. The real nitrate ion is a resonance hybrid, not a molecule switching among three separate drawings or a chemical equilibrium. Electron density is delocalized, so all three N–O bonds have the same length and an average bond order of 4/3 (about 1.33).
Resonance hybrid
Average N–O bond order = (2 + 1 + 1) / 3 = 4/3 ≈ 1.33
The three equivalent contributors make all three N–O bonds equivalent in the resonance hybrid.
NO₃⁻ Lewis Structure with Formal Charges
Formula
Formal charge = valence e⁻ − nonbonding e⁻ − ½ × bonding e⁻
| Atom | Calculation | Formal charge |
|---|---|---|
| N | 5 − 0 − ½(8) = +1 | +1 |
| O | 6 − 4 − ½(4) = 0 | 0 |
| O | 6 − 6 − ½(2) = −1 | −1 |
| O | 6 − 6 − ½(2) = −1 | −1 |
The formal charges sum to -1, matching the −1 ion. Smaller formal charges mark the preferred structure.
Central N: 5 − 0 − ½(8) = +1. Nitrogen has no nonbonding electrons and eight bonding electrons.
Double-bonded O: 6 − 4 − ½(4) = 0.
Each single-bonded O: 6 − 6 − ½(2) = −1.
The formal charges add to −1, the net charge of nitrate. In the other two resonance contributors, the double-bonded oxygen and the two single-bonded oxygens exchange positions. Nitrogen's +1 formal charge is not the same thing as its +5 oxidation state in nitrate.
NO₃⁻ Octet Rule and Lone Pairs
NO₃⁻ obeys the octet rule in each resonance drawing. Nitrogen is a period-2 element, so it cannot expand its octet: one double bond plus two single bonds gives it exactly eight bonding electrons. Drawing three N=O double bonds would put 12 bonding electrons around nitrogen and is not a valid nitrate Lewis structure, even though it appears to remove formal charges.
NO₃⁻ Molecular Geometry and Bond Angle
| Electron geometry | Trigonal planar |
|---|---|
| Molecular geometry / shape | Trigonal planar |
| VSEPR notation | AX₃ |
| Bond angle | 120° |
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. NO₃⁻ has 3 bonding domains and no lone pairs (AX₃), for 3 electron domains in total.
NO₃⁻ has trigonal planar molecular geometry because nitrogen has three bonding domains and no lone pairs (AX₃).
NO₃⁻ is AX₃: nitrogen is A and the three N–O bonding regions are X₃. With three electron domains and no lone pairs on nitrogen, electron-domain geometry and molecular geometry are both trigonal planar. The three N–O bonds lie in one plane at approximately 120°.
Why Is NO₃⁻ Trigonal planar?
A double bond counts as one electron domain in VSEPR, just as a single bond does. Each resonance contributor therefore gives nitrogen three bonding domains, not four. With no lone pair on nitrogen, those domains spread 120° apart in a trigonal planar arrangement.
NO₃⁻ Bond Angle
The O–N–O bond angles in nitrate are approximately 120°. Three bonding domains repel one another equally around nitrogen, and nitrogen has no lone pair to compress any angle. Resonance makes all three N–O bonds equivalent in the real ion.
Charge Distribution and Symmetry of NO₃⁻
Nitrate is a trigonal planar polyatomic ion with three equivalent N–O bonds in its resonance hybrid. The negative charge is spread over the oxygen atoms rather than permanently fixed on the two oxygens that carry −1 formal charges in any one resonance drawing. Because NO₃⁻ has a net charge of −1, it should not be labelled simply polar or nonpolar by applying the neutral-molecule dipole-cancellation rule.
NO₃⁻ Hybridization
Nitrogen in NO₃⁻ is sp²-hybridized.
Nitrogen has three N–O bonding domains and no lone pairs. In a common introductory description, three electron domains correspond to sp² hybridization and a trigonal planar arrangement. The remaining unhybridized p orbital participates in the delocalized π bonding represented by the three equivalent resonance contributors.
NO₃⁻ vs NO₂⁻: What's the Difference?
Nitrate and nitrite are both negatively charged nitrogen oxyanions with resonance, but adding one oxygen changes the electron-domain count at nitrogen. Nitrate has three bonding domains and no nitrogen lone pair, so it is trigonal planar. Nitrite has two bonding domains and one nitrogen lone pair, so it is bent.
Both ions have a −1 overall charge, but nitrate has three equivalent contributors and an average N–O bond order of 4/3; nitrite has two contributors and an average N–O bond order of 1.5.
| Property | NO₃⁻ | NO₂⁻ (nitrite ion) |
|---|---|---|
| Total valence electrons | 24 | 18 |
| Total oxygen atoms | 3 | 2 |
| Resonance contributors | 3 | 2 |
| Lone pairs on N | 0 | 1 |
| Molecular geometry | Trigonal planar | Bent |
| Bond angle | 120° | ≈115° |
| Average N–O bond order | 1⅓ | 1.5 |
| Overall charge | −1 | −1 |
NO₃⁻ Lewis Structure Summary
| Chemical name | Nitrate ion |
|---|---|
| Total valence electrons | 24 |
| Central atom | Nitrogen |
| N=O | Double bond |
| N–O | 2 × single bond |
| Lone pairs on nitrogen | 0 |
| Lone pairs on outer atoms | O: 2, 3, 3 |
| Electron geometry | Trigonal planar |
| Molecular geometry | Trigonal planar |
| VSEPR | AX₃ |
| Bond angle | 120° |
| Hybridization | sp² |
| Polarity | Charged ion |
| Formal charge | see formal charge table |
| Resonance | Yes |
| Overall charge | −1 |
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Related Lewis Structures
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Frequently Asked Questions
What is the Lewis structure of NO₃⁻?
The nitrate ion has nitrogen in the center with one N=O double bond and two N–O single bonds in each Lewis contributor. It is drawn in brackets with an overall −1 charge. The double bond can be placed on any of the three oxygens, giving three equivalent resonance structures.
How many valence electrons are in NO₃⁻?
NO₃⁻ has 24 valence electrons: 5 from nitrogen, 18 from three oxygens, and 1 added for the −1 charge. The calculation is 5 + 3(6) + 1 = 24.
How many resonance structures does NO₃⁻ have?
NO₃⁻ has three equivalent resonance contributors. Each has one N=O double bond and two N–O single bonds, but the double bond can be assigned to any one of the three oxygens.
What are the formal charges in NO₃⁻?
In one nitrate resonance contributor, nitrogen has formal charge +1, the double-bonded oxygen has 0, and the two single-bonded oxygens each have −1. These formal charges add to the ion's overall −1 charge.
Why are all N–O bonds in nitrate equivalent?
The three resonance contributors are equivalent, so the π bonding and negative charge are delocalized over all three oxygens. The real nitrate ion is a resonance hybrid with three equal N–O bonds, each with an average bond order of 4/3.
What is the molecular geometry of NO₃⁻?
NO₃⁻ has trigonal planar molecular geometry. Nitrogen has three bonding domains and no lone pairs (VSEPR AX₃), so the three N–O bonds lie in one plane about 120° apart.
Does NO₃⁻ obey the octet rule?
Yes. In each standard contributor nitrogen has one double bond and two single bonds, which gives it eight bonding electrons. Each oxygen also has an octet. Nitrogen cannot be drawn with three N=O double bonds because that would give this period-2 atom 12 electrons.