Understanding the Context

The nitrite ion (NO₃⁻) consists of one nitrogen atom bonded to three oxygen atoms, with a net negative charge of −1. A proper Lewis structure illustrates covalent bonds and lone pairs, showing how valence electrons are arranged while obeying the octet rule.

Step-by-Step Construction:

1. Total Valence Electrons
   Nitrogen (N): 5
   Each Oxygen (O): 6 → 3 × 6 = 18
   Add 1 for the −1 charge → Total: 5 + 18 + 1 = 24 valence electrons

2. Central Atom Selection
   Nitrogen is less electronegative than oxygen and is typically the central atom.

Key Insights

3. Bond Formation
   Draw single bonds: N—O between N and each O (using 6 electrons).
   This leaves 24 − 6 = 18 electrons.

4. Distributing Remaining Electrons
   Each oxygen must complete its octet. Place a lone pair (2 electrons) on each outer O first:
   3 × 2 = 6 electrons used
   Remaining: 18 − 6 = 12 electrons

5. Forming Double Bonds
   Use the remaining pairs to form double bonds starting with the least electronegative atom.
   Convert some lone pairs from outer O into lone pairs, then share one electron pair to form a double bond with N.

6. Resonance Structures
   The actual structure is a resonance hybrid:
   Three equivalent resonance forms exist where the double bond shifts among the three O atoms.

Final NO₃⁻ Lewis Structure:

Final Thoughts

[O=N–O]⁻
| [O–N–O]⁻ ↔ [O–N=O]⁻

In reality, the nitrogen shares a double bond with one oxygen and single bonds with the other two, which collectively represent resonance stabilization.

The Resonance Revelation: Why It Matters

A common misconception is drawing a single fixed structure. The true nature of NO₃⁻ is a resonance hybrid, meaning electron density is spread across all three N–O bonds. This delocalization explains why all N–O bonds are of equal length—intermediate between single and double bonds—unlike typical polar covalent molecules.

Understanding this transformation: