NO2 Lewis Structure: Draw the Dot and Cross Diagram

If you’ve ever tried drawing the Lewis structure for nitrogen dioxide only to end up with one stray electron, you’re not alone. That unpaired electron is what makes NO₂ a radical — and it changes everything about its shape, polarity, and how you represent it on paper.

Total valence electrons in NO₂: 17 · Bond angle: ~134° · Dipole moment: 0.316 D · Molecular shape: Bent

Key facts about NO₂

Six data points capture the essential chemistry of nitrogen dioxide at a glance.

The implication: every structural property of NO₂ stems from its odd‑electron count — a rare case that forces students to move beyond the octet rule.

How to write Lewis dot structure for NO₂?

Step‑by‑step process

1. Count total valence electrons: 5 from nitrogen and 12 from two oxygens gives 17 (YouTube tutorial on NO₂ Lewis structure).
2. Place nitrogen in the center because it is the least electronegative atom (YouTube tutorial on NO₂ Lewis structure).
3. Connect the two oxygens to nitrogen with single bonds, using 4 electrons (2 per bond).
4. Distribute the remaining 13 electrons as lone pairs on the oxygens. Each oxygen needs an octet; after giving each oxygen three lone pairs (12 electrons), you have one electron left — place it on nitrogen (YouTube video tutorial on NO₂). The central nitrogen ends up with only 7 electrons, violating the octet rule.
5. To reduce formal charge, convert one N–O single bond to a double bond by moving a lone pair from an oxygen to the bond, creating two major resonance forms (Ace Organic Chemistry Online tutorial).

Determining total valence electrons

• Nitrogen (Group 15): 5 valence electrons
• Each oxygen (Group 16): 6 valence electrons × 2 = 12
• Total = 17 valence electrons (Ace Organic Chemistry Online tutorial)

Because 17 is odd, one electron remains unpaired — making NO₂ a radical species (YouTube video tutorial on NO₂).

Placing atoms and electrons

Connect the two oxygens to nitrogen with single bonds. That uses 4 electrons (2 per bond). Of the remaining 13 electrons, distribute them as lone pairs on the oxygens. Each oxygen needs an octet; after giving each oxygen three lone pairs (12 electrons), you have one electron left — place it on nitrogen (YouTube video tutorial on NO₂). The central nitrogen ends up with only 7 electrons, violating the octet rule.

Resonance structures

To reduce formal charge, one of the N–O single bonds can be converted to a double bond by moving a lone pair from an oxygen to the bond. This creates two major resonance forms (Ace Organic Chemistry Online tutorial). The unpaired electron remains on the nitrogen in both forms, delocalised over the oxygens. A common teaching model represents NO₂ with resonance between O=N–O• and •O–N=O (Ace Organic Chemistry Online tutorial).

The pattern: you cannot draw a single structure that obeys the octet rule; you must accept resonance as the model.

Is NO₂⁻ polar?

The nitrite ion, NO₂⁻, has 18 valence electrons — one more than neutral NO₂ (University of Maryland educational resource). This extra electron fills the octet and gives the ion a bent geometry with a bond angle near 115°.

Molecular geometry of NO₂⁻

With nitrogen at the centre and one double bond to an oxygen, the ion is bent. Two resonance structures are possible: O=N–O⁻ and O⁻–N=O (Studocu student notes).

Electronegativity difference

Oxygen is more electronegative than nitrogen, so each N–O bond is polar. The bent shape means the bond dipoles do not cancel — NO₂⁻ has a net dipole and is therefore polar (Studocu student notes).

Dipole moment direction

The dipole points from the partially positive nitrogen toward the more negative oxygen atoms. The asymmetry of the lone pair on nitrogen also contributes to the overall polarity.

What this means: both NO₂ and NO₂⁻ are polar, but the magnitude differs because of the extra electron and slightly different bond angle.

Is NO₂⁺ polar or nonpolar?

Removing one electron from NO₂ yields NO₂⁺ (the nitronium ion), which has only 16 valence electrons — an even number. The loss of the unpaired electron drastically changes the geometry.

Linear geometry of NO₂⁺

With no odd electron, the central nitrogen can form two double bonds (N=O) using sp hybridisation. The molecule is linear with a bond angle of 180° (MadSci science Q&A).

Symmetry leads to nonpolar

Because the two N=O bonds are identical and point in opposite directions along a straight line, their dipole moments cancel exactly. NO₂⁺ has no net dipole — it is nonpolar.

The trade‑off: linear geometry makes NO₂⁺ a strong electrophile, used in nitration reactions. Its nonpolar nature also means it behaves differently in solution compared to NO₂ and NO₂⁻.

What is the Criss Cross method of NO₂?

The “criss‑cross” method is a technique for writing empirical formulas of ionic compounds by swapping the charges of the cation and anion. Its application to NO₂ is limited because NO₂ is a covalent molecule.

Applicability to covalent compounds

Criss‑cross works for ionic salts like NaCl or MgO. For covalent compounds, the formula is determined by the valence electron count and the desire to achieve stable octets — not by charge neutralisation (YouTube tutorial on NO₂ Lewis structure).

Steps for NO₂

If one tried to apply criss‑cross to nitrogen (charge +5 in its highest oxidation state) and oxygen (–2), the resulting formula would be N₂O₅ — not NO₂. The method fails because it does not account for the odd‑electron nature of NO₂.

Limitations

Always check whether the compound is ionic or covalent before using criss‑cross. For NO₂, the Lewis structure approach is the correct tool.

The implication: criss-cross is not a shortcut for covalent molecules like NO₂ — stick to valence electron counting.

What is the difference between NO₂ and NO₂⁻?

Four key distinctions separate neutral nitrogen dioxide from the nitrite ion. The table below summarises them side by side.

One extra electron, one fundamental difference: NO₂ is a radical, while NO₂⁻ is a stable closed‑shell ion.

The pattern: adding one electron turns a reactive radical into a stable ion while preserving bent geometry and polarity — but the bond angle narrows and the charge changes reactivity entirely.

Clarity: confirmed facts vs. what remains unclear

Expert perspectives on NO₂

“Nitrogen dioxide is a classic example of an odd‑electron molecule that cannot be described by a single Lewis structure — resonance is required.”

— Chemistry: The Central Science (university textbook)

“The bent geometry of NO₂ means the bond dipoles do not cancel, making it polar despite the symmetrical formula.”

— Chemistry Steps (tutorial site)

Summary: what this means for chemistry students

NO₂ is more than a footnote in the octet rule. Its 17 valence electrons force a resonance description, a bent geometry, and a net dipole that sets it apart from both its anion and cation cousins. For students learning Lewis structures, the lesson is clear: always count electrons first. When the count is odd, expect a radical — and reach for resonance.

Related reading: NO2 Dot and Cross Diagram: Lewis Structure Explained

Frequently asked questions