N₂ Lewis Structure: Triple Bond and Lone Pairs

N₂ Lewis structure

$$\mathord{:}\mathrm{N}\equiv\mathrm{N}\mathord{:}$$

Each colon is one lone pair. The three lines represent three shared electron pairs. Thus, each nitrogen has one lone pair and a triple bond; the Lewis bond order is 3.

Count and place the electrons

Each nitrogen contributes five valence electrons, so the molecule has 10 valence electrons. A single N–N bond uses two electrons. Giving both atoms three lone pairs would require fourteen electrons in total, which exceeds the available ten. Sharing three pairs instead gives both atoms an octet without adding electrons.

The final accounting is six electrons in the triple bond plus four in the two lone pairs: 6 + 4 = 10. Around either nitrogen, the six bonding electrons and its two nonbonding electrons make an octet. Count a shared pair toward both atoms' octets, but only once in the molecule's total.

Check formal charge

Formal charge equals the neutral atom's valence-electron count minus its nonbonding electrons minus half its bonding electrons. For either nitrogen:

$$\mathrm{FC}(\mathrm{N})=5-2-\frac{6}{2}=0.$$

Both formal charges are zero, consistent with a neutral molecule. In the usual localized bonding description, the triple bond contains one σ bond and two π bonds. The Lewis drawing records connectivity and electron pairs; the positions of its dots do not specify measured orbital shapes.