Q. \( \text{n}_2 \) Lewis structure: polar or nonpolar?

To decide whether an \( \mathrm{N_2} \) Lewis structure is polar or nonpolar, we use two key ideas:

1) What is the electron (Lewis) structure telling you about bonding?
A Lewis structure for \( \mathrm{N_2} \) shows a nitrogen–nitrogen double bond, with both nitrogens completing their valence shells.

2) What matters for polarity is molecular geometry and bond dipoles?
Even if a bond has some polarity, a molecule is overall polar only if those bond dipoles do not cancel out.

Step 1: Write the Lewis structure idea for \( \mathrm{N_2} \)

\( \mathrm{N} \) has \(5\) valence electrons. With two nitrogen atoms total:

\[
\mathrm{N_2:}\quad 2 \times 5 = 10 \text{ valence electrons}
\]

Each nitrogen wants an octet (8 electrons). A standard Lewis structure places a double bond between the two nitrogens and leaves a lone pair on each nitrogen.

That means the bonding is:

\[
\mathrm{N{=}N}
\]

Important detail: \( \mathrm{N_2} \) is a diatomic molecule (two atoms only).

Step 2: Determine whether the molecule has a bond dipole

The molecule would have a bond dipole only if the bonded atoms have different electronegativities.

Here, both atoms are nitrogen, so the electronegativity values are identical.

Therefore, the \( \mathrm{N{-}N} \) bond is effectively nonpolar because there is no electronegativity difference to create a net dipole in one direction.

Step 3: Check whether any dipoles could cancel

For diatomic molecules with identical atoms (like \( \mathrm{N_2} \)), any possible bond dipoles would be equal in magnitude and opposite in direction, so they cancel (and in this case, the bond dipole is already zero by symmetry).

So the molecule has:

\[
\text{Net dipole moment} = 0
\]

Final conclusion

\( \mathrm{N_2} \) (with its Lewis structure \( \mathrm{N{=}N} \)) is nonpolar.