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I'm really struggling to understand how to apply the concept of electronegativity when predicting the bond type in a molecule like SO2. My textbook says it's polar covalent, but when I look at the difference on the Pauling scale, it seems like it should be closer to ionic, and I'm getting confused about where that line actually is.

That SO2 has S–O bonds with a difference around 0.86 on the Pauling scale makes it polar covalent. The line between polar covalent and ionic isn’t a hard cutoff; ionic bonds usually involve much larger differences and a transfer of electron density, often in solids. In molecules you’re still sharing electrons, just unevenly.

I looked at the actual bonding: the S–O bond lengths sit shorter than a typical single bond and closer to a double bond, which fits a resonance picture with S=O like characters. The molecule is bent, so the two polar bonds don’t cancel their dipoles, giving a net dipole. All of that points to covalent with polarity, not ion transfer.

Practically, the polar covalent label is about the nature of the bond, not the whole story. The geometry means the two dipoles add to give a net molecular polarity, but there isn’t full electron transfer.

Partial charges appear on S and O in the resonance forms; S ends up slightly positive, O slightly negative. That distribution supports a covalent framework with resonance rather than an ionic bond.

Sometimes it helps to remember the boundary is context dependent: gas‑phase molecules behave differently from salts. The idea of a sharp line is a simplification, and SO2 sits in that cloudy region where you share electrons but with significant polarity.

Do you think the issue is the threshold itself or the bigger point that resonance and geometry change how we classify bonds, and should we focus on percent ionic character instead of a single label?

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