Topic 2.6 Resonance and Formal Charge: draw resonance structures and use formal charge to select the most reasonable Lewis diagram, and explain how resonance describes delocalised bonding.

What this topic is asking

The College Board (Topic 2.6) wants you to draw resonance structures when more than one valid Lewis diagram exists, to understand the resonance hybrid as the true delocalised structure, and to calculate formal charge and use it to choose the most reasonable Lewis diagram among alternatives.

Resonance

The classic example is the nitrate ion $\text{NO}_3^{-}$, which can be drawn with the N=O double bond in any of three positions. No single drawing is correct: the real ion has all three N-O bonds identical, each one a blend of single and double bond character. The double-headed arrow between resonance structures means "the truth is the average", not "the molecule flips between forms".

Why resonance matters

Resonance solves a contradiction. A single Lewis structure of $\text{NO}_3^{-}$ predicts one short (double) bond and two longer (single) bonds, but experiment finds all three N-O bonds equal, with a length between that of a single and a double bond. The resonance hybrid resolves this: the extra bonding pair is delocalised equally over all three positions, so every bond has the same intermediate length and strength. Delocalisation also lowers the energy, which is why species with resonance are especially stable.

Formal charge

Formal charge is a way to compare competing Lewis structures. The best structure is the one that:

1. gives every atom a formal charge as close to zero as possible, and
2. if a nonzero formal charge is unavoidable, places the negative formal charge on the most electronegative atom.

This is how you decide, for example, the preferred arrangement of atoms or the placement of a double bond when several Lewis diagrams satisfy the octet rule. Formal charge is bookkeeping, not a real charge, but it is a reliable guide to which structure best represents the molecule, and it links forward to acid strength and reactivity in later units.

Try this

Q1. Calculate the formal charge on the central nitrogen in the ammonium ion $\text{NH}_4^{+}$ (four N-H bonds, no lone pairs on N). [1 point]

• Cue. $\text{FC} = 5 - 0 - \tfrac{1}{2}(8) = 5 - 4 = +1$.

Q2. Explain why all three sulfur-oxygen bonds in the sulfite ion, $\text{SO}_3^{2-}$, have the same length. [2 points]

• Cue. The ion is a resonance hybrid; the bonding electrons are delocalised equally over the three S-O positions, so each bond is identical and intermediate between single and double.