Resonance

Sometimes, applying the five-step Lewis-structure recipe honestly to a single molecule still yields MORE THAN ONE equally valid structure -- not because one is wrong, but because the true electronic structure cannot be captured by any single dot-and-line diagram. The carbonate ion, CO₃²⁻, is the textbook example.

Working through CO₃²⁻. Label the three oxygens O(A), O(B) and O(C) for reference. Total valence electrons = [1×4 (C)] + [3×6 (O)] + [2 (the ion's own 2− charge)] = 4 + 18 + 2 = 24. Distributing these electrons by the usual method, and then moving one lone pair from O(A) to complete carbon's octet, gives a structure with one C=O(A) double bond and two C-O single bonds (to O(B) and O(C)), with the 2− charge marked on the ion overall (Fig 10.14(a)).

But this choice of O(A) to carry the double bond was arbitrary -- two further, EQUALLY VALID Lewis structures can be drawn simply by moving the double bond to O(B) or to O(C) instead. All three structures have the atoms in exactly the same relative positions; they differ only in WHICH bonding and lone pairs sit where (Fig 10.14(b)). Such structures -- identical in atomic arrangement, differing only in electron placement -- are called resonance structures (or canonical structures), and the phenomenon itself is called resonance.

The resonance hybrid. Experimentally, all three carbon-oxygen bonds in the carbonate ion are found to be EQUIVALENT (same bond length, intermediate between a C-O single and a C=O double bond). The actual structure of CO₃²⁻ is therefore said to be a resonance hybrid -- an average of the three resonance/canonical forms, NOT an actual physical oscillation between them (the ion does not flip from one structure to another and back; it exists, at all times, as the single hybrid). No single Lewis structure can be drawn for this hybrid, but a composite picture -- partial (fractional) bonds to all three oxygens, each carrying a fractional −2/3 charge -- gives a reasonable qualitative sense of it (Fig 10.14(c)). …