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When more than one valid Lewis structure can be drawn for a molecule, the true structure is a resonance hybrid — a weighted average of all contributors, with electrons delocalized over multiple atoms rather than fixed in one structure. Formal charge (charge assigned to each atom assuming equal electron sharing) identifies the most stable resonance contributor: structures with formal charges closest to zero, and with negative formal charge on the most electronegative atom, are most significant. Resonance explains equal bond lengths in species like benzene and carbonate.
Draw all resonance structures for ozone, carbonate, nitrate, and benzene. Calculate formal charges for each contributor and rank stability. Connect the concept of delocalization to the observed equal bond lengths in benzene (all bonds intermediate between single and double).
You already know how to draw Lewis structures — assigning electrons to atoms so that each achieves an octet (or duet for hydrogen). But sometimes you can draw more than one perfectly valid Lewis structure for the same molecule, and those structures differ only in where you place the double bonds or lone pairs. Each of these is called a resonance structure (or resonance contributor), and the real molecule is not any single one of them. It is a resonance hybrid — a blend of all contributors, the way a mule is a hybrid of a horse and a donkey, not something that flickers between the two.
Consider the carbonate ion, CO₃²⁻. You can draw three Lewis structures, each placing the double bond on a different oxygen. If one of those structures were "the" structure, you would expect one short C=O bond and two longer C–O bonds. But experiments show all three bonds are identical in length — intermediate between a single and a double bond. That is the hybrid in action: the electrons are delocalized across all three C–O bonds simultaneously, spread out rather than pinned to one location. The same logic explains why benzene's six C–C bonds are all the same length, midway between single and double.
Formal charge is the bookkeeping tool that tells you which resonance structures matter most. To calculate it, take the number of valence electrons an atom "should" have (from its group number), subtract its lone-pair electrons, and subtract half of its bonding electrons. The result is the formal charge on that atom in that particular resonance structure. Two rules then rank the contributors: structures where every atom has a formal charge of zero (or as close to zero as possible) are more significant, and when negative formal charge is unavoidable, it should sit on the more electronegative atom. A structure with a negative charge on carbon and a positive charge on oxygen is a less important contributor than one with negative charge on oxygen.
Resonance and formal charge work together to predict molecular behavior. Delocalization stabilizes molecules — spreading charge over more atoms lowers energy. That is why the carboxylate ion (RCO₂⁻) is far more stable than an alkoxide (RO⁻): the negative charge is shared between two oxygens rather than concentrated on one. When you encounter new molecules, drawing all reasonable resonance structures and evaluating their formal charges will tell you where the electron density actually sits, which bonds are stronger or weaker than a single Lewis structure suggests, and which sites are most reactive.