Sometimes, there are several correct Lewis structures for a given molecule. Ozone () is one example. The compound is a chain of three oxygen atoms, and minimizing the charges while giving each atom an octet of electrons requires that the central oxygen atom form a single bond with one terminal oxygen and a double bond with the other terminal oxygen.
When drawing the Lewis structure, the choice of placement for the double bond is arbitrary, and either choice is equally correct. The multiple correct ways of drawing the Lewis structure are called the resonance forms.
Based on the resonance forms, a beginning chemistry student might wonder if ozone has bonds of two different lengths, since single bonds are generally longer than double bonds. However, the ozone molecule is perfectly symmetrical, with bonds that are the same length. None of the resonance forms represent the true structure of the molecule. Rather, the negative charge of the electrons that would form a double bond are delocalized, or distributed evenly across the three oxygen atoms. The true structure is a composite, with bonds shorter than what would be expected for single bonds, but longer than the expected double bonds.
The resonance hybrid for ozone is found by identifying the multiple resonance structures for the molecule.
Thus for the two structures (I and II) shown above constitute the canonical structures or resonance structures and their hybrid (i.e. the III structure) represents the structure of more accurately. Resonance is represented by a double-headed arrow between the resonance structures, as illustrated above.
The resonance hybrid is more stable than its canonical forms, i.e. the actual compound (hybrid) is at a lower energy state than its canonical forms. Resonance stability increases with increased number of resonance structures.
The difference in the experimental and calculated energies is the amount of energy by which the compound is stable. This difference is known as resonance energy or delocalisation energy.
All resonance structures are not equivalent. The following rules help determine whether or not a resonance structure will contribute significantly to the hybrid structure.
Rules of Resonance
Rule 1: The most significant resonance contributor has the greatest number of full octets (or if applicable, expanded octets).
Rule 2: The most significant resonance contributor has the fewest atoms with formal charges.
Rule 3: If formal charges cannot be avoided, the most significant resonance contributor has the negative formal charges on the most electronegative atoms, and the positive formal charges on the least electronegative atoms.
Rule 4: The most significant resonance contributor has the greatest number of covalent bonds.
Rule 5: If a pi bond is present, the most significant resonance contributor has this pi bond between atoms of the same row of the periodic table (usually carbon pi bonded to boron, carbon, nitrogen, oxygen, or fluorine).
Rule 6: Aromatic resonance contributors are more significant than resonance contributors that are not aromatic.