A single Lewis structure often fails to accurately represent the true electron distribution within a molecule or ion, especially for species exhibiting delocalized bonding. The concept of resonance addresses this limitation by proposing multiple theoretical structures, known as resonance forms or contributors. These forms collectively describe the molecule’s actual, time-averaged structure, called the resonance hybrid. Resonance forms are imaginary representations, not distinct molecules that rapidly interconvert, but they serve as essential tools for understanding delocalized bonding. To be considered valid, these forms must adhere to a strict set of rules that define the invariant features of the molecule.
Fixed Spatial Arrangement of Atoms
A defining feature shared by all valid resonance forms is the fixed spatial arrangement of the atomic nuclei. The connectivity, or the sigma bond framework, of the molecule or ion must remain completely unchanged across all contributing structures. Only the electrons—specifically pi electrons found in double or triple bonds and non-bonding lone pairs—are permitted to shift their positions.
This rule means that an atom bonded to specific neighbors in one form must be bonded to the exact same neighbors in all other resonance forms. The geometry of the molecule, such as the bond angles and hybridization of the atoms, must also be maintained. If the position of an atom were to change, the resulting structure would represent a different compound, known as a structural isomer or a tautomer, not a resonance form.
For instance, in the carbonate ion (\(\text{CO}_3^{2-}\)), the central carbon atom is always bonded to the three oxygen atoms, and the relative positions of these four nuclei never change. Only the location of the double bond and the corresponding non-bonding electrons on the oxygen atoms are redistributed to create the three equivalent resonance structures. The immobility of the atomic nuclei is the physical basis for this invariant feature.
Conservation of Total Electron Count and Net Charge
The most rigorous accounting principle that all resonance forms share is the conservation of both the total number of valence electrons and the net electrical charge. Since resonance involves only the reorganization of existing electrons, no electrons are ever added to or removed from the system when drawing contributing structures. This means that if a species, such as the nitrate ion (\(\text{NO}_3^{-}\)), has a total of 24 valence electrons, every resonance form proposed for that ion must also contain exactly 24 valence electrons.
The net charge of the entire species must likewise remain constant in every contributor. For the nitrate ion, the overall charge is \(-1\), and the sum of the formal charges on all atoms in any resonance structure must total precisely \(-1\). While the formal charge on individual atoms will frequently change, the algebraic sum of these charges is an invariant property.
Formal charge is a tool used to track the distribution of electrons within a single Lewis structure. The movement of pi bonds and lone pairs causes these localized formal charges to shift from one atom to another. However, the foundational principle is that the total count of electrons and the overall net charge of the molecule or ion cannot be altered by the process of drawing resonance forms.
Consistency in Electron Pairing
A final feature that must be consistent across all resonance forms is the state of electron pairing, specifically the spin multiplicity. The number of paired electrons and, consequently, the number of unpaired electrons must remain the same in every contributing structure. This rule ensures that a stable, closed-shell molecule, where all electrons are paired, cannot have a resonance form that is an open-shell species with an unpaired electron, also known as a free radical.
If the original molecule has all its electrons paired, all of its resonance forms must also be entirely closed-shell. If the original molecule were a free radical, such as the allyl radical, every resonance form must also contain exactly one unpaired electron. The rule prevents the transformation of a radical into a non-radical species, or vice versa, because a change in spin state represents a chemical change, not just electron delocalization.