1.9:
Resonance and Hybrid Structures
Most molecules and ions are represented by a single Lewis structure. However, in particular compounds, some electrons are delocalized over multiple bonds or atoms rather than localized to a specific bond or atom. These compounds can be represented accurately with multiple Lewis structures.
Consider the Lewis structure for sulfur trioxide. The single bonds between each oxygen and the central sulfur atoms satisfy the octet for the oxygen atoms. However, to reach a full octet for the sulfur, an additional bond must be formed between sulfur and one of the oxygen atoms.
These multiple Lewis structures of sulfur trioxide are called resonance structures or contributing structures. The spatial positions of all of the component atoms remain the same; however, the valence electrons are distributed differently.
These double-headed arrows can be thought of as commas and should not be confused with resonance structures being in equilibrium. In these molecules, the actual structure is the weighted average or a hybrid of its resonance structures.
Resonance structures are identified by ‘electron pushing’, or transforming lone pairs into bonds and vice versa, as denoted with curved arrows. This maps the delocalized areas.
Such delocalization results in resonance stabilization—that is, a molecule with a lower potential energy than that of any theoretical non-delocalized structure.
If a contributing structure is lower in energy than another, it more closely resembles the actual molecular structure. Certain preferences are used to estimate the relative energies of various contributing structures.
Firstly, structures in which all atoms have filled valence shells are more stable. Secondly, structures with a greater number of covalent bonds are more stable. Thirdly, structures that minimize formal charges are more stable.
Finally, structures carrying negative charges on more electronegative atoms and positive charges on less electronegative atoms are more stable.
1.9:
Resonance and Hybrid Structures
According to the theory of resonance, if two or more Lewis structures with the same arrangement of atoms can be written for a molecule, ion, or radical, the actual distribution of electrons is an average of that shown by the various Lewis structures.
Resonance Structures and Resonance Hybrids
The Lewis structure of a nitrite anion (NO2−) may actually be drawn in two different ways, distinguished by the locations of the N–O and N=O bonds.
If nitrite ions contain a single and a double bond, the two bond lengths are expected to be different. A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms. However, experiments show that both N–O bonds in NO2− have the same strength and length, and are identical in all other properties. It is not possible to write a single Lewis structure for NO2− in which nitrogen has an octet, and both bonds are equivalent.
Instead, the concept of resonance is used: The actual distribution of electrons in each of the nitrogen–oxygen bonds in NO2− is the average of a double bond and a single bond.
The individual Lewis structures are called resonance structures. The actual electronic structure of the molecule (the average of the resonance forms) is called a resonance hybrid of the individual resonance forms. A double-headed arrow between Lewis structures indicates that they are resonance forms.
The carbonate anion, CO32−, provides a second example of resonance.
- • One oxygen atom must have a double bond to carbon to complete the octet on the central atom.
- • All oxygen atoms, however, are equivalent, and the double bond could form from any one of the three atoms. This gives rise to three resonance forms of the carbonate ion.
- • Since three identical resonance structures can be written, the actual arrangement of electrons in the carbonate ion is known to be the average of the three structures.
- • Again, experiments show that all three C–O bonds are exactly the same.
A molecule described as a resonance hybrid never possesses an electronic structure described by either resonance form. It does not fluctuate between resonance forms; rather, the actual electronic structure is always the average of that shown by all resonance forms.
This text is adapted from Openstax, Chemistry 2e, Chapter 7.4 Formal Charges and Resonance.