9.10: Resonance

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 do indeed 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: if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by the various Lewis structures. 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 forms. 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 a 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.

Always remember that 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.

George Wheland, one of the pioneers of resonance theory, used a historical analogy to describe the relationship between resonance forms and resonance hybrids. A medieval traveler, having never before seen a rhinoceros, described it as a hybrid of a dragon and a unicorn because it had many properties in common with both. Just as a rhinoceros is neither a dragon sometimes nor a unicorn at other times, a resonance hybrid is neither of its resonance forms at any given time.

Like a rhinoceros, it is a real entity that experimental evidence has shown to exist. It has some characteristics in common with its resonance forms, but the resonance forms themselves are convenient, imaginary images (like the unicorn and the dragon).

This text is adapted from Openstax, Chemistry 2e, Section 7.4: Formal Charges and Resonance.

Most molecules and ions can be represented using unique Lewis structures. However, certain compounds can be shown by multiple, equally valid Lewis structures. 

Consider the Lewis structure for sulfur trioxide. The single bonds between each oxygen and the central sulfur atoms satisfy the octet for 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. Since any of the three oxygens can form the double bond with sulfur, three different Lewis structures can be drawn. 

These multiple Lewis structures are called resonance structures, where the skeletal structures remain the same, but electrons are distributed differently.

All three structures are valid and equivalent representations of the molecule, yet all are non-existent in nature. The actual structure does not oscillate between the resonance structures but is a hybrid or an average of the three Lewis structures, which can be measured in bond lengths. 

In sulfite, the length of a sulfur-oxygen single bond is 1.51 angstrom, while in sulfur trioxide, the sulfur-oxygen bond length is 1.42 angstrom. Thus, in the hybrid molecule, the bond length is an intermediate between single and double bonds. 

In hybrid molecules, electrons participating in double bonds or lone pairs are often delocalized over multiple bonds or atoms, meaning they are not stationary on one particular atom. The delocalization reduces the potential energy of electrons resulting in stabilization called resonance stabilization. 

Resonance is also observed for aromatic compounds such as benzene. Benzene is a hexagonal carbon-ring with one hydrogen bonded to each of the carbon atoms, and alternating single and double bonds between the carbon atoms. Based on the location of carbon-carbon double bonds, benzene can have two resonance structures. 

Recall that double bonds are usually shorter than single bonds. However, all the carbon-carbon bonds in benzene have equal bond lengths, which are intermediate between carbon-carbon single and double bonds. 

Hence, benzene exists as a resonance hybrid and can be represented as a hexagon with a circle inside. The circle indicates that benzene is a blend of two resonance structures, and the double bonds cannot be localized to any two specific carbon atoms. 

• Resonance - A concept in which electrons are delocalized over three or more centres.
• Resonance Structures - Different Lewis structures for the same molecule.
• Resonance Hybrid - Actual structure of the molecule represented as the average of the resonance forms.
• Lewis Structure - A graphical representation of the molecular structure of a compound.
• Delocalization - The spread of electrons over more than two atoms.

• Define Resonance – An understanding of electron delocalization (e.g., resonance in chemistry).
• Contrast Lewis Structure vs Resonance Structures – Identify the difference (e.g., bonding in NO₂^-).
• Explore Resonance of Nitrite Ion – Visualize how resonance structures are drawn (e.g., Lewis structure of NO₂^-).
• Explain Resonance Hybrid – Better understanding of the actual molecular structure.
• Apply Delocalization in Context – Understand how electrons spread over different atoms.

• How does the concept of resonance explain electron delocalization?
• What is resonance hybrid and how is it different from resonance structures?
• What is delocalization and how does it affect molecular structure?

• Students – Understand resonance, how it affects molecular structure, and its importance in chemistry.
• Educators – Get a clear understanding of resonance and its importance in teaching related topics in chemistry.
• Researchers – Important for understanding molecular structures in research work.
• Chemistry Enthusiasts – Offers insights into a complex concept in chemistry and enhances knowledge.