Resonance Structure
Resonance structures (also called canonical forms or contributing structures) are two or more valid Lewis structures for the same molecule that differ only in the placement of electrons (not atoms), used collectively to represent the actual electron distribution. No single resonance structure accurately depicts the molecule; the true structure is a hybrid (weighted average) of all contributing structures, with electron density delocalised across multiple bonds. The concept is essential for understanding the stability of molecules like benzene, ozone, and carbonate ion, where observed bond lengths are intermediate between single and double bonds.
Resonance Structures of Common Molecules
| Molecule | Number of Resonance Structures | Observed Bond Length | Notes |
|---|---|---|---|
| Ozone (O₃) | 2 | O–O: 128 pm (between single 148 pm and double 121 pm) | Bent geometry, equivalent bonds |
| Carbonate (CO₃²⁻) | 3 | C–O: 129 pm | Three equivalent C–O bonds |
| Benzene (C₆H₆) | 2 (Kekulé) | C–C: 140 pm (between 154 pm and 134 pm) | Fully delocalised π system |
| Nitrate (NO₃⁻) | 3 | N–O: 124 pm | Trigonal planar, equivalent bonds |
| SO₂ | 2 | S–O: 143 pm | Equivalent S–O bonds |
Interactive Tools
Khan Academy – Resonance Structures
Step-by-step guide to drawing and evaluating resonance structures
Open Tool
Wikimedia Commons, CC BY-SA
Related Terms
Formal Charge
Formal charge is a bookkeeping tool used to assess the distribution of electrons in a Lewis structure by assigning each atom a hypothetical charge, assuming electrons in covalent bonds are shared equally. It is calculated as the number of valence electrons minus the number of lone pair electrons minus half the number of bonding electrons for each atom. The sum of formal charges equals the overall charge of the molecule or ion, and the most stable Lewis structure is generally the one with the smallest formal charges and any negative formal charges on the more electronegative atoms.
Pi Bond
A pi bond (π bond) is a covalent bond formed by the lateral (side-by-side) overlap of unhybridised p orbitals above and below the internuclear axis. Pi bonds are always formed in addition to an existing sigma bond, making up the second bond in a double bond and the second and third bonds in a triple bond. Unlike sigma bonds, pi bonds restrict rotation around the bond axis, which is critical for cis-trans isomerism in alkenes.
Sigma Bond
A sigma bond (σ bond) is the strongest type of covalent bond, formed by the direct head-on overlap of atomic orbitals along the internuclear axis. It is the first bond formed between two atoms in any covalent bond and allows free rotation around the bond axis. Sigma bonds are present in all single, double, and triple bonds and are responsible for the overall framework and shape of molecules.
From Latin "resonare" (to resound, echo back), referring to the oscillation or "resonance" between equivalent structures. The concept was introduced by Linus Pauling in the 1930s based on quantum mechanical valence bond theory, though the terminology was inspired by earlier German chemical literature.