1. σ and π bonds
σ and π bonds In these images, atomic orbitals and bonds are depicted as line drawings indicating shape and as isosurfaces, regions of space enclosing a defined fraction of the electron density. (A) When two s-orbitals overlap, a σ-bond is formed. In σ bonds, the atomic orbitals overlap ‘head on’ and there is electron density between the two nuclei in the bond. As a result, these bonds are typically strong compared with π bonds between the same elements. (B) In a π bond, two p-orbitals overlap laterally, and the nuclei lie in a plane in which there is no electron density. There is less overlapping between orbitals in a π bond than in a σ bond and so these bonds are typically weaker. It should be noted that a double bond, consisting of a σ and π bond is stronger than a single bond between the same two elements as the strength of both the σ and π components are included in the bond strength. (C) When a bond is formed between carbon and oxygen, there will be more electron density located near the oxygen atom, as illustrated here in the isosurface for a carbon–oxygen π bond.
Amanda L. Jonsson et al. Essays Biochem. 2017;61:
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