Bonding Orbital
Definition and meaning of Bonding Orbital in chemistry.
A bonding orbital is a molecular orbital formed when atomic orbitals combine in-phase (constructively), producing increased electron density between two nuclei and an energy lower than that of the original atomic orbitals.
In more detail
When atomic orbital wavefunctions overlap with the same phase, they interfere constructively, concentrating electron probability directly between the nuclei. This buildup of electron density attracts both positively charged nuclei toward it, lowering the overall energy of the system relative to the separated atoms and holding the atoms together. Because bonding orbitals are lower in energy, electrons fill them preferentially according to the aufbau principle; the more bonding orbitals are occupied (relative to antibonding orbitals), the higher the bond order and the stronger the bond. Every bonding orbital has a companion antibonding orbital, formed from out-of-phase overlap, that is higher in energy and contains a node between the nuclei.
Key facts
| Field | Physical Chemistry |
|---|---|
| Formation | constructive (in-phase) overlap of atomic orbitals |
| Relative energy | lower than the combining atomic orbitals |
| Counterpart | antibonding orbital (denoted with an asterisk, e.g., σ*, π*) |
In molecular hydrogen (H2), the two hydrogen 1s atomic orbitals overlap in-phase to form a single σ1s bonding molecular orbital. Both electrons occupy this lower-energy orbital, giving a bond order of 1 and accounting for the stability of the covalent H–H bond.
Frequently asked questions
How does a bonding orbital differ from an antibonding orbital?
A bonding orbital arises from in-phase overlap of atomic orbitals, is lower in energy, and increases electron density between the nuclei, stabilizing the bond. An antibonding orbital arises from out-of-phase overlap, is higher in energy, and has a node between the nuclei that weakens or breaks the bond if occupied.
Are all bonding orbitals the same type?
No. Bonding orbitals are classified as sigma (σ) or pi (π) depending on their symmetry: σ orbitals form from head-on overlap along the bond axis, while π orbitals form from side-on overlap of parallel p orbitals, as seen together in double bonds.