bonding orbital
/ BOND-ing OR-bih-tal /
Picture two ripples on a pond meeting in step, crest lining up with crest. Where they overlap, the water piles higher than either ripple alone. Electron waves do the same thing. When two atomic orbitals come together in phase, the electron density swells in the space between the two nuclei. A bonding orbital is a molecular orbital of exactly this kind — one whose electrons collect between the atoms and glue them together.
Why does that gluing happen? An electron sitting between two positive nuclei is attracted to both at once, and it also screens the nuclei from repelling each other. That shared electron density lowers the molecule's energy below the energy of the two separate atoms. So a bonding orbital always sits lower in energy than the atomic orbitals it was built from. Putting electrons into it is favorable, like a ball rolling into a valley; the deeper the valley, the stronger the bond.
Bonding orbitals are the reason molecules form at all. Counting how many electrons fall into bonding orbitals versus antibonding ones gives the bond order, the molecular orbital theory's measure of bond strength. If electrons only ever filled bonding orbitals, every pair of atoms would stick; the existence of antibonding orbitals is what stops, for example, two helium atoms from bonding.
In H2, both electrons occupy the bonding sigma orbital, where the electron density is densest right between the two protons. That central pile of negative charge holds the two nuclei together and is why H2 is a stable molecule but two lone H atoms are not.
Bonding orbitals concentrate electron density between nuclei and lie below the parent atomic orbitals.
Filling a bonding orbital lowers energy, but only the difference between bonding and antibonding electrons matters: that net count, not the bonding electrons alone, sets whether and how strongly a bond forms.