Orbital Overlap Concept

The Orbital Overlap Concept

The formation of a covalent bond is not simply a matter of two atoms coming close together. There is a specific physical picture that explains why a bond forms and how strong it will be. This picture is the orbital overlap concept.

When two hydrogen atoms approach each other, their atomic orbitals — the 1s orbitals that contain their single electrons — begin to interact. At a certain distance, these orbitals do not just touch; they partially merge or interpenetrate. This partial merging of atomic orbitals is called overlapping. During this overlap, the two electrons, one from each atom, pair up with opposite spins. The result is a covalent bond.

The key idea is that the bond forms because of this overlap. The greater the extent of overlap between the two atomic orbitals, the stronger the resulting covalent bond. A small overlap gives a weak bond; a large overlap gives a strong bond.

Where This Sits on the Energy Curve

This overlap picture is the orbital-level view of the potential-energy story of Section 4.5 (Fig. 4.8): partial interpenetration of the two 1s orbitals is what carries the system down to its minimum energy at the equilibrium bond length, and pushing the atoms closer than that distance drives the energy sharply back up.

The Role of Electron Spin

The orbital overlap concept also explicitly states the condition for the electrons involved. The bond forms by the pairing of electrons present in the valence shell having opposite spins. This is a direct consequence of the Pauli exclusion principle. If two electrons with the same spin tried to occupy the same overlapping region, they would violate this principle, and the system's energy would increase dramatically — no bond would form. Only electrons with opposite spins can occupy the same spatial region (the overlapping orbitals) and form a stable bond.

Summary of the Orbital Overlap Concept

1. A covalent bond forms when atomic orbitals of two atoms overlap (partially interpenetrate). …