Chemistry · Ch 4 — Chemical Bonding and Molecular Structure
Directional Properties of Bonds
Directional Properties of Bonds
Why Shape Matters in Covalent Bonds
A covalent bond forms when atomic orbitals overlap. In a diatomic molecule like H₂, the bond is simply the result of two 1s orbitals overlapping along the line joining the nuclei. But in polyatomic molecules — CH₄, NH₃, H₂O — the geometry of the whole molecule becomes just as important as the bonds themselves.
Why does CH₄ have a tetrahedral shape with all H–C–H angles exactly 109.5°? Why is NH₃ pyramidal rather than planar? The valence bond theory answers these questions by introducing two key ideas: the directional nature of atomic orbitals, and the concept of hybridisation. The direction in which a bond points is not arbitrary — it is determined by the shapes and orientations of the orbitals that overlap.
The Core Idea: Overlap Determines Direction
For a bond to form, the overlapping orbitals must have the correct relative orientation. An s orbital is spherical — it has no preferred direction. But a p orbital is dumbbell-shaped, with two lobes pointing along a specific axis (x, y, or z). When a p orbital overlaps with another orbital, the strongest bond forms when the overlap occurs along the axis of the p orbital — that is, when the two nuclei lie on the line through the centres of the overlapping lobes.
This is the origin of directional properties — and also of a famous failure. If oxygen in H₂O bonded through its two half-filled 2p orbitals alone, the two O–H bonds should sit at the 90° angle those p orbitals make with each other. The measured angle is 104.5° — nowhere near 90°. Pure atomic-orbital overlap cannot account for the actual directions of bonds; something more is needed.
The Problem with Simple p-Orbital Overlap
If we try to explain the shapes of CH₄, NH₃, and H₂O using only pure s and p orbitals, we run into trouble. Carbon has the ground-state configuration 1s²2s²2pₓ¹2pᵧ¹ — only two unpaired electrons, which would suggest it forms only two bonds. But methane has four equivalent C–H bonds. Clearly, something else is happening.
The valence bond theory resolves this by proposing that before bonding occurs, the atomic orbitals on the central atom mix or hybridise to form a set of new, equivalent orbitals. These hybrid orbitals have specific directional properties that directly determine the molecular geometry.
The directional properties of bonds in polyatomic molecules are explained by the overlap of hybrid orbitals, not pure atomic orbitals. The geometry of the molecule is a direct consequence of the spatial arrangement of these hybrid orbitals.
How hybridisation resolves this — the full sp³ treatment of CH₄ (109.5°), NH₃ (107°) and H₂O (104.5°) — is developed in Sections 4.6 and 4.6.1.