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Hybridization describes how atomic orbitals mix to form new hybrid orbitals suited for bonding. Common hybridization schemes are sp (linear), sp² (trigonal planar), and sp³ (tetrahedral), each yielding different molecular geometries and bond angles. Hybridization explains why bonding geometry often differs from pure orbital geometry.
Determine the number of electron groups around a central atom, predict hybridization, and sketch the geometry. Verify predictions using molecular models.
From your study of electron configuration, you know that carbon's ground state has two electrons in the 2s orbital and two in separate 2p orbitals — which would suggest carbon should form only two bonds. But carbon routinely forms four equivalent bonds, as in methane (CH₄). Something about the simple atomic orbital picture does not match reality. Hybridization is the model that resolves this discrepancy: it proposes that atomic orbitals on the same atom can mathematically combine — or "mix" — to produce a new set of equivalent hybrid orbitals that are better suited for forming bonds.
The number of orbitals you mix equals the number of hybrid orbitals you get out. In sp³ hybridization, one s orbital and three p orbitals combine to produce four identical sp³ hybrid orbitals, each pointing toward the corner of a tetrahedron with 109.5° angles between them. This is exactly what we observe in methane: four equivalent C–H bonds arranged tetrahedrally. In sp² hybridization, one s orbital mixes with two p orbitals to produce three sp² hybrid orbitals in a trigonal planar arrangement (120° angles), leaving one unhybridized p orbital available for pi bonding — which you already know from your study of sigma and pi bonds. This is the bonding picture in ethylene (C₂H₄), where the double bond consists of one sigma bond (from sp² overlap) and one pi bond (from the leftover p orbitals). In sp hybridization, one s and one p orbital mix to give two sp hybrids pointing in opposite directions (180°), leaving two unhybridized p orbitals for pi bonds — as in acetylene (C₂H₂) with its triple bond.
The practical rule for predicting hybridization is straightforward: count the number of electron groups (bonds plus lone pairs) around the central atom. Two electron groups mean sp, three mean sp², and four mean sp³. A lone pair counts as an electron group just like a bond does — it occupies a hybrid orbital and affects the geometry. This is why ammonia (NH₃) is sp³ hybridized even though it has only three bonds: the lone pair occupies the fourth sp³ orbital, pushing the three N–H bonds into a pyramidal shape rather than a flat trigonal arrangement.
It is important to understand that hybridization is a model, not a physical process that atoms undergo. Atoms do not literally remix their orbitals before bonding. Rather, hybridization is a mathematical description that correctly predicts observed bond angles, bond equivalence, and molecular geometry. Its predictive power is what makes it valuable: given only the molecular formula, you can determine how many electron groups surround each atom, assign hybridization, and predict the three-dimensional shape of the molecule — connecting electron configuration directly to molecular geometry.