Hybridization
Hybridization is a theoretical concept in chemistry describing the mixing of atomic orbitals of similar energy within the same atom to form new hybrid orbitals of equivalent energy and shape, oriented to minimize electron repulsion. Developed by Linus Pauling in 1931, hybridization explains molecular geometry that cannot be accounted for by simple orbital overlap — for example, carbon's four equivalent C–H bonds in methane despite having distinct 2s and 2p orbitals. The type of hybridization (sp, sp², sp³, sp³d, sp³d²) determines bond angles, molecular geometry, and the presence of pi bonds.
Worked Example
Problem
Determine the hybridization of carbon in ethene (C₂H₄) and explain what type of bonds are present.
Solution
Step 1: Draw the Lewis structure of C₂H₄: each C is double-bonded to the other (C=C) and single-bonded to 2 H atoms. Step 2: Count electron domains around each carbon: 1 double bond (counted as 1 domain) + 2 single bonds = 3 electron domains. Step 3: 3 electron domains → sp² hybridization (1 s + 2 p orbitals mix). Step 4: The 3 sp² orbitals point at 120° apart, forming 3 sigma (σ) bonds: C–C σ bond + 2 C–H σ bonds. Step 5: The remaining unhybridized p orbital on each carbon overlaps sideways to form 1 pi (π) bond. Step 6: Total bonds in C=C: 1 σ + 1 π = double bond.
Answer
Each C in C₂H₄ is sp² hybridized (3 σ bonds at 120°, trigonal planar geometry), and the C=C double bond consists of 1 σ bond + 1 π bond.
Types of Hybridization and Corresponding Molecular Geometry
| Hybridization | Orbitals Mixed | Geometry | Bond Angle | Example |
|---|---|---|---|---|
| sp | 1s + 1p | Linear | 180° | BeCl₂, C₂H₂ |
| sp² | 1s + 2p | Trigonal planar | 120° | BF₃, C₂H₄ |
| sp³ | 1s + 3p | Tetrahedral | 109.5° | CH₄, NH₃ |
| sp³d | 1s + 3p + 1d | Trigonal bipyramidal | 90°/120° | PCl₅ |
| sp³d² | 1s + 3p + 2d | Octahedral | 90° | SF₆ |
Interactive Tools
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Related Terms
sp3 Hybridization
sp³ hybridization occurs when one s orbital and three p orbitals of the same atom mix to form four equivalent sp³ hybrid orbitals, each oriented at approximately 109.5° from one another in a tetrahedral arrangement. This type of hybridization is exhibited by carbon in saturated compounds like methane (CH₄), by nitrogen in ammonia (NH₃), and by oxygen in water (H₂O), though the presence of lone pairs in NH₃ and H₂O slightly distorts the ideal tetrahedral angle. sp³ hybridization is fundamental to understanding the three-dimensional structure of organic molecules, including all alkanes and the carbon backbone of biological macromolecules.
Molecular Geometry
Molecular geometry (or molecular shape) refers to the three-dimensional spatial arrangement of atoms within a molecule, determined by the positions of the atoms — not the lone pairs — around the central atom. The geometry is predicted using VSEPR theory or hybridization models and directly influences physical properties such as polarity, reactivity, phase of matter, colour, magnetism, and biological activity. Common geometries include linear, bent, trigonal planar, trigonal pyramidal, tetrahedral, and octahedral.
VSEPR Theory
Valence Shell Electron Pair Repulsion (VSEPR) theory is a model used to predict the three-dimensional geometry of molecules based on the principle that electron pairs in the valence shell of a central atom repel each other and arrange themselves as far apart as possible to minimize repulsion. The theory considers both bonding pairs and lone pairs, with lone pairs exerting greater repulsive force than bonding pairs, which distorts ideal bond angles. VSEPR theory was developed by Ronald Gillespie and Ronald Nyholm in 1957 and remains one of the most useful and accessible tools for predicting molecular shape.
From Latin "hybrida" (offspring of two different species) — the concept of "mixing" orbital types to form new ones. Formally introduced by Linus Pauling in 1931 in his valence bond theory papers.