Hybridization Rules
Determine hybrid orbital types
Understanding Hybridization
Hybridization is a theoretical concept in valence bond theory that explains molecular geometry by proposing that atomic orbitals mix to form new hybrid orbitals with specific spatial arrangements. Introduced by Linus Pauling in the 1930s, hybridization theory resolved the puzzle of how carbon forms four equivalent bonds in methane despite having only two unpaired electrons in its ground state. The theory proposes that one 2s and three 2p orbitals combine to form four equivalent sp³ hybrid orbitals, each capable of forming a sigma bond with hydrogen atoms in a tetrahedral arrangement at 109.5° angles.
The type of hybridization (sp, sp², sp³, sp³d, sp³d²) correlates directly with the steric number—the sum of bonding electron pairs and lone pairs on the central atom. This steric number approach, combined with VSEPR (Valence Shell Electron Pair Repulsion) theory, allows chemists to predict both electron geometry (determined by hybridization) and molecular geometry (the actual shape considering lone pairs). For example, water has sp³ hybridization (steric number 4), but its molecular geometry is bent rather than tetrahedral because two of the four hybrid orbitals contain lone pairs instead of bonding pairs.
Understanding hybridization is essential for predicting bond angles, explaining molecular shapes, understanding reactivity patterns in organic chemistry, and interpreting spectroscopic data. Different hybridizations produce characteristic bond angles: sp = 180° (linear), sp² = 120° (trigonal planar), sp³ = 109.5° (tetrahedral). In organic chemistry, recognizing that carbon in alkanes uses sp³ hybridization, alkenes use sp², and alkynes use sp hybridization helps explain differences in reactivity, bond strength, and acidity. The concept also extends to coordination chemistry where transition metals use d-orbitals in hybridization schemes like sp³d² (octahedral) and sp³d (trigonal bipyramidal).
Steric Number Method
Steric Number = (bonding pairs) + (lone pairs)
SN = 2 → sp (linear)
Example: CO₂, BeCl₂
SN = 3 → sp² (trigonal planar)
Example: BF₃, CH₂=CH₂
SN = 4 → sp³ (tetrahedral)
Example: CH₄, NH₃, H₂O
SN = 5 → sp³d (trigonal bipyramidal)
Example: PCl₅
SN = 6 → sp³d² (octahedral)
Example: SF₆
Additional Example: SO₂
Step 1: Draw Lewis structure — S has 2 bonding regions (to O atoms) and 1 lone pair
Step 2: Count bonding regions: 2 (double bonds count as ONE region)
Step 3: Count lone pairs on central S: 1
Step 4: Steric number = 2 + 1 = 3
Step 5: SN = 3 → sp² hybridization
Note: Electron geometry is trigonal planar, but molecular geometry is bent (~119°) due to lone pair repulsion.
Answer: sp² hybridization, bent molecular geometry
Key Concepts
Bond Angles & Hybridization
sp = 180°, sp² = 120°, sp³ = 109.5°, sp³d = 90°/120°, sp³d² = 90°. Lone pairs compress bond angles slightly (H₂O: 104.5° instead of 109.5°).
Multiple Bonds
Double/triple bonds count as ONE bonding region. Sigma bonds use hybrid orbitals; pi bonds use unhybridized p orbitals. sp² has one unhybridized p; sp has two.
Organic Chemistry Applications
Alkanes (C-C): sp³. Alkenes (C=C): sp². Alkynes (C≡C): sp. Hybridization affects acidity: sp > sp² > sp³ (more s character = more acidic).
Common Mistakes
Counting multiple bonds as multiple regions
A double bond (C=O) counts as ONE bonding region, not two. Only count the number of atoms bonded, not the number of bonds.
Confusing electron geometry with molecular geometry
Hybridization predicts electron geometry. Molecular geometry considers only atom positions, affected by lone pairs.
Forgetting lone pairs in steric number
BOTH bonding pairs AND lone pairs count toward steric number. NH₃: 3 bonds + 1 lone pair = SN 4 (sp³).
Quick Tips
- Multiple bonds (double, triple) count as ONE bonding region.
- Lone pairs on central atom count toward steric number.
- Hybridization = electron geometry; molecular geometry considers lone pairs.