VSEPR and Bond Hybridization

VSEPR theory

• Valence Shell Electron Pair Repulsion Theory

• Molecules arrange themselves so that electron domains (bonds and lone pairs) are as far apart as possible to minimize repulsion

Electron domain

• A region around a central atom where electrons are located

• Each single, double, or triple bond = 1 domain

• Each lone pair = 1 domain

Electron geometry

• The shape based on the arrangement of all electron domains (bonding + lone pairs)

• Determines angles between domains

Molecular geometry

• The actual shape of the molecule, based only on the positions of atoms, not lone pairs

• Lone pairs affect shape but aren’t “seen” in molecular geometry

Linear geometry

• Electron domains: 2

• Electron geometry: Linear

• Molecular geometry: Linear

• Bond angle: 180°

• Example: CO₂

Trigonal planar geometry

• Electron domains: 3, no lone pairs

• Electron geometry: Trigonal planar

• Molecular geometry: Trigonal planar

• Bond angle: 120°

• Example: BF₃

Bent (3 electron domains)

• Electron domains: 3 (2 bonding + 1 lone pair)

• Electron geometry: Trigonal planar

• Molecular geometry: Bent

• Bond angle: <120° (lone pair pushes bonds closer)

• Example: SO₂

Tetrahedral geometry

• Electron domains: 4, no lone pairs

• Electron geometry: Tetrahedral

• Molecular geometry: Tetrahedral

• Bond angle: 109.5°

• Example: CH₄

Trigonal pyramidal

• Electron domains: 4 (3 bonding + 1 lone pair)

• Electron geometry: Tetrahedral

• Molecular geometry: Trigonal pyramidal

• Bond angle: ~107° (lone pair repels more than bonds)

• Example: NH₃

Bent (4 electron domains)

• Electron domains: 4 (2 bonding + 2 lone pairs)

• Electron geometry: Tetrahedral

• Molecular geometry: Bent

• Bond angle: ~104.5°

• Example: H₂O

Hybridization

The mixing of atomic orbitals (s, p, d) to form new hybrid orbitals that match the observed molecular geometry

sp hybridization

• Mixing: one s + one p orbital → two sp orbitals

• Electron domains: 2

• Geometry: Linear

• Remaining unhybridized p orbitals form π bonds

• Example: CO₂ (carbon)

sp² hybridization

• Mixing: one s + two p orbitals → three sp² orbitals

• Electron domains: 3

• Geometry: Trigonal planar

• One unhybridized p orbital forms π bond

• Example: BF₃, C₂H₄ (ethene)

sp³ hybridization

• Mixing: one s + three p orbitals → four sp³ orbitals

• Electron domains: 4

• Geometry: Tetrahedral

• Examples: CH₄, NH₃, H₂O

How to determine hybridization

• Count the number of electron domains around the central atom:

• 2 domains → sp

• 3 domains → sp²

• 4 domains → sp³

Sigma bond

• Formed by end-to-end overlap of orbitals along the bond axis

• Present in ALL bonds (single, double, triple)

• Stronger than π bonds

Pi bond

• Formed by sideways overlap of unhybridized p orbitals

• Only in double and triple bonds

• Weaker than σ bond

• Prevents rotation around double bond

Double bond

• In a double bond (like C=C or C=O):

• One sigma bond (from sp²–sp² or sp²–p overlap)

• One pi bond (from p–p sideways overlap)

Triple bond

• In a triple bond (like N≡N or C≡C):

• One sigma bond (from sp–sp overlap)

• Two pi bonds (from two sets of p–p overlaps at right angles)

Why hybridization happens

• Because pure s and p orbitals don’t point in the right directions for observed bond angles

• Hybrid orbitals do, they align perfectly with VSEPR-predicted geometries