Valence Bond Theory and Hybridization in Chemical Bonding

Valence Bond Theory and Hybridization

Introduction to Valence Bond Theory

Valence Bond (VB) Theory explains the formation of chemical bonds as the overlap of atomic orbitals from different atoms, resulting in a region of increased electron density between the nuclei. This theory helps describe the geometry and strength of covalent bonds in molecules.

• Valence Bond Theory: Bonds form when atomic orbitals on adjacent atoms overlap and the electrons are shared between the atoms.

• Orbital Overlap: The strength of a covalent bond depends on the extent of overlap between the atomic orbitals.

• Bonding and Antibonding Interactions: Constructive overlap leads to bonding orbitals, while destructive overlap leads to antibonding orbitals.

• Example: The H2 molecule forms when two hydrogen 1s orbitals overlap, creating a bond.

Conventions of the Orbital Overlap Model

The orbital overlap model follows several conventions to describe bonding:

• Each electron is assigned to a specific atomic orbital, following the Pauli exclusion principle.

• The directionality of bonds arises from the orientation of the atomic orbitals involved.

• Only unpaired electrons in atomic orbitals can participate in bonding.

• Bonding occurs when atomic orbitals overlap, and the resulting bond has a specific geometry determined by the orientation of the orbitals.

Hybridization of Atomic Orbitals

Hybridization is the process by which atomic orbitals mix to form new, equivalent hybrid orbitals that are oriented to maximize bonding and minimize electron repulsion. This concept explains the observed shapes of molecules.

• sp3 Hybridization: Mixing one s and three p orbitals forms four sp3 hybrid orbitals, arranged tetrahedrally (109.5° angles).

• sp2 Hybridization: Mixing one s and two p orbitals forms three sp2 hybrid orbitals, arranged trigonal planar (120° angles).

• sp Hybridization: Mixing one s and one p orbital forms two sp hybrid orbitals, arranged linearly (180° angles).

• Example: In methane (CH4), carbon undergoes sp3 hybridization to form four equivalent C–H bonds.

Summary Table: Atomic Orbital Hybridization

Type	Number of Hybrid Orbitals	Geometry	Bond Angle
sp	2	Linear	180°
sp2	3	Trigonal planar	120°
sp3	4	Tetrahedral	109.5°

Multiple Bonds: Sigma (σ) and Pi (π) Bonds

Multiple bonds consist of one sigma (σ) bond and one or more pi (π) bonds:

• Sigma (σ) Bond: Formed by head-on overlap of orbitals (e.g., sp3–sp3 or s–s).

• Pi (π) Bond: Formed by side-on overlap of unhybridized p orbitals.

• Double Bond: One σ bond and one π bond (e.g., in ethene, C2H4).

• Triple Bond: One σ bond and two π bonds (e.g., in ethyne, C2H2).

Bonding in Selected Molecules

• Methane (CH4): Carbon is sp3 hybridized, forming four σ bonds with hydrogen.

• Ethene (C2H4): Each carbon is sp2 hybridized, forming a σ bond framework and a π bond from unhybridized p orbitals.

• Ethyne (C2H2): Each carbon is sp hybridized, forming a σ bond and two π bonds.

Key Equations

• Hybridization: Number of hybrid orbitals = number of atomic orbitals mixed

• Bond Order: $\text{Bond Order} = \frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2}$

Key Terms

• Bond Energy: The energy required to break a bond between two atoms.

• Bond Length: The average distance between the nuclei of two bonded atoms.

• Hybrid Orbital: An orbital formed by the combination of two or more atomic orbitals.

• Sigma (σ) Bond: A covalent bond formed by head-on overlap of orbitals.

• Pi (π) Bond: A covalent bond formed by side-on overlap of p orbitals.

Examples and Applications

• Example 1: In ethene (C2H4), each carbon forms three σ bonds (sp2 hybridized) and one π bond (unhybridized p orbital).

• Example 2: In ethyne (C2H2), each carbon forms two σ bonds (sp hybridized) and two π bonds (unhybridized p orbitals).

Review Table: Summary of Atomic Orbital Hybridization

Type	Number of Hybrid Orbitals	Geometry	Bond Angle
sp	2	Linear	180°
sp2	3	Trigonal planar	120°
sp3	4	Tetrahedral	109.5°

Additional info:

• Hybridization explains molecular geometry as predicted by VSEPR theory.

• Unhybridized p orbitals are responsible for π bonding in multiple bonds.

• Hybridization can be extended to d orbitals for expanded octets (not covered in detail here).