Chapter 6: Chemical Bonding II – Valence Bond Theory and Molecular Orbital Theory

Valence Bond Theory

Introduction to Valence Bond Theory

Valence Bond (VB) Theory explains how atomic orbitals combine to form chemical bonds in molecules. It focuses on the overlap of atomic orbitals from different atoms to create a shared electron pair, resulting in a covalent bond.

• Atomic Orbitals: Regions in an atom where electrons are likely to be found (e.g., 1s, 2s, 2p).

• Bond Formation: Occurs when orbitals from two atoms overlap and share electrons.

• Example: In methane (CH4), carbon's 2s and 2p orbitals combine with hydrogen's 1s orbitals.

Theoretical Prediction vs. Observed Reality

• Theoretical Prediction: If carbon used its unhybridized 2s and 2p orbitals, the H–C–H bond angles would be 90°.

• Observed Reality: Experimental data shows that the H–C–H bond angles in methane are 109.5°, indicating a tetrahedral geometry.

Hybridization

To explain the observed geometry, VB theory introduces the concept of hybridization, where atomic orbitals mix to form new, equivalent hybrid orbitals.

• sp3 Hybridization: One 2s and three 2p orbitals combine to form four sp3 hybrid orbitals, each oriented toward the corners of a tetrahedron (109.5° angles).

• Hybrid Orbitals: Have larger lobes than standard p orbitals, allowing for greater overlap and stronger bonds.

Formation of Hybrid Orbitals

• sp3 Hybridization: Four equivalent sp3 orbitals (as in CH4 and NH3).

• sp2 Hybridization: Three sp2 hybrid orbitals and one unhybridized p orbital (as in C2H4).

• sp Hybridization: Two sp hybrid orbitals and two unhybridized p orbitals (as in C2H2).

Types of Bonds

• Sigma (σ) Bonds: Formed by the head-on overlap of orbitals (e.g., sp3-1s in CH4).

• Pi (π) Bonds: Formed by the side-on overlap of unhybridized p orbitals (present in double and triple bonds).

Examples of Hybridization

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

Orbital Diagrams and Electron Configurations

• Before Hybridization (Carbon): 1s2 2s2 2p2

• After Hybridization (sp3): Four sp3 orbitals, each with one electron, ready to form four bonds.

Molecular Orbital Theory

Introduction to Molecular Orbital (MO) Theory

MO Theory describes the formation of molecular orbitals by the linear combination of atomic orbitals from all atoms in a molecule. Electrons are delocalized over the entire molecule, not just between two atoms.

• Bonding Orbitals: Lower in energy, result from constructive interference of atomic orbitals.

• Antibonding Orbitals: Higher in energy, result from destructive interference.

MO Diagrams and Bond Order

• Bond Order Formula:

$\text{Bond Order} = \frac{1}{2}[(\text{number of electrons in bonding orbitals}) - (\text{number of electrons in antibonding orbitals})]$

• Example: O2 molecule has a bond order of 2, indicating a double bond.

Application to Diatomic Molecules

• MO theory explains magnetic properties (paramagnetism/diamagnetism) and bond orders in molecules like O2, N2, and F2.

Practice and Application

Sample Problems and Applications

• Draw orbital diagrams for atoms and molecules before and after hybridization.

• Assign hybridization to atoms in molecules (e.g., CH4, NH3, CO2).

• Predict molecular geometry using VSEPR theory and hybridization.

• Use MO diagrams to predict bond order and magnetic properties.

Bond Energies and Trends

Bond	Bond Energy (kJ/mol)
Br–Cl	218
Br–Br	193
I–Br	175

Trend: Bond energy decreases as the size of the atoms increases, due to less effective orbital overlap.

Key Definitions

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.

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

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

• Bond Order: Number of chemical bonds between a pair of atoms.

• Paramagnetic: Molecule with unpaired electrons, attracted to a magnetic field.

• Diamagnetic: Molecule with all electrons paired, not attracted to a magnetic field.

Summary Table: Hybridization and Geometry

Hybridization	Number of Electron Domains	Geometry	Example
sp3	4	Tetrahedral	CH4
sp2	3	Trigonal planar	BF3
sp	2	Linear	BeCl2

Additional info: These notes integrate both visual and textual content from the provided materials, expanding on the concepts of hybridization, molecular geometry, and molecular orbital theory as covered in a standard General Chemistry curriculum.