Chemistry Chp 9 - Molecular Geometry and Bonding Theories

9.1 Molecular Geometry

• Molecular shapes are crucial for chemical processes and can be predicted using Lewis structures and the VSEPR model.

• Focus on ABx molecules, where A is the central atom and B are surrounding atoms.

• Use Lewis structures and VSEPR model to determine shape.

The VSEPR Model

• Electron pairs repel each other; electron domains (lone pairs or bonds) arrange to maximize separation.

• Two electron pairs: 180^\circ (linear).

• Three electron pairs: 120^\circ (trigonal planar).

• Four electron pairs: 109.5^\circ (tetrahedral).

Electron-Domain Geometry and Molecular Geometry

• Electron-domain geometry: arrangement of all electron domains.

• Molecular geometry: arrangement of bonded atoms.

Steps to determine geometries:

1. Draw the Lewis structure.

2. Count electron domains on the central atom.

3. Determine electron-domain geometry using VSEPR.

4. Determine molecular geometry by considering atom positions.

Deviation from Ideal Bond Angles

• Lone pairs take more space than bonding pairs.

• Multiple bonds repel more strongly than single bonds.

9.2 Molecular Geometry and Polarity

• Molecular geometry affects physical/chemical properties.

• Polarity depends on bond polarity and molecular geometry.

• Structural isomers: same formula, different arrangements.

9.3 Valence Bond Theory

• Combines Lewis's shared pairs with quantum mechanics.

• Atoms share electrons via overlapping atomic orbitals with single, unpaired electrons of opposite spins.

• Bond forms if molecule's potential energy is lower than isolated atoms.

• Covalent bond formation is exothermic.

• Bonds form when singly occupied atomic orbitals overlap, resulting in lower potential energy.

9.4 Hybridization of Atomic Orbitals

• Hybridization (mixing) of atomic orbitals extends overlap concept.

• Starts with molecular geometry to explain bonds/angles.

Hybridization of s and p Orbitals

Hybridization of s, p, and d Orbitals

• Steps for hybridized bonding:

1. Draw Lewis structure.

2. Count electron domains (hybrid orbitals).

3. Draw ground-state orbital diagram.

4. Maximize unpaired valence electrons by promotion.

5. Combine atomic orbitals to generate hybrid orbitals.

6. Place electrons in hybrid orbitals singly before pairing.

9.5 Hybridization in Molecules Containing Multiple Bonds

• Rotation exists about single bonds but not double bonds, leading to structural isomers.

9.6 Molecular Orbital Theory

• Explains properties Lewis structures and valence bond theory cannot.

• Diatomic oxygen is paramagnetic (attracted to magnetic fields).

• Diamagnetic species are repelled by magnetic fields.

• Magnetic properties result from electron configuration.

• Paired electrons: diamagnetic; unpaired: paramagnetic.

• Atomic orbitals combine into molecular orbitals with specific shapes/energies.

• Molecular orbitals accommodate max two electrons with opposite spins.

• Number of molecular orbitals equals combined atomic orbitals.

Bonding and Antibonding Molecular Orbitals

Bond Order

• bond \ order = \frac{number \ of \ electrons \ in \ bonding \ MOs - number \ of \ electrons \ in \ antibonding \ MOs}{2}

• Indicates molecule stability.

Molecular Orbital Diagrams

• Lower-energy orbitals fill first, obeying Hund’s rule.

Molecular Orbitals in Heteronuclear Diatomic Species

• Different atoms (e.g., NO).

• More electronegative atom's orbitals are lower in energy.

• Lower-energy atomic orbital contributes more to bonding MO, higher-energy to antibonding.

9.7 Bonding Theories and Descriptions of Molecules with Delocalized Bonding

Strengths and Weaknesses of Bonding Theories

• Lewis Theory:

  • Strength: Qualitative predictions of bond strengths/lengths.

  • Weakness: Two-dimensional, doesn't explain bond formation or differences in compounds like H2, F2, and HF.

• VSEPR Model:

  • Strength: Predicts shapes of molecules/ions.

  • Weakness: Based on Lewis theory, doesn't explain bond formation.

• Valence Bond Theory:

  • Strength: Covalent bond formation via atomic orbital overlap; bonds form due to lower potential energy.

  • Weakness: Fails to explain bonding in molecules like BeCl2, BF3, CH_4.

• Hybridization of Atomic Orbitals:

  • Strength: Extends valence bond theory, explains bonding/geometry of more molecules.

  • Weakness: Fails to predict paramagnetism of O_2.

• Molecular Orbital Theory:

  • Strength: Predicts magnetic and other properties.

  • Weakness: Complex molecular orbital pictures.

Delocalized π Orbitals