Chapter 7 Notes

Chapter 7: Molecular Geometry, Intermolecular Forces, and Bonding Theories

7.1 Molecular Geometry

Molecular shape prediction is done using the VSEPR (Valence Shell Electron Pair Repulsion) model.
Notation: ABx represents a molecule where A is the central atom, and x is the number of surrounding B atoms. x can vary from 2 to 6.

VSEPR Model

Electrons repel each other, arranging to minimize repulsion.
Electron Domains: Include lone pairs, single bonds, double bonds, and triple bonds.
Geometry types based on electron domains:
2 domains: Linear
3 domains: Trigonal planar
4 domains: Tetrahedral
5 domains: Trigonal bipyramidal
6 domains: Octahedral

Electron-Domain vs. Molecular Geometry

Electron Domain Geometry: Arrangement of electron domains around a central atom
Molecular Geometry: Arrangement of bonded atoms only
A bond angle is defined between two adjacent A–B bonds.
Example geometries:
AB2: linear
AB3: trigonal planar
AB4: tetrahedral

Examples of Molecular Geometry Determination

Molecule SO3: Three bonding domains (trigonal planar).
Molecule ICl4−: Six domains (octahedral); two lone pairs create square planar geometry.

Deviation from Ideal Bond Angles

Lone pairs take more space than bonded pairs, influencing bond angles.
Multiple bonds (like double bonds) exert greater repulsion than single bonds.

Geometry of Molecules with More Than One Central Atom

Complex molecules can be analyzed considering multiple central atoms.

7.2 Molecular Geometry and Polarity

Molecular polarity is influenced by the geometry and bond character:
Diatomic molecules: Difference in electronegativities leads to polarity.
Complex molecules: Requires analysis of individual bond polarities and their shape.
Example: CO2 is nonpolar despite having polar bonds; H2O is polar.

7.3 Intermolecular Forces

Intermolecular forces: Attractiveness between neighboring molecules. Key types include:
Dipole-Dipole interactions: Between polar molecules.
Hydrogen bonding: Special type of dipole-dipole involving H bonded to N, O, or F.
Dispersion forces (London forces): Present in nonpolar molecules, caused by temporary dipoles.
Ion-Dipole interactions: Between ions and polar molecules.

7.4 Valence Bond Theory

Bonds form when atomic orbitals overlap.
Conditions for a bond to form: electrons must be of opposite spin, and overlap must result in lower potential energy.

7.5 Hybridization of Atomic Orbitals

Hybridization: Mixing of atomic orbitals to explain molecular shape.
Types based on number of domains:
sp: 2 domains (linear)
sp²: 3 domains (trigonal planar)
sp³: 4 domains (tetrahedral)
sp³d: 5 domains (trigonal bipyramidal)
sp³d²: 6 domains (octahedral)

7.6 Hybridization in Molecules with Multiple Bonds

In double/triple bonds, hybridized orbitals account for sigma (σ) bonds while unhybridized orbitals form pi (π) bonds.
Example: Ethylene (C2H4) involves sp² hybridization and includes one π bond from unhybridized 2p orbitals.

7.7 Molecular Orbital Theory

Describes bonding beyond valence bond theory; considers molecular orbitals formed from atomic orbitals.
Types of orbitals: Bonding (lower energy, stable) and antibonding (higher energy, less stable).
Bond order: Indicates molecule stability, calculated as (number of bonding electrons - number of antibonding electrons) / 2.

7.8 Bonding Theories and Delocalized Bonding

Lewis Theory: Qualitative, fails to explain bond formation.
Valence Bond Theory: Overlapping atomic orbitals form bonds, but may not predict all molecular properties.
Molecular Orbital Theory: Accurately predicts magnetic properties and delocalization in molecules like benzene.
Resonance structures: Used for species like carbonate ion to represent delocalized electrons.

Summary of Concepts

Geometry and bond types significantly influence molecular properties (e.g., polarity).
Intermolecular forces play crucial roles in states of matter and physical properties.
Understanding bonding theories enhances comprehension of molecular behavior and reactivity.