Lecture 11: Chemical Bonding and Bonding Theories

These flashcards cover key concepts related to chemical bonding theories, hybridization, molecular orbitals, and bond properties, aiding in the understanding of covalent bonding.

Valence Bond Theory

Describes covalent bonds as the overlap of atomic orbitals.

Hybrid Orbitals

Mixtures of atomic orbitals with intermediate energy, corresponding to the number of hybrid orbitals formed.

sp Hybrid Orbitals

Formed from one s and one p orbital, leading to linear geometry.

sp² Hybrid Orbitals

Formed from one s and two p orbitals, leading to trigonal planar geometry.

sp³ Hybrid Orbitals

Formed from one s and three p orbitals, leading to tetrahedral geometry.

Sigma (σ) Bond

A bond resulting from the end-to-end overlap of orbitals, with maximum electron density along the bond.

Pi (π) Bond

A bond resulting from side-to-side overlap of p orbitals, with zero electron density along the bond.

Bond Order

A measure of the stability of a bond calculated as 1/2( # bonding electrons - # antibonding electrons).

Magnetic Properties

Paramagnetic molecules contain unpaired electrons and are attracted to magnetic fields, while diamagnetic molecules have all paired electrons and are weakly repelled.

Molecular Orbitals (MOs)

Formed from the combination of atomic orbitals, describing the distribution of electrons in a molecule.

Bonding Molecular Orbital

An orbital where electron density is greatest between bonded atoms.

Antibonding Molecular Orbital

An orbital with zero electron density at a point between bonded atoms.

Diatomic Molecules

Molecules consisting of two atoms, where only atomic orbitals of similar energies contribute to molecular orbitals.

Bonding in C2H4

Involves sp² hybridization and includes both σ and π bonds.

Bonding in HCN

A linear molecule with sp hybridization, containing both σ and π bonds.

MO Diagram for H₂

Shows the bonding and antibonding molecular orbitals, with a bond order indicating stability.

MO Diagram for He2

Indicates that He2 does not exist due to equal numbers of bonding and antibonding electrons.