Advanced Theories of Covalent Bonding – Key Vocabulary

Thirty vocabulary flashcards covering key terms from Chapter 8: Advanced Theories of Covalent Bonding, including valence bond theory, hybridization types, σ and π bonding, and foundational concepts of molecular orbital theory.

Valence Bond Theory

Model that explains covalent bonding as the overlap of singly-occupied atomic orbitals with paired spins.

Atomic Orbital Overlap

The sharing of electron density that occurs when two atomic orbitals occupy the same region of space, forming a covalent bond.

Sigma (σ) Bond

A covalent bond formed by end-to-end (head-on) overlap of orbitals along the internuclear axis; all single bonds are σ bonds.

Pi (π) Bond

Bond produced by side-by-side overlap of two unhybridized p orbitals; present in double (one π) and triple (two π) bonds.

Bond Length

The internuclear distance at which the energy of a bonded atom pair is minimized (e.g., 74 pm in H₂).

Hybridization

Mixing of atomic orbitals on an atom to create equivalent hybrid orbitals oriented for optimal bonding geometry.

sp Hybrid Orbitals

Set of two hybrid orbitals formed from one s and one p orbital; oriented 180° apart, giving linear geometry.

sp2 Hybrid Orbitals

Three equivalent hybrids made from one s and two p orbitals; arranged 120° apart, giving trigonal planar geometry.

sp3 Hybrid Orbitals

Four hybrids produced from one s and three p orbitals; directed toward the corners of a tetrahedron (109.5°).

sp3d Hybrid Orbitals

Five hybrids from one s, three p, and one d orbital; yield trigonal bipyramidal electron geometry.

sp3d2 Hybrid Orbitals

Six hybrids formed from one s, three p, and two d orbitals; produce an octahedral arrangement.

Region of Electron Density

Any lone pair, single bond, double bond, or triple bond counted by VSEPR to determine molecular shape and hybridization.

VSEPR Theory

Valence-Shell Electron-Pair Repulsion model that predicts molecular geometry based on minimizing electron-pair repulsions.

Multiple Bond

A double or triple covalent bond consisting of one σ bond plus one or two π bonds respectively.

Resonance

Concept that some molecules are best described by two or more Lewis structures with delocalized π electrons.

Electron Delocalization

Distribution of electrons across several atoms, as in the π system of benzene, lowering overall energy.

Molecular Orbital Theory

Quantum-mechanical description in which atomic orbitals combine to form molecular orbitals extending over the whole molecule.

Bonding Molecular Orbital

Lower-energy MO formed by constructive interference of atomic orbitals; electron occupancy stabilizes the molecule.

Antibonding Molecular Orbital (σ, π)

Higher-energy MO produced by destructive interference; electron occupancy destabilizes the molecule.

Bond Order

Half the difference between bonding and antibonding electrons; indicates bond strength and stability.

Constructive Interference

In-phase combination of wave functions that increases amplitude, creating a bonding orbital.

Destructive Interference

Out-of-phase combination of waves that reduces amplitude, producing an antibonding orbital with nodes.

Homonuclear Diatomic Molecule

Molecule composed of two identical atoms (e.g., O₂, N₂) often analyzed with MO diagrams for bonding.

Hybrid Atomic Orbitals

New orbitals of identical energy formed by mixing standard atomic orbitals on the same atom for bonding.

Side-by-Side Overlap

Lateral interaction of two parallel p orbitals that produces a π bond.

End-to-End Overlap

Head-on orbital interaction along the internuclear axis, forming a σ bond.

Trigonal Planar Geometry

Molecular shape with 120° bond angles resulting from three regions of electron density (sp2 hybridization).

Linear Geometry

Arrangement with 180° bond angle produced by two regions of electron density (sp hybridization).

Tetrahedral Geometry

Shape with four regions of electron density arranged at 109.5° angles (sp3 hybridization).

Octahedral Geometry

Molecular shape with six regions of electron density oriented 90° apart (sp3d2 hybridization).