1/29
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.
Name | Mastery | Learn | Test | Matching | Spaced |
---|
No study sessions yet.
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).