Chapter 5 Review40 Total Terms
Chapter 5 Key Terms & Vocabulary
Comprehensive index of all scientific terms, definitions, and examples across Chapter 5: Advanced Theories of Bonding.
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From Section 5.1: Valence Bond Theory
Valence bond theoryBonding Theory
A model that describes a covalent bond as the overlap of half-filled atomic orbitals, resulting in a pair of shared electrons between bonded nuclei.
Example: Explains bonding in diatomic molecules like H₂ and HCl.
From Section 5.1: Valence Bond Theory
overlapQuantum Mechanics
The condition where a portion of an atomic orbital from one atom and a portion of an orbital from another atom occupy the same region of space.
Example: Greater orbital overlap produces stronger, more stable covalent bonds.
From Section 5.1: Valence Bond Theory
sigma bonds (σ bonds)Bond Types
Covalent bonds formed when electron density is concentrated along the internuclear axis via end-to-end orbital overlap.
Example: All single covalent bonds are sigma bonds (e.g., s-s, s-p, or end-to-end p-p overlap).
From Section 5.1: Valence Bond Theory
pi bond (π bond)Bond Types
A covalent bond formed by the side-by-side overlap of parallel p orbitals, with regions of electron density residing on opposite sides of the internuclear axis.
Example: Present in double (1 , 1 ) and triple (1 , 2 ) bonds.
From Section 5.1: Valence Bond Theory
nodeQuantum Mechanics
A plane or spatial region along the internuclear axis of a pi bond where the probability of finding an electron is exactly zero.
Example: Located in the nodal plane directly between bonded atoms in a pi bond.
From Section 5.2: Hybrid Atomic Orbitals
hybridizationBonding Theory
The mathematical combination of atomic wave functions to form new, specialized hybrid orbitals with tailored shapes and orientations during covalent bonding.
From Section 5.2: Hybrid Atomic Orbitals
hybrid orbitalsBonding Theory
New, equivalent orbitals formed by combining atomic orbitals on a bonded atom; their shapes and energies match the molecule's spatial geometry.
From Section 5.2: Hybrid Atomic Orbitals
sp hybrid orbitalsHybridization Types
A set of two equivalent hybrid orbitals oriented linearly at 180° resulting from the mixing of one s and one p orbital.
Example: Found in linear molecules like BeCl2 and CO2.
From Section 5.2: Hybrid Atomic Orbitals
sp2 hybrid orbitalsHybridization Types
A set of three equivalent hybrid orbitals arranged in a trigonal planar geometry (120°) resulting from the mixing of one s and two p orbitals.
Example: Found in trigonal planar molecules like BH3 and formaldehyde.
From Section 5.2: Hybrid Atomic Orbitals
sp3 hybrid orbitalsHybridization Types
A set of four equivalent hybrid orbitals arranged in a tetrahedral geometry (109.5°) resulting from the mixing of one s and three p orbitals.
Example: Found in tetrahedral molecules like CH4, NH3, and H2O.
From Section 5.2: Hybrid Atomic Orbitals
sp3d hybrid orbitalsHybridization Types
A set of five equivalent hybrid orbitals arranged in a trigonal bipyramid resulting from mixing one s, three p, and one d orbital.
Example: Found in expanded octet molecules like PCl5.
From Section 5.2: Hybrid Atomic Orbitals
sp3d2 hybrid orbitalsHybridization Types
A set of six equivalent hybrid orbitals arranged in an octahedron resulting from mixing one s, three p, and two d orbitals.
Example: Found in octahedral molecules like SF6.
From Section 5.3: Multiple Bonds
Sigma Bond (σ Bond)Bonding Theory
A covalent bond formed by the end-to-end overlap of atomic orbitals, featuring electron density concentrated symmetrically along the internuclear axis.
Example: C–H single bonds in ethene
From Section 5.3: Multiple Bonds
Pi Bond (π Bond)Bonding Theory
A covalent bond formed by the side-by-side overlap of unhybridized parallel p orbitals, with electron density located in lobes above and below the internuclear axis.
Example: The second bond in a carbon-carbon double bond
From Section 5.3: Multiple Bonds
sp² HybridizationHybridization
The mixing of one s and two p atomic orbitals to create three equivalent hybrid orbitals directed in a trigonal planar arrangement.
Example: Carbon atoms in ethene (C₂H₄)
From Section 5.3: Multiple Bonds
sp HybridizationHybridization
The mixing of one s and one p atomic orbitals to create two equivalent hybrid orbitals directed in a linear arrangement.
Example: Carbon atoms in acetylene (C₂H₂)
From Section 5.3: Multiple Bonds
Unhybridized p OrbitalAtomic Orbitals
An atomic p orbital that does not participate in hybridization, remaining perpendicular to the hybrid orbital plane to participate in π bonding.
Example: The remaining 2p orbital on carbon in ethene
From Section 5.3: Multiple Bonds
ResonanceMolecular Structure
A condition where two or more valid Lewis structures with identical atomic frameworks but different π-electron placements describe a single molecule.
Example: Alternate double bond arrangements in benzene
From Section 5.3: Multiple Bonds
DelocalizationBonding Theory
The sharing of π electrons across three or more adjacent atoms rather than being localized between a specific pair of bonded atoms.
Example: π electrons distributed across the entire benzene ring
From Section 5.3: Multiple Bonds
Ethene (C₂H₄)Molecules
A simple hydrocarbon featuring a carbon-carbon double bond and sp² hybridized carbon atoms.
Example: C₂H₄
From Section 5.3: Multiple Bonds
Acetylene (C₂H₂)Molecules
A linear hydrocarbon featuring a carbon-carbon triple bond and sp hybridized carbon atoms.
Example: HC≡CH
From Section 5.3: Multiple Bonds
Benzene (C₆H₆)Molecules
A cyclic hydrocarbon with a planar hexagonal ring structure where π electrons are fully delocalized through resonance.
Example: C₆H₆
From Section 5.4: Molecular Orbital Theory
ParamagnetismMagnetic Properties
Property of being attracted to an external magnetic field due to the presence of unpaired electrons.
From Section 5.4: Molecular Orbital Theory
GouyExperimental Methods
A balance system used to measure magnetic susceptibility and experimentally determine the number of unpaired electrons in a sample.
From Section 5.4: Molecular Orbital Theory
DiamagneticMagnetic Properties
Property of materials in which all electrons are paired, causing them to weakly repel a magnetic field.
From Section 5.4: Molecular Orbital Theory
Molecular Orbital TheoryBonding Theories
A model of chemical bonding describing electrons as delocalized across an entire molecule via quantum-mechanical wave functions.
From Section 5.4: Molecular Orbital Theory
Molecular Orbital (Ψ²)Quantum Mechanics
A mathematical region of space in a molecule where valence electrons are most likely to be found.
From Section 5.4: Molecular Orbital Theory
Homonuclear Diatomic MoleculesMolecular Structure
Covalent molecules composed of two identical atoms bonded together (e.g., H₂, O₂).
From Section 5.4: Molecular Orbital Theory
Linear Combination of Atomic Orbitals (LCAO)Mathematical Methods
The mathematical addition and subtraction of atomic orbital wave functions to generate molecular orbitals.
From Section 5.4: Molecular Orbital Theory
σs Molecular OrbitalMolecular Orbitals
A lower-energy bonding molecular orbital formed by the in-phase overlap of two s atomic orbitals.
From Section 5.4: Molecular Orbital Theory
σs* Molecular OrbitalMolecular Orbitals
A higher-energy antibonding molecular orbital formed by the out-of-phase overlap of two s atomic orbitals.
From Section 5.4: Molecular Orbital Theory
Bonding OrbitalsOrbital Interactions
Molecular orbitals that lower system energy and draw nuclei together when occupied by electrons.
From Section 5.4: Molecular Orbital Theory
Antibonding OrbitalsOrbital Interactions
Molecular orbitals that increase system energy, feature a nodal plane between nuclei, and pull nuclei apart.
From Section 5.4: Molecular Orbital Theory
Pi (π) Bonding Molecular OrbitalMolecular Orbitals
A bonding molecular orbital formed by the side-by-side overlap of parallel p atomic orbitals.
From Section 5.4: Molecular Orbital Theory
π* Antibonding Molecular OrbitalMolecular Orbitals
An antibonding molecular orbital resulting from out-of-phase side-by-side overlap of p atomic orbitals.
From Section 5.4: Molecular Orbital Theory
Degenerate OrbitalsQuantum Mechanics
Molecular or atomic orbitals that possess exactly the same energy level.
From Section 5.4: Molecular Orbital Theory
KohnScientists
Theoretical physicist Walter Kohn, co-developer of density functional theory for computing molecular orbital structures.
From Section 5.4: Molecular Orbital Theory
Molecular Orbital DiagramVisualization
A visual chart displaying the relative energy levels of atomic and molecular orbitals and their electron occupancies.
From Section 5.4: Molecular Orbital Theory
Bond OrderBond Properties
A measure of bond strength calculated as half the difference between bonding and antibonding electrons.
From Section 5.4: Molecular Orbital Theory
s-p MixingOrbital Interactions
Mathematical interaction between s and p orbitals of similar energy that alters relative molecular orbital energy ordering.
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5.4 Molecular Orbital Theory
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