Chemical Bonding, Molecular Geometry, and Valence Bond Theory

Vocabulary practice flashcards covering chemical bonding, VSEPR molecular geometry, resonance, hybridization, valence bond theory, and course exam preparation procedures.

Chemical Bond

An interaction that forms when the potential energy of the bonded atoms is lower than the potential energy of the separate, isolated atoms.

The Octet Rule (Energy Valley Model)

The principle that atoms (such as C, N, O, and F) reach the bottom of their potential energy valley with $8$ valence electrons ($s^2p^6$), achieving maximum nuclear attraction before any additional electron is forced across a large energy gap into a higher, shielded shell.

Ionic Bond

A chemical bond characterized by the complete transfer of valence electrons from a metal to a nonmetal, forming oppositely charged ions held together entirely by electrostatic attraction without electron sharing.

Covalent Bond

A chemical bond formed by the sharing of valence electron pairs between nonmetal or metalloid atoms, which lowers potential energy by attracting the nuclei of both atoms.

Electronegativity

A relative scale describing the ability of a bonded atom to attract shared electrons toward itself; it generally increases across a period and decreases down a group.

Polar Covalent Bond

A covalent bond formed between atoms with an electronegativity difference ($\Delta\chi$) between $0.5$ and $2.0$, leading to unequal electron sharing, partial charges ($\delta^+$ and $\delta^-$), and a bond dipole moment.

Nonpolar Covalent Bond

A bond between atoms with an electronegativity difference ($\Delta\chi$) less than $0.5$, characterized by an essentially equal sharing of electron density.

Formal Charge

A bookkeeping charge assigned to an atom in a Lewis structure, calculated as: $\text{Formal Charge} = \text{Total valence electrons} - (\text{lone pair electrons} + \text{one electron per bond})$.

Resonance Structures

Two or more valid Lewis structures for the same molecule that possess identical nuclear arrangements but differ in the spatial distribution of valence electrons.

Resonance Hybrid

The true, single chemical species that exists in nature as an intermediate average of all valid resonance contributors, rather than oscillating between individual structures.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a model stating that electron domains (bonding pairs and lone pairs) around a central atom repel each other and position themselves as far apart as possible to minimize potential energy.

Lone Pair Repulsion Effect

The steric and electrostatic phenomenon where nonbonding lone pairs occupy more physical space than bonding pairs, exerting greater repulsive forces and compressing adjacent bond angles.

Equatorial Preference in Trigonal Bipyramidal Systems

The geometric requirement in five-domain systems (such as $\text{SF}_4$) that lone pairs reside in equatorial positions to encounter only two $90^{\circ}$ lone pair-bonding pair repulsions instead of three $90^{\circ}$ repulsions at axial sites.

Dipole Moment

A vector quantity indicating the magnitude and direction of net molecular charge separation, depicted by an arrow pointing from the more positive region to the more negative region.

Molecular Polarity

The overall polarity of a molecule determined by the vector addition of all individual bond dipole moments based on the three-dimensional molecular geometry.

Valence Bond (VB) Theory

A bonding model positing that covalent bonds form when two half-filled atomic or hybrid orbitals overlap in space and their electrons pair up with opposite spins between both nuclei.

Hybrid Orbitals

Mathematical combinations of standard atomic orbitals ($s$, $p$, and $d$) localized on an atom that create directional sets of orbitals matching observed molecular geometries.

$sp^3$ Hybridization

The hybridization resulting from mixing one $2s$ orbital and three $2p$ orbitals to yield four equivalent orbitals oriented at $109.5^{\circ}$ (tetrahedral geometry), each having $\frac{1}{4}$ s character and $\frac{3}{4}$ p character.

$sp^2$ Hybridization

The hybridization resulting from mixing one $2s$ orbital and two $2p$ orbitals to yield three equivalent orbitals oriented at $120^{\circ}$ (trigonal planar geometry), each with $\frac{1}{3}$ s character and $\frac{2}{3}$ p character, leaving one unhybridized $p$ orbital.

$sp$ Hybridization

The hybridization resulting from mixing one $2s$ orbital and one $2p$ orbital to produce two equivalent orbitals oriented at $180^{\circ}$ (linear geometry), each with $\frac{1}{2}$ s character and $\frac{1}{2}$ p character, leaving two unhybridized $p$ orbitals.

$sp^3d$ Hybridization

The mixing of one $s$, three $p$, and one $d$ orbital to form five hybrid orbitals arranged in a trigonal bipyramidal geometry for expanded-octet atoms.

$sp^3d^2$ Hybridization

The mixing of one $s$, three $p$, and two $d$ orbitals to create six hybrid orbitals arranged in an octahedral geometry for expanded-octet atoms.

$\sigma$ (Sigma) Bond

A covalent bond formed by end-to-end overlap between atomic or hybrid orbitals, concentrating electron density directly along the internuclear axis.

$\pi$ (Pi) Bond

A covalent bond formed by side-to-side overlap of parallel unhybridized $p$ atomic orbitals, concentrating electron density above and below the internuclear plane.

Allene ($\text{H}_2\text{C}=\text{C}=\text{CH}_2$) Geometry

A non-planar molecule whose central carbon is $sp$ hybridized with two mutually perpendicular unhybridized $2p$ orbitals, forcing the terminal $sp^2$ hybridized $\text{CH}_2$ groups to lie in perpendicular planes.

Five-Minute Mistake Log

A reflection practice completed after each LBLA, BIT, and Expo where students record missed or guessed questions, identify what caused the error, and write the correct reasoning to build a personal exam study guide.

Collaborative Prep Exam Protocol

A structured two-stage exam review session consisting of a $30\text{-minute}$ individual attempt in total silence (Stage 1), a $3\text{-minute}$ pivot roster check, and a $40\text{-minute}$ collaborative consensus discussion (Stage 2) earning course credit for participation and group correctness.