Chemical Bonding, Molecular Geometry, and Bonding Theories

Comprehensive vocabulary flashcards covering chemical bonding types, Lewis structures, lattice energies, electronegativity, VSEPR theory, hybridization, orbital overlap, and molecular orbital theory.

Chemical Bond

The attractive force holding two or more atoms or ions together, categorized into covalent, ionic, and metallic bonds.

Covalent Bond

A chemical bond resulting from the sharing of valence electron pairs between atoms, typically occurring between nonmetals.

Ionic Bond

A chemical bond resulting from the electrostatic attraction between oppositely charged ions formed by the transfer of electrons from a metal to a nonmetal.

Metallic Bond

The attractive force that binds atoms of pure metals together in a lattice.

Valence Electrons

The electrons residing in an atom's outermost shell that participate in chemical bonding and determine the number and strength of bonds that can form.

Lewis Symbol

A pictorial representation developed by Gilbert N. Lewis where an element's chemical symbol is surrounded by dots that represent its valence electrons.

Octet Rule

The principle that atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons, matching the stable $s^2p^6$ configuration of noble gases.

Lattice Energy

The energy required to completely separate one mole of a solid ionic compound into its gaseous ions; it increases with higher ionic charges and decreasing ion separation distance according to Coulomb's law ($E_{el} = \frac{k Q_1 Q_2}{d}$).

Electronegativity

The ability of an atom in a molecule to attract shared electrons to itself, scaled by Linus Pauling from $0.7$ for $\text{Cs}$ to $4.0$ for $\text{F}$.

Bond Polarity

A measure of how equally or unequally electrons in a covalent bond are shared, determined by the electronegativity difference between the bonded atoms.

Formal Charge

The charge assigned to an individual atom in a Lewis structure, computed as: $\text{Valence electrons} - (\text{Number of bonds} + \text{Number of lone pair electrons})$.

Resonance Structures

Two or more equally valid Lewis structures used in combination to accurately represent a molecule or ion with delocalized electrons that cannot be described by a single structure.

Expanded Octet

An exception to the octet rule where a central atom from the third period or below accommodates more than eight valence electrons by utilizing low-energy $d$-orbitals.

Bond Dissociation Enthalpy ($D$)

The energy required to break a specific covalent bond in the gas phase; it is always a positive quantity because bond breaking is endothermic.

Reaction Enthalpy ($\Delta H_{rxn}$)

An estimate of net reaction heat calculated from bond dissociation energies via the formula: $\Delta H_{rxn} = \sum D(\text{bonds broken}) - \sum D(\text{bonds formed})$.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a model stating that valence electron pairs (domains) around a central atom repel each other and adopt a 3D geometry that minimizes this repulsion.

Electron Domain

A region of electron density around a central atom, where each nonbonding electron pair, single bond, double bond, or triple bond counts as a single domain.

Electron-Domain Geometry

The three-dimensional arrangement in space of all electron domains (both bonding and nonbonding pairs) surrounding a central atom.

Molecular Geometry

The three-dimensional arrangement defined exclusively by the positions of bonded atoms in a molecule, excluding nonbonding electron pairs.

Axial and Equatorial Positions

The two distinct positions in a trigonal bipyramidal geometry: equatorial positions form an equilateral triangle in a single plane ($120^\circ$ apart), while axial positions lie directly above and below this plane ($90^\circ$ to the equatorial plane).

Dipole Moment ($\mu$)

A quantitative measure of molecular polarity defined by $\mu = Qr$, where $Q$ is the magnitude of separated charge, $r$ is the separation distance, and units are reported in debyes (D).

Valence Bond Theory

A model of chemical bonding stating that covalent bonds form when valence atomic or hybrid orbitals overlap and share two electrons of opposite spin.

Hybrid Orbitals

Equivalent degenerate orbitals formed by mathematical mixing of standard atomic orbitals ($s, p, d$) on an atom to adopt shapes and spatial orientations suitable for bonding.

$sp$ Hybrid Orbitals

Two degenerate hybrid orbitals formed by mixing one $s$ and one $p$ atomic orbital, oriented $180^\circ$ apart in a linear geometry and leaving two unhybridized $p$ orbitals.

$sp^2$ Hybrid Orbitals

Three degenerate hybrid orbitals formed by mixing one $s$ and two $p$ atomic orbitals, oriented in a trigonal plane with $120^\circ$ angles and leaving one unhybridized $p$ orbital.

$sp^3$ Hybrid Orbitals

Four degenerate hybrid orbitals formed by mixing one $s$ and three $p$ atomic orbitals, pointing toward the vertices of a tetrahedron with angles of $109.5^\circ$.

Sigma ($\sigma$) Bond

A covalent bond characterized by direct, head-to-head overlap of orbitals along the internuclear axis with cylindrical symmetry of electron density.

Pi ($\pi$) Bond

A covalent bond formed by side-to-side overlap of parallel unhybridized $p$ orbitals, resulting in regions of electron density located above and below the internuclear axis.

Delocalized $\pi$ Bonding

A bonding arrangement where $\pi$ electrons are smeared over three or more adjacent atoms through continuous $p$-orbital overlap, stabilizing the species.

Molecular Orbital (MO) Theory

A quantum mechanical approach to bonding that treats electrons as delocalized across an entire molecule by combining atomic orbitals into molecular wavefunctions.

Bonding Molecular Orbital

A molecular orbital of lower energy than the original atomic orbitals, formed by constructive wave overlap that concentrates electron density directly between the nuclei.

Antibonding Molecular Orbital ($\sigma^*$ or $\pi^*$)

A molecular orbital of higher energy than the original atomic orbitals, formed by destructive wave overlap and characterized by a nodal plane with zero electron density between the nuclei.

Bond Order

A measure of bond strength and stability defined by the formula: $\text{Bond order} = \frac{1}{2}(\text{bonding electrons} - \text{antibonding electrons})$.

$2s-2p$ Orbital Interaction

The mixing of $2s$ and $2p$ orbitals in homonuclear diatomics ($B_2$, $C_2$, $N_2$) that shifts the $\sigma_{2p}$ molecular orbital higher in energy than the $\pi_{2p}$ orbitals.

Paramagnetism

A magnetic behavior exhibited by molecules possessing one or more unpaired electrons, causing them to be drawn into an external magnetic field.

Diamagnetism

A magnetic behavior exhibited by molecules with all paired electrons, resulting in weak repulsion by an external magnetic field.