Basic Concepts of Chemical Bonding, Molecular Geometry, and Bonding Theories

Vocabulary practice flashcards covering basic chemical bonding concepts, Lewis structures, lattice energies, electronegativity, VSEPR theory, hybridization, and molecular orbital theory.

Valence Electrons

The electrons in the highest occupied primary energy level ($s$ and $p$ electrons in a condensed electron configuration), which can reach a maximum of $8$.

Octet Rule

The principle that atoms tend to gain, lose, or share electrons until they achieve a stable arrangement of $8$ valence electrons, becoming isoelectronic with a noble gas.

Metallic Bond

A type of chemical bond where metal atoms are held together by a mobile "sea of electrons" surrounding the nuclei.

Ionic Bond

A chemical bond resulting from local electrostatic attractions between oppositely charged ions, typically formed between a metallic cation and a nonmetallic anion.

Covalent Bond

A chemical bond in which atoms share pairs of electrons in localized bonds to achieve a stable octet.

Lattice Energy

The amount of energy required to completely separate one mole of a solid ionic compound into its gaseous ions; this highly endothermic separation corresponds to a large energy release when the lattice forms.

Born-Haber Cycle

A thermodynamic cycle used to analyze the formation of ionic solids by relating lattice energy to enthalpy of formation, ionization energy, electron affinity, and atomization enthalpies.

Electronegativity

The unitless ability of an atom in a molecule to attract bonding electrons to itself, characterized by Linus Pauling on a thermochemical scale.

Dipole Moment ($\mu$)

A quantitative measure of polarity defined by $\mu = Q r$, where $Q$ is the magnitude of separated charge and $r$ is the distance between charges.

Nonpolar Covalent Bond

A covalent bond formed between two atoms having a small electronegativity difference of less than $0.5$.

Polar Covalent Bond

A covalent bond in which electrons are shared unequally due to a medium electronegativity difference between $0.5$ and $2.0$.

Formal Charge

The charge assigned to an individual atom in a Lewis structure, computed as valence electrons minus nonbonding electrons minus half of bonding electrons.

Resonance Structures

Two or more valid Lewis structures for a single molecule or ion having identical arrangements of atoms but differing distributions of electrons.

Expanded Octet

A hypervalent condition occurring in elements from Period $3$ and below that accommodate more than $8$ valence electrons around a central atom.

Bond Enthalpy

The positive energy required to break one mole of a specific chemical bond in the gas phase; overall reaction enthalpy is calculated as $\Delta H = \sum(\text{bonds broken}) - \sum(\text{bonds formed})$.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a model predicting geometric shapes based on the principle that electron domains repel one another and stay as far apart as possible.

Electron Domain

A region of electron density around a central atom consisting of a single bond, multiple bond, or nonbonding lone pair.

Electron-Domain Geometry

The three-dimensional arrangement determined by considering all bonding and nonbonding electron domains surrounding a central atom.

Molecular Geometry

The actual spatial conformation and arrangement of only the bonded atoms in a molecule, influenced by the presence of nonbonding lone pairs.

Linear Geometry

The geometry characteristic of two electron domains around a central atom ($AB_2$) with a predicted bond angle of $180^\circ$.

Trigonal Planar Geometry

The geometry produced by three electron domains in one plane around a central atom ($AB_3$) with predicted bond angles of $120^\circ$.

Tetrahedral Geometry

The geometry produced by four electron domains around a central atom ($AB_4$) with predicted bond angles of $109.5^\circ$.

Trigonal Bipyramidal Geometry

The geometry formed by five electron domains ($AB_5$) consisting of two axial positions and three equatorial positions with predicted bond angles of $90^\circ$ and $120^\circ$.

Octahedral Geometry

The geometry formed by six electron domains ($AB_6$) symmetrically oriented around a central atom with predicted bond angles of $90^\circ$.

Valence-Bond Theory

A model stating that covalent bonds form when valence atomic orbitals overlap, concentrating electron density between two nuclei to balance attractive and repulsive forces.

Hybrid Orbitals

Identical atomic orbitals created by the mathematical mixing (hybridization) of nonequivalent atomic orbitals on a central atom during covalent bonding.

$sp$ Hybridization

The combination of one $s$ orbital and one $p$ orbital to produce two equivalent hybrid orbitals oriented at $180^\circ$, giving a linear geometry.

$sp^2$ Hybridization

The combination of one $s$ orbital and two $p$ orbitals forming three equivalent hybrid orbitals oriented at $120^\circ$, yielding a trigonal planar geometry.

$sp^3$ Hybridization

The combination of one $s$ orbital and three $p$ orbitals producing four equivalent hybrid orbitals directed toward tetrahedral vertices at angles of $109.5^\circ$.

Sigma ($\sigma$) Bond

A covalent bond formed by end-to-end overlap of atomic or hybrid orbitals with maximum electron density along the internuclear axis.

Pi ($\pi$) Bond

A covalent bond formed by side-by-side overlap of parallel unhybridized $p$ orbitals with electron density concentrated above and below the internuclear axis.

Delocalized Pi Bond

A pi bonding system in which electrons extend over three or more adjacent atoms simultaneously rather than being confined between two specific nuclei.

Molecular Orbital (MO) Theory

A bonding model that describes electrons in molecules through wave functions called molecular orbitals, which extend across the entire molecule rather than individual atoms.

Bonding Molecular Orbital

A molecular orbital with high electron density between the nuclei that is lower in energy than the original parent atomic orbitals, stabilizing the molecule.

Antibonding Molecular Orbital

A high-energy molecular orbital possessing a nodal plane between nuclei that destabilizes the molecule when occupied by electrons.

Bond Order

A value indicating net bond stability calculated as $\frac{\text{No. of } e^- \text{ in bonding orbitals} - \text{No. of } e^- \text{ in antibonding orbitals}}{2}$, where higher values denote shorter, stronger bonds.

Paramagnetism

The magnetic property of a substance containing one or more unpaired electrons, causing it to be attracted toward an external magnetic field.

Diamagnetism

The magnetic property of a substance in which all electrons are paired, resulting in weak repulsion by an external magnetic field.