CHEM 142 Unit 2: Chemical Bonding, Lewis Structures, Resonance, and VSEPR

Vocabulary practice flashcards covering chemical bonds, Lewis structures, formal charges, resonance, octet rule exceptions, and VSEPR theory based on CHEM 142 Unit 2 lecture notes.

Coulomb's Law

A principle stating that electrostatic force is directly proportional to the magnitude of charges and inversely proportional to the square of the distance between them, expressed as $F = K_e \cdot \frac{q_1 q_2}{r^2}$.

Chemical Bond

An interaction between atoms that lowers the overall energy of the system compared to isolated atoms.

Bond Energy

The quantity of energy required to break a chemical bond, representing the depth of the potential energy well.

Exothermic Reaction

A reaction where product bonds are stronger than reactant bonds, resulting in a net release of energy.

Endothermic Reaction

A reaction where product bonds are weaker than reactant bonds, resulting in a net absorption of energy.

Bond Breaking Energetics

The thermodynamic principle stating that breaking a bond always requires energy and never releases energy; energy released in a reaction originates from forming new, stronger bonds.

Ionic Bond

A bond formed by the transfer of electrons between a metal and a nonmetal, resulting in electrostatic attraction between cations and anions in a 3D lattice.

Nonpolar Covalent Bond

A covalent bond formed between nonmetals with identical or similar electronegativities where electrons are shared equally (e.g., $F-F$, $H-H$).

Polar Covalent Bond

A covalent bond between nonmetals with different electronegativities where electrons are shared unequally, producing partial charges (e.g., $H-F$, $O-H$).

Electronegativity (EN)

An atom's ability to attract shared bonding electrons toward itself, which increases from left to right across a period and from bottom to top up a group.

Bond Dipole

A vector pointing toward the more electronegative atom with a crossed tail ($+$) on the less electronegative atom, representing unequal electron distribution.

Partial Ionic Character

A measure comparing actual electron transfer to complete transfer along the bonding spectrum; a bond with greater than about $50\%$ ionic character is classified as ionic.

Octet Rule

The rule that atoms tend to share electrons until each attains a stable configuration of $8$ valence electrons ($ns^2 np^6$).

Duet Rule

The stability guideline applicable to hydrogen, which requires only $2$ valence electrons because the $n = 1$ shell accommodates a maximum of $2$ electrons.

Formal Charge (FC)

The apparent charge assigned to an atom in a molecule, calculated as: $FC = \text{valence electrons} - \text{number of bonds} - \text{unshared electrons}$.

Structural Isomers

Molecules that share the identical molecular formula but possess different physical connections among their constituent atoms.

Resonance Structures

Two or more equivalent Lewis representations of a single molecule that maintain identical atomic connectivity while altering the distribution of double bonds and lone pairs, denoted with double-headed arrows ($\leftrightarrow$).

Delocalization

The distribution of an electron pair across three or more bonded atoms rather than being confined between two specific nuclei.

Resonance Hybrid

The actual physical molecule of a resonance system, representing a weighted composite blend of all individual resonance structures.

Bond Order Relationship

The inverse relationship between bond length and bond strength: single bonds are the longest and weakest, double bonds are intermediate, and triple bonds are the shortest and strongest.

Sub-octet Exception

An exception to the octet rule where a central atom remains stable with fewer than $8$ valence electrons, such as $H$ (duet, $2$), $Be$ (quartet, $4$), and $B$ (sextet, $6$).

Expanded Octet

An exception to the octet rule occurring in Period $3$ and below (e.g., $P$, $S$, $Cl$, $Br$, $I$) where a central atom accommodates more than $8$ electrons by using empty d orbitals.

Free Radical

A chemical species characterized by an odd number of total valence electrons, resulting in at least one unpaired electron (e.g., $NO_2$).

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a model stating that electron domains surrounding a central atom orient as far apart as possible to minimize mutual electrostatic repulsion.

Electron Domain

A region of electron density around a central atom composed of a single lone pair, a single bond, a double bond, or a triple bond (each counting as exactly one domain).

Electron Domain Geometry (EDG)

The three-dimensional spatial arrangement of all electron domains (both bonding pairs and lone pairs) surrounding the central atom.

Molecular Geometry (MG)

The spatial arrangement of the bonded atoms only, determined by the locations of nuclei with nonbonding lone pairs omitted from the shape's name.

AXE Notation

A systematic notation where $A$ denotes the central atom, $X$ represents the number of bonded atoms, and $E$ denotes the number of lone pairs residing on the central atom.

Equatorial Positions

The three coplanar positions spaced $120^\circ$ apart in a trigonal bipyramidal geometry where lone pairs preferentially locate due to greater spatial volume.

Axial Positions

The two linear positions positioned at $90^\circ$ angles relative to the equatorial plane in a trigonal bipyramidal electron domain geometry.

Square Planar

The molecular geometry ($AX_4E_2$) resulting from an octahedral electron domain geometry containing two lone pairs positioned $180^\circ$ directly opposite each other.

Molecular Polarity

The net dipole moment of a molecule determined by both bond dipole vectors and molecular symmetry; dipoles in symmetric geometries cancel to produce nonpolar molecules, whereas asymmetric dipoles yield polar molecules.