Unit 2.2 Molecular Structures, VSEPR Theory, and Intermolecular Forces

Vocabulary practice flashcards covering Lewis dot structures, octet rule exceptions, molecular geometries (VSEPR theory), bond and molecular polarity, and intermolecular forces (IMFs) with their effects on boiling points.

Lewis dot structures

Diagrams that show the valence electrons and bonding in a molecule or ion.

Valence electrons

The electrons in the outermost shell of an atom that are the only electrons involved in chemical bonding.

Octet rule

The principle that most atoms react by losing, gaining, or sharing electrons to achieve a stable configuration of eight valence electrons, similar to a noble gas.

Cation

A positive ion formed when a metal atom loses electrons to become stable.

Anion

A negative ion formed when a nonmetal atom gains electrons to become stable.

Covalent bond

A chemical bond formed when atoms share pairs of electrons.

Bonding pair

A shared pair of electrons between two atoms, represented in a Lewis structure by a line.

Lone pair

An unshared pair of electrons on an atom, represented in a Lewis structure by two dots.

Central atom

The atom in the middle of a molecule to which other atoms bond; it is typically the least electronegative atom (except hydrogen) and is often written first in the chemical formula.

Constitutional (Structural) Isomers

Molecules that have the same molecular formula but a different connectivity or bonding sequence of atoms.

Stereoisomers

Molecules that have the same molecular formula and atom-to-atom connectivity but a different arrangement of atoms in three-dimensional space.

VSEPR theory

Valence Shell Electron Pair Repulsion theory; a model based on minimizing the repulsion of shared and unshared electron pairs around the central atom to predict molecular geometry.

Bond angle

The angle formed by any two surrounding atoms and the central atom in a molecule.

Linear molecular shape

A molecular geometry exhibited by molecules such as $CO_2$ (2 bonded atoms, 0 lone pairs on the central atom) and $CO$.

Trigonal planar molecular shape

A molecular geometry formed when there are 3 atoms bonded to the central atom and 0 lone pairs on the central atom, as seen in $BF_3$.

Tetrahedral molecular shape

A molecular geometry formed when there are 4 atoms bonded to the central atom and 0 lone pairs on the central atom, as seen in $CH_4$.

Trigonal pyramidal molecular shape

A molecular geometry formed when there are 3 atoms bonded to the central atom and 1 lone pair on the central atom, as seen in $PI_3$.

Bent molecular shape

A molecular geometry formed when there are 2 atoms bonded to the central atom and 2 lone pairs on the central atom, as seen in $H_2O$.

Nonpolar covalent bond

A bond formed when atoms equally share the bonding electrons.

Polar covalent bond

A bond formed when atoms unequally share electrons due to a difference in electronegativity.

Electronegativity

The ability of an atom in a bond to attract shared electrons; it increases from left to right across periods, decreases top to bottom in groups, and is highest in fluorine.

Polar molecule

A molecule that has a partial positive charge end and a partial negative charge end depending on the covalent bonds it contains and its shape.

Intermolecular forces (IMFs)

Attractive forces that hold molecules together in the liquid phase; their strength directly determines boiling point because thermal energy must overcome them.

London Dispersion Forces (LDF)

Weak, temporary attractive forces arising from random fluctuations in electron density that create temporary dipoles; they are present in all atoms and molecules and are the only IMF in nonpolar substances.

Dipole-dipole interactions

Attractive forces occurring between polar molecules due to the attraction between their permanent dipoles.

Hydrogen bond

A particularly strong type of dipole-dipole interaction occurring when a hydrogen atom covalently bonded to nitrogen, oxygen, or fluorine is attracted to a lone pair on a nearby nitrogen, oxygen, or fluorine atom.

General IMF strength ranking

Ionic bonds $\text{>>}$ Hydrogen bonding $>$ Dipole-dipole interactions $>$ London dispersion forces.