Molecules, Chemical Bonding, and Intermolecular Forces

Comprehensive vocabulary flashcards covering chemical bonding, molecular structures, valence bond theory, molecular orbital theory, and intermolecular forces.

Molecule

A unit of matter that results when two or more atoms are joined by covalent bonds.

Ionic Bond

A bond resulting from the electrostatic attraction between a cation and an anion, formed by a complete transfer of one or more electrons from one atom to another.

Polyatomic Ion

A charged group of more than one kind of atom joined together by covalent bonds.

Electrostatic Potential

The energy a charged particle has due to its position relative to another charged particle; directly proportional to the product of the charges and inversely proportional to the distance between them.

Lattice Energy

The energy released when one mole of an ionic compound forms from its free ions in the gas phase.

Molecular-Level Ionic vs. Covalent Representation

In molecular-level representations, ionic compounds form an extended 3D lattice array of alternating cations and anions (left), whereas covalent compounds exist as distinct, individual molecular units (right).

Percent Yield

The ratio of the actual yield to the theoretical yield expressed as a percentage: $\text{Percent Yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$.

Nonpolar Covalent Bond (Pure Covalent Bond)

A covalent bond characterized by an even distribution of charge, where electrons in the bond are shared equally or almost evenly between two atoms.

Polar Covalent Bond

A bond resulting from unequal sharing of bonding pairs of electrons between two atoms, where the electrons are more attracted to the more electronegative atom.

Bond Length

The distance between the nuclear centers of two atoms joined together in a bond, which corresponds to the internuclear distance at the potential energy minimum.

Bond Energy

The energy needed to break $1\,\text{mole}$ of a particular covalent bond in a molecule or polyatomic ion in the gas phase.

Octet Rule

The guideline stating that atoms of main group elements make bonds by gaining, losing, or sharing electrons to achieve a valence shell containing 8 electrons, or four electron pairs.

Bonding Pair

A pair of electrons shared between two atoms in a covalent bond.

Electronegativity

A relative measure of the ability of an atom in a bond to attract electrons to itself when bonded to another atom.

Resonance

The characteristic of electron distributions when two or more equivalent Lewis structures can be drawn for one compound.

Resonance Structure

One of two or more Lewis structures with the same arrangement of atoms but different arrangements of bonding pairs of electrons.

Resonance Stabilization

The increased stability of a molecular structure resulting from the delocalization of its electrons, which lowers their potential energy.

Formal Charge

A value calculated for an atom in a molecule or polyatomic ion: $\text{Formal Charge} = (\#\text{ valence } e^-) - \frac{1}{2}(\#\text{ bonding } e^-) - (\#\text{ nonbonding } e^-)$.

Bond Order

The number of bonds between atoms: 1 for a single bond, 2 for a double bond, and 3 for a triple bond.

Bond Angle

The angle (in degrees) defined by lines joining the centers of two atoms to a third atom to which they are chemically bonded.

Electronic Geometry

The three-dimensional arrangement of bonding pairs and lone pairs of electrons (electron domains) about a central atom.

Molecular Geometry

The three-dimensional arrangement of the atoms in a molecule.

Bond Dipole

The separation of electrical charge created when atoms with different electronegativities form a covalent bond.

Dipole Moment

A vector representing the polarity of a bond, or a quantitative expression of the net polarity of a molecule.

Valence Bond Theory

A bonding theory stating that covalent bonds are formed by the overlap of atomic orbitals, each containing 1 electron of opposite spin, where overlapping lobes must share the same phase.

Hybridization

In valence bond theory, the mixing of atomic orbitals to generate new sets of orbitals that are available to form covalent bonds with other atoms.

Hybrid Atomic Orbital

In valence bond theory, one of a set of equivalent orbitals about an atom created when specific atomic orbitals are mixed.

Sigma ($\sigma$) Bond

A covalent bond formed by the direct 'head-on' overlap of orbitals along the internuclear axis.

Pi ($\pi$) Bond

A covalent bond formed by the parallel 'side-on' overlap of atomic orbitals.

Atomic Orbital

A wave function whose square gives the probability of finding an electron within a given region of space in an atom.

Molecular Orbital (MO)

A wave function whose square gives the probability of finding an electron within a given region of space in a molecule.

Bonding Orbital

A molecular orbital describing regions of increased electron density between nuclear centers that serves to hold atoms together, being lower in energy than the starting atomic orbitals.

Antibonding Orbital

A molecular orbital (designated with an asterisk, $*$) describing regions of electron density that destabilize a molecule because they do not increase electron density between nuclear centers, being higher in energy than the starting atomic orbitals.

MO Bond Order Formula

The formula used to calculate bond order in molecular orbital theory: $\text{Bond Order} = \frac{\#\text{ bonding } e^- - \#\text{ antibonding } e^-}{2}$.

Intramolecular Forces

Any force that holds together the atoms making up a molecule or compound, including all types of chemical bonds.

Intermolecular Forces (IMFs)

Forces of attraction or repulsion between neighboring particles (multiple molecules or compounds).

Van der Waals Force

Any interaction between neutral atoms and molecules, including hydrogen bonds, dipole-dipole interactions, and London dispersion forces (does not apply to interactions involving ions).

Ion-Dipole Interaction

An attractive electrostatic force between an ion and a molecule that possesses a permanent dipole.

Dipole-Dipole Interaction

An electrostatic attraction between regions of polar molecules that have partial charges of opposite sign.

London Dispersion Force

A weak intermolecular force experienced by all atoms and molecules, regardless of structure, caused by temporary dipoles arising from electron motion.

Hydrogen Bond

The strongest dipole-dipole interaction, occurring between a hydrogen atom bonded to an electronegative atom ($\text{O}$, $\text{N}$, or $\text{F}$) and an electron-rich lone pair on another $\text{O}$, $\text{N}$, or $\text{F}$ atom.

Polarizability

The relative ease with which the electron cloud in a molecule, ion, or atom can be distorted to induce a temporary dipole.

Permanent Dipole

A permanent separation of electrical charge in a molecule resulting from an unequal distribution of bonding and/or lone pairs of electrons.

Temporary Dipole

A transient intermolecular separation of charge caused by the momentary displacement of electrons within molecules.