Unit 7 – Ionic & Covalent Bonding (Vocabulary Review)

18 vocabulary flashcards covering ionic bonding, lattice energy, ionic radii, covalent bonding, molecular-orbital concepts, σ and π bonds, and related terminology from the Unit 7 lecture.

Ionic bond

A chemical bond formed by the transfer of electrons from a metal to a non-metal, creating oppositely charged ions that are held together by strong electrostatic attraction in a crystal lattice.

Lattice energy

The stabilizing energy released when gaseous ions assemble into an ionic solid; responsible for the high stability of ionically bonded compounds like NaCl.

Sodium chloride formation

Process in which Na atoms donate electrons to Cl atoms, producing Na⁺ and Cl⁻ ions that arrange in a cubic lattice to form common table salt.

Ionic radius

The effective size of an ion in a crystal, determined experimentally from electron-density maps where density approaches zero along the line between neighboring ions.

Electron-density map

A contour plot showing regions of equal electron density (e⁻ Å⁻³); used to locate ionic radii and bonding regions in crystals such as NaCl.

Covalent bond

A chemical bond in which two atoms share one or more pairs of electrons, also called an electron-pair bond.

H–H bond energy

−436 kJ mol⁻¹; the energy change associated with forming the covalent bond in H₂ at its equilibrium distance.

Equilibrium bond length (H₂)

0.74 Å; the internuclear distance at which the potential energy of the H₂ molecule is at a minimum.

Molecular orbital (MO)

An orbital that extends over an entire molecule, formed by the linear combination of atomic orbitals (LCAO).

Bonding orbital

An MO resulting from constructive overlap of atomic orbitals; electron occupancy lowers the system’s energy and promotes bond formation.

Antibonding orbital

An MO (denoted with an asterisk, e.g., σ* or π*) formed by destructive overlap; electron occupancy raises the energy and weakens bonding.

Bond order

Half the difference between the number of bonding and antibonding electrons; quantifies bond strength (e.g., H₂ has bond order 1).

σ (sigma) bond

A covalent bond with cylindrical symmetry around the bond axis, typically formed by head-on overlap of s or p orbitals.

π (pi) bond

A covalent bond formed by side-by-side overlap of p orbitals above and below the bond axis; possesses nodal planes along the axis.

Constructive overlap

In MO theory, the in-phase (same sign) combination of orbital lobes that increases electron density between nuclei, generating bonding MOs.

Destructive overlap

Out-of-phase (opposite sign) combination of orbital lobes that creates a node between nuclei, producing antibonding MOs.

Non-bonding interaction

An orbital combination that yields neither stabilization nor destabilization; electrons in such orbitals do not affect bond order.

He₂ molecule (hypothetical)

A non-existent diatomic species where both σ1s and σ1s* orbitals would be filled, giving bond order 0 and resulting in no stable bond.