Unit 2: Molecular and Ionic Compound Structure and Properties

Structure–property relationship

The idea that a substance’s observable properties (melting point, hardness, conductivity, solubility) are determined by what particles it contains and how they are arranged and held together.

Chemical stability (lower potential energy)

Bonded arrangements form because they typically place the system at lower potential energy (more stable) than separated particles.

Coulomb’s law

Electrostatic force between charges: F = k(q1q2)/r^2; larger charges increase attraction/repulsion magnitude, and smaller distance greatly increases the force.

Potential energy curve (bond well)

Graph of potential energy vs. internuclear distance showing a minimum (stable bond); too close gives strong repulsion, too far approaches zero interaction.

Equilibrium bond length

The internuclear distance at the minimum of the potential energy curve, where attractive and repulsive forces balance (most stable separation).

Bond energy (bond enthalpy)

Energy required to break 1 mole of a specific bond in the gas phase (typically reported as an average over many molecules); deeper potential-energy well implies a stronger bond.

Ionic bonding

Bonding model where electrons are treated as transferred, forming cations and anions held together by strong electrostatic attraction in a 3D lattice (often metal + nonmetal).

Covalent bonding

Bonding model where atoms (often nonmetals) share electrons, forming discrete molecules or extended covalent networks.

Metallic bonding

Bonding model for metals in which valence electrons are delocalized across many atoms, producing conductivity and the ability to deform without shattering.

Partial ionic character

The concept that many bonds are not purely ionic or purely covalent; electrons may be shared unequally, giving the bond some ionic behavior.

Electronegativity

An atom’s tendency to attract shared electrons in a bond; large electronegativity differences often correspond to more ionic behavior, small differences to more covalent behavior.

Cation

A positively charged ion (formed when an atom loses electrons).

Anion

A negatively charged ion (formed when an atom gains electrons).

Ionic lattice

A repeating 3D arrangement of alternating cations and anions; ionic compounds exist as lattices rather than as separate “molecules.”

Formula unit

The smallest whole-number ratio of ions represented by an ionic compound’s formula (used instead of “molecule” for ionic substances).

Empirical formula (ionic compound)

The simplest whole-number ratio of ions in an ionic solid, reflecting the repeating lattice rather than discrete molecules.

Charge balance (electrical neutrality)

Rule for ionic formulas: total positive charge must equal total negative charge, producing an overall neutral compound.

Lattice energy

Energy released when gaseous ions form an ionic solid (or equivalently, energy required to separate an ionic solid into gaseous ions); increases with higher ionic charges and smaller ionic radii (shorter distances).

Brittleness (ionic crystals)

Property of ionic solids where shifting layers can align like charges, causing strong repulsion and fracture instead of deformation.

Ionic conductivity (phase dependence)

Ionic solids conduct poorly as solids (ions locked in place) but conduct when molten or aqueous because ions become mobile charge carriers.

Electron-sea model

Model of metallic bonding: a lattice of metal cations surrounded by a “sea” of delocalized valence electrons that can move through the solid.

Malleability and ductility

Metal properties explained by metallic bonding: metals can be hammered into sheets (malleable) or drawn into wires (ductile) because the electron sea maintains attraction as layers shift.

Alloy

A mixture of elements with metallic properties; often stronger/harder than pure metals because different atoms disrupt regular lattice sliding.

Substitutional alloy

Alloy formed when atoms of similar size replace host metal atoms in the lattice (e.g., brass: Cu and Zn).

Interstitial alloy

Alloy formed when much smaller atoms occupy gaps between metal atoms (e.g., steel: C in interstitial spaces of Fe).

Lewis structure

A diagram showing valence electrons, bonds, and lone pairs in molecules or polyatomic ions; used to predict connectivity, bond order, resonance, and (with VSEPR) shape/polarity.

Octet rule

Guideline that many atoms form bonds to reach 8 valence electrons around them (noble-gas-like configuration).

Duet rule

Guideline that hydrogen tends to be stable with 2 valence electrons (a filled 1s shell).

Central atom rule (Lewis)

When drawing Lewis structures, the least electronegative atom is often central (never H), then atoms are connected and electrons distributed to satisfy octets when possible.

Incomplete octet

Octet-rule exception where a stable structure has fewer than 8 electrons around an atom (commonly B with 6, and sometimes Be).

Radical (odd-electron species)

Species with an odd total number of valence electrons; cannot give all atoms complete octets (e.g., NO).

Expanded octet

Octet-rule exception where period 3 or lower elements (e.g., P, S, Cl) can have more than 8 electrons around the central atom in valid structures (period 2 elements typically do not).

Resonance structures

Multiple valid Lewis structures differing only in electron placement (not atom positions), used when one structure cannot represent electron delocalization accurately.

Resonance hybrid

The real electron distribution of a resonant species: a delocalized blend of the resonance structures (not a rapid switching between drawings).

Formal charge (FC)

Electron-accounting tool for judging Lewis structures; FC = V − N − (B/2). Preferred structures usually minimize formal charge magnitude and place negative FC on more electronegative atoms.

VSEPR theory

Valence Shell Electron Pair Repulsion theory: electron regions around a central atom arrange to minimize repulsions, determining 3D molecular shape.

Electron domain

A region of electron density around a central atom; each single, double, or triple bond counts as 1 domain, and each lone pair counts as 1 domain.

Electron geometry vs. molecular geometry

Electron geometry describes the arrangement of all electron domains (bonds + lone pairs); molecular geometry describes the arrangement of atoms only (ignore lone pairs when naming shape).

Lone-pair repulsion (bond-angle compression)

Lone pairs repel more strongly than bonding pairs, reducing bond angles from ideal values (e.g., CH4 109.5° > NH3 ~107° > H2O ~104.5°).

Sigma bond

Bond formed by end-to-end orbital overlap along the internuclear axis; every single bond is a sigma bond, and the first bond in a multiple bond is sigma.

Pi bond

Bond formed by side-by-side overlap of unhybridized p orbitals; the second bond in a double bond is pi, and the second and third bonds in a triple bond are pi.

Hybridization (domain-counting)

Model linking electron-domain geometry to orbitals: 2 domains → sp, 3 → sp2, 4 → sp3, 5 → sp3d, 6 → sp3d2.

Molecular polarity (net dipole)

Determined by bond polarity (electronegativity differences) and molecular shape; a molecule is polar if bond dipoles do not cancel by symmetry (vector sum ≠ 0).

Covalent network solid

Solid made of atoms covalently bonded in an extended lattice (e.g., diamond, quartz); typically very hard with very high melting point and poor conductivity (graphite is a notable conducting exception due to delocalized electrons in layers).

Unit cell

The smallest repeating 3D “box” of a crystal lattice that can be stacked to build the full crystalline solid.

Unit cell particle-sharing fractions

Counting rule for particles on boundaries: corner = 1/8, edge-centered = 1/4, face-centered = 1/2, completely inside = 1.

Simple cubic (SC)

Cubic unit cell with atoms only at the corners; total atoms per unit cell = 1 (8 corners × 1/8).

Body-centered cubic (BCC)

Cubic unit cell with corner atoms plus one atom in the center of the cube; total atoms per unit cell = 2 (1 from corners + 1 body atom).

Face-centered cubic (FCC)

Cubic unit cell with corner atoms plus atoms centered on each face; total atoms per unit cell = 4 (1 from corners + 3 from faces).

Crystal density equation

Density of a cubic unit cell: ρ = (Z·M)/(NA·a^3), where Z = atoms per unit cell, M = molar mass, NA = Avogadro’s number, and a = edge length (in cm).