Unit 1 Deep Dive: Periodicity, Valence, and Ionic Bonding

Periodic trends

Repeating patterns in element properties across the periodic table that can be predicted using nuclear charge, shielding, distance, and electrostatic attraction.

Nuclear charge

The positive charge of the nucleus (from protons) that attracts electrons; increases as the number of protons increases.

Shielding (screening)

Reduction of the nucleus’s pull on valence electrons due to repulsion from inner (core) electrons.

Effective nuclear charge (Zeff)

The net positive charge “felt” by a valence electron after accounting for shielding; generally increases across a period.

Coulomb’s law (qualitative use)

Electrostatic force increases with greater charge and decreases strongly as distance increases (F ∝ q1q2/r^2), used to explain attraction trends.

Atomic radius

A measure of atom size inferred from bonding distances; generally decreases across a period (higher Zeff) and increases down a group (higher energy levels and shielding).

Ionic radius

The effective size of an ion in an ionic crystal; cations are smaller than their atoms, and anions are larger than their atoms.

Cation

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

Anion

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

Isoelectronic series

A set of species with the same number of electrons; within the series, the species with more protons is smaller (greater attraction).

Ionization energy (IE)

Energy required to remove an electron from a gaseous atom or ion; reflects how tightly an atom holds its outermost electron.

First ionization energy

Energy required to remove the first electron from a neutral gaseous atom: X(g) → X⁺(g) + e⁻.

Electron affinity (EA)

Energy change when a gaseous atom gains an electron: X(g) + e⁻ → X⁻(g); more favorable EA generally means more energy released when gaining an electron.

Electronegativity (EN)

An atom’s ability to attract shared electrons in a chemical bond; increases across a period and decreases down a group (fluorine is highest).

Metallic character

How strongly an element behaves like a metal (tends to lose electrons and form cations); increases down a group and decreases across a period.

Valence electrons

Electrons in the outermost occupied energy level (main-group elements); they control bonding, reactivity, and common ion charges.

Lewis dot diagram

A representation using dots around an element symbol to show only valence electrons and track electrons in bonding.

Ionic compound

A substance made of cations and anions held together by electrostatic attraction in a repeating 3D structure (not discrete molecules).

Ionic lattice

The repeating 3D crystal structure of an ionic solid where each ion is attracted to many neighboring opposite charges.

Empirical formula (for ionic compounds)

The simplest whole-number ratio of ions in an ionic lattice (e.g., “NaCl” represents a ratio, not a single molecule).

Electrical neutrality (charge balance)

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

Lattice energy

Energy associated with separating one mole of an ionic solid into gaseous ions (or the magnitude of energy released when the lattice forms); increases with higher ionic charges and smaller ionic radii.

Electrolyte

A substance that conducts electricity when molten or dissolved in water because its ions are free to move (ionic solids do not conduct).

Hydration (ion–dipole) attraction

Attraction between dissolved ions and polar water molecules; competes with lattice energy and helps determine ionic solubility.

Brittleness (of ionic solids)

Tendency of ionic crystals to cleave when shifted so like charges line up, causing strong repulsion and fracture.