The ionic model

An atom (or group of atoms) with an overall **charge** because it has **lost or gained electrons**.

A **positive** ion, formed when an atom **loses** electrons (more protons than electrons).

A **negative** ion, formed when an atom **gains** electrons (more electrons than protons).

**Cations** — metals **lose** their outer electrons to form **positive** ions.

**Anions** — non-metals **gain** electrons to form **negative** ions.

To reach a **full outer shell** — the stable **noble-gas configuration** of a Group 18 atom.

**1+, 2+, 3+** — these metals lose 1, 2 or 3 outer electrons.

**3−, 2−, 1−** — these non-metals gain 3, 2 or 1 electrons.

The **electrostatic attraction between oppositely charged ions** (a cation and an anion).

Group 17, so it **gains 1** electron → a **1−** ion (reaching 2, 8, 8).

**Na⁺** (sodium loses 1 e⁻) and **Cl⁻** (chlorine gains 1 e⁻).

Ionic = **attraction between charged ions** (electrons transferred); covalent = a **shared pair** of electrons.

A **positively** charged ion (a metal, or NH_{4}⁺).

A **negatively** charged ion (a non-metal, or a polyatomic ion).

A charged group of **bonded atoms** that acts as a single unit (e.g. SO_{4}²⁻, NO_{3}⁻).

The ratio of ions is chosen so the **total positive and negative charges cancel** to zero.

Write each ion with its charge, **cross over** the charge sizes as subscripts, then **simplify** to the smallest whole-number ratio.

**MgCl_{2}** — Mg²⁺ needs two Cl⁻ to balance.

**Al_{2}O_{3}** — crossover of Al³⁺ and O²⁻ (6+ balances 6−).

When a **polyatomic ion appears two or more times**, e.g. Ca(NO_{3})_{2}, (NH_{4})_{2}SO_{4}.

**SO_{4}²⁻** — a 2− polyatomic ion.

**NH_{4}⁺** — a 1+ polyatomic cation.

Cation name first, then the non-metal anion ending in **-ide** (e.g. magnesium ox**ide**).

**Ca_{3}N_{2}** — Ca²⁺ and N³⁻ crossover (6+ balances 6−).

A regular, repeating **3-D array** of oppositely charged ions, with each ion surrounded by ions of the opposite charge.

**Strong electrostatic forces of attraction** between the oppositely charged ions (this is the ionic bond).

Many **strong electrostatic attractions** between the ions must be broken, which needs a **large amount of energy**.

**Higher ionic charge** and **smaller ionic radius** — both increase the electrostatic attraction.

When **molten** or **dissolved in water (aqueous)** — the ions are then **free to move**. Not as a solid.

The ions are held in **fixed positions** in the lattice, so no charged particles are free to move.

A force makes layers **shift**, bringing **like charges** next to each other; they **repel** and split the crystal.

Water is **polar**: its δ⁻ oxygen attracts cations and δ⁺ hydrogens attract anions, pulling ions out of the lattice (hydration).

Solid = ions **fixed**, does **not** conduct. Molten = lattice broken, ions **free to move**, **conducts**.

High melting point + does **not** conduct as a solid + **conducts when molten/aqueous** = ionic.

Mg^{2+} and O^{2−} carry **higher charges** than Na^{+} and Cl^{−}, so the electrostatic attraction is much stronger.