• Evidence of ions: electron density maps, melting/boiling points, electrical conductivity, solubility, and electrolysis all provide proof that ions actually exist.
• Ionic bonding & structure: how electrons transfer from metals to non-metals to form oppositely charged ions, which then pack into giant ionic lattices held together by electrostatic attraction.
• Dot-and-cross diagrams: how to draw the electron transfer for compounds like NaCl, MgO, CaF₂, Li₃N, and Al₂O₃.
• Ionic bond strength: how atomic radius and ionic radius change across periods and groups, and why this affects bond strength (charge density).
• Polarisation: how a cation can distort an anion's electron cloud, introducing "covalent character" into an otherwise ionic bond.
• Metallic bonding & lattices: the "sea of delocalised electrons" model, and how it explains metals' physical properties (conductivity, malleability, high melting points).

Giant Ionic Lattices

Think of an ionic bond like a very committed, permanent handshake between a positively charged metal ion (a cation) and a negatively charged non-metal ion (an anion). The metal atom has given away electrons (becoming positive), and the non-metal has taken them (becoming negative).

Unlike a covalent bond, which only pulls two specific atoms together, an ionic bond pulls in every direction at once. This means ions don't just pair up — they stack into a huge, repeating 3D pattern called a giant ionic lattice (also called a giant ionic structure). Every positive ion is surrounded by negative ions, and vice versa, like a 3D checkerboard.

Electron density maps are one piece of hard evidence for ions. These maps show where electrons are most likely to be found. For NaCl, the map shows two separate "islands" of electron density (one smaller, around Na⁺, one larger, around Cl⁻) with the density dropping to zero in between. This gap is exactly what you'd expect if electrons have been completely transferred from Na to Cl — there's no "shared" electron cloud like you'd see in a covalent bond.

Melting & Boiling Point

Most ionic compounds are solid at room temperature, and they need a lot of heat to melt or boil. Why? Because you're not just wiggling a few atoms loose — you're trying to rip apart a giant lattice held together by strong electrostatic forces acting in every direction. That takes serious energy.

The bigger the ionic charge, the higher the melting point, because the electrostatic attraction is stronger. Compare:

Notice CaO's melting point is over twice as high. This isn't about ion size (Na⁺ and Ca²⁺ are actually similar in size) — it's purely because CaO has a 2+ and 2− pairing instead of Na₂O's 1+ and 2−, so the electrostatic pull is much stronger.

Electrical Conductivity

For any substance to conduct electricity, it needs freely moving charged particles — either electrons or ions.

• Solid ionic compounds do NOT conduct — the ions are locked in fixed positions in the lattice, vibrating but not moving around.
• Molten or dissolved ionic compounds DO conduct — once melted or dissolved, the ions are free to move and carry charge to the electrodes.

Solubility

Many ionic compounds dissolve in polar solvents like water. This happens in two steps:

1. The polar water molecules attack and break down the lattice.
2. Water molecules then surround each ion — the δ+ (slightly positive) hydrogen ends of water cluster around negative anions, and the δ− (slightly negative) oxygen ends cluster around positive cations.

This process is called hydration, and the bonds formed between water and the ions are called ion-dipole bonds.

Electrolysis — The Clearest Proof

This is probably the most convincing evidence of all. In electrolysis, positive ions migrate toward the negative electrode (cathode) and negative ions migrate toward the positive electrode (anode). If ions didn't exist, there'd be no reason for this directional movement to happen.

A brilliant visual example uses copper(II) chromate(VI), CuCrO₄, which is olive green in solution (a mix of blue Cu²⁺ and yellow CrO₄²⁻). When electrolysed:

• Blue colour (Cu²⁺) creeps toward the negative electrode.
• Yellow colour (CrO₄²⁻) creeps toward the positive electrode.

You can literally watch the two colours separate and migrate in opposite directions — visual, undeniable proof that oppositely charged particles exist and move independently.