The big idea: An ionic compound is not made of separate molecules. The cations and anions arrange themselves into a giant ionic lattice — a regular, repeating 3-D pattern that extends in every direction.
Each ion is surrounded by ions of the opposite charge, and the whole structure is held together by strong electrostatic forces of attraction between the oppositely charged ions. Every physical property of the compound comes from this lattice.
Sodium chloride, NaCl: a giant 3-D lattice of alternating Na+ and Cl− ions, each ion surrounded by oppositely charged neighbours. The whole solid is held by strong electrostatic attractions.
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Key terms: - Giant ionic lattice — a regular 3-D array of alternating positive and negative ions (no individual molecules). - Electrostatic force of attraction — the pull between oppositely charged particles; this is the ionic bond. - Ionic bond strength depends on the ionic charge (higher charge → stronger) and the ionic radius (smaller ions sit closer → stronger).
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To melt an ionic solid you must break the strong electrostatic attractions holding every ion to its neighbours throughout the giant lattice. Because there are so many strong attractions, a large amount of energy is needed — so ionic compounds have high melting and boiling points and are solids at room temperature.
What makes the melting point even higher?: The melting point rises when the electrostatic attraction is stronger:
- Higher ionic charges — MgO (Mg2+ and O2−) melts far higher than NaCl (Na+ and Cl−). - Smaller ionic radii — the ions sit closer, so the attraction is stronger.
More charge and smaller ions → stronger bond → higher melting point.
Conducting electricity — only when the ions can move: Electrical conductivity needs charged particles that are free to move.
- Solid ionic compound → does not conduct: the ions are locked in fixed positions in the lattice. - Molten (liquid) or dissolved in water (aqueous) → conducts: the lattice has broken up, so the ions are now free to move and carry charge.
Worked example — why salt only conducts when molten
Solid sodium chloride does not conduct electricity, but molten sodium chloride does. Explain this difference.
Solution
- Conduction needs mobile charge carriers. A substance conducts only if it contains charged particles that are free to move.
- Solid: the Na+ and Cl− ions are held in fixed positions in the lattice, so no charges can move — it does not conduct.
- Molten: heating breaks down the lattice, so the ions are free to move and carry charge — it conducts.
Final answer
In the solid the ions are fixed; when molten the lattice breaks and the ions are free to move and carry charge.
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Ionic solids are hard but brittle — a sharp knock shatters them. This is a direct result of the ordered lattice of charges.
Why ionic solids are brittle: When a force makes one layer of ions slide over the next, ions of the same charge come to lie next to each other.
These like charges repel strongly, pushing the layers apart, so the crystal cracks and shatters along that plane rather than bending.
Solubility in water: Many ionic compounds dissolve in water because water is a polar molecule. The slightly negative oxygen of water is attracted to the cations and the slightly positive hydrogens to the anions.
These ion–water attractions pull the ions out of the lattice and surround them (hydration), breaking the lattice apart so the solid dissolves. (Not all ionic compounds dissolve — it depends on how the lattice and hydration energies compare.)
Solid ionic compound
- Ions fixed in the lattice.
- Does not conduct electricity.
- Brittle — shatters when layers shift.
Molten or aqueous
- Lattice broken; ions free to move.
- Conducts electricity.
- Ions surrounded by water (if dissolved).
How this is tested: S2.1.3 turns up two ways. On Paper 1A you are given properties — a high melting point plus no conduction as a solid but good conduction when molten — and must identify the substance as ionic. On Paper 2 you explain a property (melting point, conductivity or brittleness) directly from the lattice.
The markers always want the property linked back to the ions: the strong electrostatic attractions, or whether the ions are free to move.
Always name the cause: Never just state the property. For each mark, name the reason: strong electrostatic attractions (melting point), ions free to move / fixed (conductivity), or layers shift so like charges repel (brittleness).
IB-style question — explain two properties (a)
(a) Magnesium oxide, MgO, is an ionic compound. Explain why it has a very high melting point and does not conduct electricity when solid. [2]
How to score the marks
- Mark 1 — melting point. MgO is a giant ionic lattice held by strong electrostatic attractions between Mg2+ and O2−; a large amount of energy is needed to break all of these, so the melting point is very high. (The 2+ / 2− charges make it even stronger than NaCl.)
- Mark 2 — conductivity. In the solid the ions are held in fixed positions in the lattice, so there are no charged particles free to move — it cannot conduct.
Final answer
High melting point: strong electrostatic forces between the ions need lots of energy to break. No conduction as a solid: the ions are fixed and cannot move.
IB-style question — explain brittleness (b)
(b) Explain why solid magnesium oxide is brittle and shatters when struck. [2]
How to score the marks
- Mark 1 — the layers shift. A force makes one layer of ions slide over the layer below.
- Mark 2 — like charges repel. This brings ions of the same charge next to each other; they repel strongly, splitting the layers apart so the crystal shatters.
Final answer
When struck, the layers of ions shift so that like charges line up; they repel, forcing the layers apart and cracking the crystal.