Imagine stacking a billion tiny magnets in a perfectly alternating grid — positive, negative, positive, negative — in all three dimensions. The electrostatic attraction holds everything together with enormous force. Now imagine nudging one row sideways. Suddenly like charges align. The whole thing shatters. This is why your table salt is hard, brittle, and melts at 801°C.
Core Content
Wrong: Ionic compounds conduct electricity in the solid state because they contain charged ions.
Right: Ionic compounds only conduct electricity when molten or dissolved in water. In the solid state, the ions are locked in a fixed lattice and cannot move. Conductivity requires mobile charge carriers, which are only present when the lattice breaks down.
Ionic compounds do not form discrete molecules. Instead, ions arrange themselves into a giant regular 3D lattice where every cation is surrounded by anions and every anion is surrounded by cations. The arrangement maximises attractive forces and minimises repulsive forces.
The lattice energy is the energy holding the ions together. It depends on ion charge and ion size:
Using the principles of lattice energy (charge and size), you can predict and explain differences between ionic compounds:
Worked Examples
| Identify the ions in each compoundNaCl: Na⁺ (+1) and Cl⁻ (−1). MgO: Mg²⁺ (+2) and O²⁻ (−2). The key difference is the charge on the ions — NaCl ions have charges of ±1; MgO ions have charges of ±2. | The charge difference is the most important factor. Always identify and compare charges before anything else when comparing lattice energies. |
| Apply Coulomb's Law reasoningElectrostatic force ∝ (charge₁ × charge₂) / distance². For NaCl: force ∝ (1×1) = 1. For MgO: force ∝ (2×2) = 4. So the electrostatic attraction in MgO is approximately four times stronger than in NaCl (assuming similar ionic radii). | This is Coulomb's Law logic — you don't need to calculate it exactly. Just note that doubling the charge on both ions quadruples the force. HSC expects conceptual understanding: 'higher charge → stronger attraction → higher lattice energy.' |
| Connect to melting pointMelting requires enough energy to overcome the electrostatic forces holding ions in the lattice. MgO requires ~4× more energy to pull Mg²⁺ and O²⁻ apart compared to NaCl's Na⁺ and Cl⁻. This is reflected in the ~3.5× higher MP. | MP is directly related to lattice energy, which is directly related to ion charge and size. More energy to break → higher temperature needed → higher MP. This three-step reasoning is the full mark-gaining chain. |
| Solid NaCl — no conductivityIn solid NaCl, Na⁺ and Cl⁻ ions are locked in fixed positions in the ionic lattice. They cannot move through the solid. Electrical conductivity requires mobile charge carriers — since no ions (or electrons) can move freely, the solid does not conduct. | This is the most common exam question on this topic. The key phrase is 'ions are fixed in the lattice'. Do not say 'no free electrons' as the explanation — that is true for metals and covalent substances. For ionic compounds, the reason is specifically immobile ions in the lattice. |
| Molten NaCl — excellent conductivityAt 801°C, the lattice collapses. Na⁺ and Cl⁻ ions are now free to move as independent particles in the liquid. When a voltage is applied, cations (Na⁺) migrate toward the negative electrode and anions (Cl⁻) migrate toward the positive electrode — charge is carried through the liquid. | The ions themselves carry the charge. This is different from metallic conductors (electrons carry charge). In electrochemistry, this distinction matters: ionic current vs electronic current. |
| Dissolved NaCl — excellent conductivityWater molecules (polar) surround Na⁺ and Cl⁻ ions, pulling them off the lattice surface (hydration). The hydrated ions are mobile in solution and carry charge just as in the molten state. | The mechanism is the same as for molten NaCl: mobile ions carrying charge. The difference is that water does the 'lattice-breaking' work instead of heat. |
Activities
1 Predict whether calcium fluoride (CaF₂, Ca²⁺ and F⁻) or potassium bromide (KBr, K⁺ and Br⁻) will have a higher melting point. Justify your prediction using ion charges and size reasoning.
2 A student heats solid magnesium chloride (MgCl₂) until it melts and applies a voltage. Describe what happens to the ions and explain why the molten liquid conducts electricity.
3 The table below lists properties of three substances. Use the data to classify each as ionic compound, metal, or covalent molecular compound.
Explain the different conductivity profiles of copper and NaCl. Both conduct as liquids — why does only copper conduct as a solid?
Glucose (C₆H₁₂O₆) does not conduct in any state, yet it dissolves in water. Explain why dissolved glucose does not conduct electricity.
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
Multiple Choice
5 random questions from a replayable lesson bank — feedback shown immediately
Short Answer
6. Describe the structure of an ionic lattice using sodium chloride (NaCl) as an example. In your answer, explain what holds the lattice together and why no discrete molecules exist in NaCl. 3 MARKS
7. Compare the electrical conductivity of solid aluminium oxide (Al₂O₃, an ionic compound) and liquid aluminium (Al, a metal). Explain why both conduct as liquids but only one conducts as a solid, referring to the charge carriers in each case. 4 MARKS
8. Magnesium oxide (MgO) is used as a refractory material — a substance that withstands very high temperatures without melting. Using your knowledge of ionic structure and lattice energy, explain why MgO is well-suited to this application. In your answer, compare MgO to NaCl. 4 MARKS
1. CaF₂ will have a higher MP. Ca²⁺ has a charge of +2 while K⁺ has +1; F⁻ and Br⁻ both have −1, but the Ca²⁺/F⁻ combination produces stronger electrostatic attraction (higher charge on Ca²⁺ → higher lattice energy). Additionally, F⁻ is smaller than Br⁻, meaning ions in CaF₂ are closer together, further increasing the attraction. Both factors (higher charge on cation + smaller anion) raise the lattice energy → higher MP for CaF₂.
2. When MgCl₂ melts, the ionic lattice breaks down and Mg²⁺ and Cl⁻ ions become free to move independently in the liquid. When a voltage is applied, Mg²⁺ ions (positive) migrate toward the negative electrode (cathode) and Cl⁻ ions (negative) migrate toward the positive electrode (anode). This movement of charged particles constitutes an electric current — hence excellent conductivity.
3. X: ionic compound — high MP, no solid conductivity but excellent molten conductivity, hard and brittle. Y: covalent molecular compound — low MP (180°C), no conductivity in any state, soft and waxy. Z: metallic element — very high MP (1538°C = iron), excellent conductivity as both solid and liquid, malleable.
A: Copper conducts as a solid because it has a sea of delocalised electrons that are free to move throughout the metallic lattice at all times — no lattice disruption is needed. NaCl does not conduct as a solid because its charge carriers (Na⁺ and Cl⁻ ions) are fixed in the ionic lattice and cannot move. Both conduct as liquids: liquid copper still has delocalised electrons; molten NaCl has freed ions that can now move.
B: Glucose dissolves in water as intact polar molecules (C₆H₁₂O₆), not as ions. There are no charged particles in the glucose solution — water molecules interact with the glucose molecules via hydrogen bonding, but no ionisation occurs. Since conductivity requires mobile charge carriers (ions or free electrons), and dissolved glucose has neither, it does not conduct electricity even in solution.
1. C — Ions fixed in lattice = correct ionic explanation. A is wrong (there ARE ions; the issue is they can't move). B and D are factually wrong.
2. B — NaCl (±1) vs MgO (±2) = massive charge difference → massive MP difference. A, C, D all involve ±1 compounds with small size differences → small MP differences.
3. A — High MP + hard/brittle + no solid conductivity + conducts dissolved = all ionic hallmarks.
4. D — Higher charges → stronger Coulomb attraction → higher lattice energy → higher MP. The correct causal chain.
5. C — Water's polarity attracts and separates ions (hydration). Mobile hydrated ions carry charge. No electrons or neutral atoms are involved.
Q6 (3 marks): In NaCl, Na⁺ and Cl⁻ ions are arranged in a regular, repeating 3D pattern — each Na⁺ is surrounded by 6 Cl⁻ and each Cl⁻ is surrounded by 6 Na⁺ (1 mark). The lattice is held together by strong electrostatic forces (ionic bonds) between oppositely charged ions acting in all directions simultaneously (1 mark). No discrete molecules exist because each ion is attracted to all its nearest neighbours, not to one specific partner — the entire crystal is one giant extended structure in which the formula NaCl simply represents the simplest whole-number ratio of ions (1 mark).
Q7 (4 marks): Solid aluminium (Al) conducts electricity because it has a sea of delocalised electrons that are free to move throughout the metallic lattice — electrons are the charge carriers (1 mark). Solid Al₂O₃ does not conduct because Al³⁺ and O²⁻ ions are fixed in the rigid ionic lattice and cannot move — no mobile charge carriers are present (1 mark). Liquid aluminium conducts because the metallic structure is maintained in the molten state — delocalised electrons remain mobile (1 mark). Molten Al₂O₃ conducts because the lattice has broken down — Al³⁺ and O²⁻ ions are now free to move and carry charge; the charge carriers in this case are ions, not electrons (1 mark).
Q8 (4 marks): MgO has a very high melting point (2852°C) because Mg²⁺ and O²⁻ carry charges of ±2, producing very strong electrostatic attraction between ions and a very high lattice energy (1 mark). In comparison, NaCl (MP 801°C) has ions with charges of only ±1 — the electrostatic attraction is roughly four times weaker (applying Coulomb's Law: force ∝ charge₁ × charge₂), so much less energy is needed to disrupt the NaCl lattice (1 mark). To use a substance as a refractory material, it must not melt at operating temperatures — MgO's 2852°C MP means it remains solid in furnaces, kilns, and industrial reactors that operate at temperatures far exceeding those where NaCl would have already melted (1 mark). Additionally, the strong lattice makes MgO chemically and thermally stable under extreme conditions — it does not readily decompose or react with other materials at high temperature (1 mark).
Return to your Think First response. You should now be able to explain the large difference in melting points:
Answer questions on ionic bonding, lattice structure and physical properties before your opponents cross the line. Fast answers = faster car. Pool: lessons 1–7.
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Use the PDF for classwork, homework or revision. It includes key ideas, activities, questions, an extend task and success-criteria proof.
Sodium chloride (NaCl) and magnesium oxide (MgO) both form ionic lattices. NaCl has a melting point of 801°C, while MgO melts at 2852°C. Both are made of positive and negative ions. What could explain why MgO requires so much more energy to melt?
Before reading on, write your best answer. Consider the charges on the ions and their sizes.