The table shows the melting points of some ionic compounds.
| Compound | Ions present | Melting point in °C |
|---|---|---|
| sodium fluoride | Na+\mathrm{Na^{+}}Na+ and F−\mathrm{F^{-}}F− | 993 |
| sodium chloride | Na+\mathrm{Na^{+}}Na+ and Cl−\mathrm{Cl^{-}}Cl− | 801 |
| sodium bromide | Na+\mathrm{Na^{+}}Na+ and Br−\mathrm{Br^{-}}Br− | 747 |
| sodium iodide | Na+\mathrm{Na^{+}}Na+ and I−\mathrm{I^{-}}I− | 661 |
| magnesium oxide | Mg2+\mathrm{Mg^{2+}}Mg2+ and O2−\mathrm{O^{2-}}O2− | 2852 |
Describe the trend in the melting points of the four sodium halides.
The halide ions get larger from fluoride to iodide. Suggest an explanation for the trend in the melting points of the sodium halides.
An oxide ion is about the same size as a fluoride ion. Explain why the melting point of magnesium oxide is much higher than the melting point of sodium fluoride.
A potassium ion is larger than a sodium ion. Predict how the melting point of potassium chloride compares with the melting point of sodium chloride. Justify your prediction.
A student writes: "When sodium chloride melts, the covalent bonds between the sodium atoms and the chlorine atoms break." Explain why the student's statement is incorrect.
56 exam-style questions on Edexcel GCSE Chemistry 2.3 Ionic bonding, covering 2.3.1 Ions and the formation of ionic bonds, 2.3.2 Naming and deducing formulae of ionic compounds, and 2.3.3 Giant ionic lattices. Each one has a worked solution and a mark scheme showing where the marks go.