Chemistry • Year 11 • Module 1 • Lesson 7
Ionic Bonding and Properties
Apply your understanding of lattice structure, lattice energy, and conductivity to real data, comparative scenarios, and an Australian industrial context.
1. Interpret data — comparing ionic compounds
The table below gives properties of five ionic compounds. Use the data to answer the questions that follow. 9 marks
| Compound | Ions (charges) | Melting point (°C) | Conducts as solid? | Conducts when dissolved? |
|---|---|---|---|---|
| NaCl | Na¹¹ (+1), Cl¯ (−1) | 801 | No | Yes |
| MgO | Mg²¹ (+2), O²¯ (−2) | 2852 | No | Yes (sparingly soluble) |
| LiF | Li¹¹ (+1), F¯ (−1) | 848 | No | Yes |
| CsI | Cs¹¹ (+1), I¯ (−1) | 632 | No | Yes |
| CaO | Ca²¹ (+2), O²¯ (−2) | 2613 | No | Slightly |
1.1 Explain why none of the five compounds conducts electricity as a solid, using the term “lattice” in your answer. 2 marks
1.2 Compare the melting points of NaCl, LiF, and CsI. All three have ions with charges of ±1. Using ion size reasoning, explain the trend in melting point from LiF to NaCl to CsI. 3 marks
1.3 MgO and CaO both have ions with charges of ±2, yet MgO has a higher melting point. Account for this difference. 2 marks
1.4 A student says: “NaCl conducts when dissolved because the ions start moving only when water hits them.” Identify one flaw in this explanation and rewrite it correctly. 2 marks
2. Interpret graph — melting point vs ion charge
The bar chart below shows the melting points of six ionic oxides and chlorides. 7 marks
Figure 1. Melting points of selected ionic compounds. Data from standard reference sources.
2.1 Describe the overall pattern in the graph relating ion charge to melting point. 2 marks
2.2 MgF2 contains Mg2+ (+2) and F− (−1) ions. Its melting point of 1263 °C is between the ±1 group and the ±2 group. Using Coulomb’s Law reasoning, explain this intermediate position. 2 marks
2.3 A scientist needs a ceramic lining for a furnace that operates at 1800 °C. Based only on the graph, which compound(s) could be used? Justify your answer. 3 marks
3. Compare conductivity across three states
Complete the three-column table below for solid NaCl, molten NaCl, and aqueous NaCl solution. For each cell, write a concise description explaining the observation. 9 marks (1 per cell)
| Feature | Solid NaCl | Molten NaCl | NaCl(aq) solution |
|---|---|---|---|
| Electrical conductivity | |||
| State of the ions | |||
| Charge carrier |
4. Predict and justify — Olympic Dam scenario
Olympic Dam in South Australia is one of the world’s largest copper-uranium mines. Copper is extracted by first dissolving copper ore (containing copper(II) sulfate, CuSO4) in water, then passing an electric current through the solution to deposit pure copper metal at a cathode electrode. This process is called electrowinning.
5 marks
4.1 Explain why an aqueous CuSO4 solution can carry an electric current, using terms “mobile ions” and “hydration” in your answer. 3 marks
4.2 An engineer suggests that solid CuSO4 powder could be used instead of the aqueous solution to carry the current in electrowinning. Predict whether this would work and justify your prediction. 2 marks
Q1.1 — Why none conduct as solids (2 marks)
In all ionic compounds, the ions are locked in fixed positions within the ionic lattice [1]. They cannot move, so there are no mobile charge carriers. Electrical conductivity requires movement of charged particles, which is not possible when ions are held rigidly in the lattice [1].
Q1.2 — LiF vs NaCl vs CsI melting points (3 marks)
All three compounds have ±1 ion charges, so charge alone does not explain the trend [1]. The key factor is ion size: Li+ and F− are the smallest ions, so they can pack most closely together, giving the strongest electrostatic attraction and the highest lattice energy — hence LiF has the highest melting point (848 °C) [1]. Cs+ and I− are the largest ions, so they are more distant from each other, giving the weakest attraction and the lowest melting point (632 °C) among the three [1].
Q1.3 — MgO vs CaO (2 marks)
Both compounds contain ±2 ions, so charge is equal [1]. Mg2+ is smaller than Ca2+, so the ions in MgO can get closer together, resulting in a stronger electrostatic attraction, higher lattice energy, and a higher melting point than CaO [1].
Q1.4 — Flaw in student explanation (2 marks)
Flaw: The ions in NaCl do not “start moving” only when touched by water; the ions already exist as Na+ and Cl− in the solid lattice [1]. Correct explanation: When NaCl dissolves, polar water molecules (via their δ+ and δ− ends) attract and surround the Na+ and Cl− ions (hydration), pulling them off the lattice and releasing them as independent mobile ions in solution; these mobile ions then carry charge when a voltage is applied [1].
Q2.1 — Graph pattern (2 marks)
Compounds with ions of higher charge (±2: MgO, CaO) have much higher melting points than compounds with ±1 ion charges (NaCl, LiF, CsI) [1]. This is because higher charges produce stronger electrostatic attraction between ions, leading to higher lattice energies that require more thermal energy to overcome [1].
Q2.2 — MgF2 intermediate position (2 marks)
In MgF2, the force between Mg2+ (+2) and each F− (−1) is proportional to (2×1) = 2 [1]. This is greater than the (1×1) = 1 interaction in NaCl or CsI, but less than the (2×2) = 4 interaction in MgO. The intermediate charge product gives an intermediate lattice energy and therefore an intermediate melting point between the ±1 and ±2 groups [1].
Q2.3 — Furnace material at 1800 °C (3 marks)
A furnace operating at 1800 °C requires a material with a melting point above 1800 °C [1]. From the graph, MgO (2852 °C) and CaO (2613 °C) both have melting points well above 1800 °C and could be used as refractory linings [1]. NaCl (801 °C), LiF (848 °C), CsI (632 °C), and MgF2 (1263 °C) all have melting points below 1800 °C and would melt inside the furnace, making them unsuitable [1].
Q3 — Conductivity comparison table
Electrical conductivity: Solid — None; Molten — Excellent; Aqueous — Excellent.
State of the ions: Solid — Fixed, immobile in the rigid lattice; Molten — Free to move independently as the lattice has broken down; Aqueous — Free to move independently, surrounded by water molecules (hydrated).
Charge carrier: Solid — None (no mobile carriers); Molten — Na+ and Cl− ions moving through the liquid; Aqueous — Hydrated Na+ and Cl− ions moving through the solution.
Q4.1 — CuSO4(aq) conductivity (3 marks)
When CuSO4 dissolves in water, the polar water molecules attract the Cu2+ and SO42− ions on the surface of the crystal, pulling them off the lattice (hydration) [1]. The ions become surrounded by water molecules and exist as independent mobile ions dispersed throughout the solution [1]. When a voltage is applied, these mobile ions move toward oppositely charged electrodes, constituting an electric current — Cu2+ toward the cathode and SO42− toward the anode [1].
Q4.2 — Solid CuSO4 in electrowinning (2 marks)
This would not work [1]. In solid CuSO4, the Cu2+ and SO42− ions are locked in fixed positions in the ionic lattice and cannot move. Without mobile charge carriers, no electric current can pass through the solid and no copper would be deposited at the cathode [1].