Going down Group 2, hydroxides dissolve better but sulphates dissolve worse. Students often learn the two trends by heart and then swap them. The book explains both with the same two energies. Understand which one falls faster and you will never mix them up.
Unit 17 of the PECTAA Chemistry 12 book is Group 2 Elements. Section 17.5 covers the solubility of Group 2 sulphates and hydroxides. This note is for FSc Part 2 (2nd year, Class 12) students in Lahore, Faisalabad, Sargodha, Rawalpindi, Multan, Gujranwala and the other Punjab boards. Gujrat district comes under the Gujranwala board. PECTAA writes the books. Your own board sets your paper.
What to know
Two energies decide solubility.
- Enthalpy of hydration (ΔH_hyd). Energy given out when gaseous ions are surrounded and held by water molecules. It is always exothermic, so its value is negative. Small ions with high charge density attract water strongly and have a more negative ΔH_hyd. Down a group the ions get bigger, so ΔH_hyd becomes less negative.
- Lattice enthalpy (ΔH_lattice). Energy needed to break one mole of a solid ionic compound into its gaseous ions. It is always endothermic. It depends on ion charges and sizes. When the cation and anion are of similar size, the attraction is strong and lattice enthalpy is high. When one ion is much larger than the other, the attraction is weaker and lattice enthalpy is lower.
Enthalpy of solution (ΔH_soln). The overall energy change when the solid dissolves. A more negative or less positive ΔH_soln generally favours higher solubility.
Hydroxides: solubility increases down the group. Mg(OH)₂ is sparingly soluble and Ba(OH)₂ is quite soluble. Down the group both ΔH_hyd (for M²⁺) and ΔH_lattice decrease. For hydroxides the drop in lattice enthalpy is the bigger one. The lattice becomes much easier to break, so more dissolves.
Sulphates: solubility decreases down the group. MgSO₄ is soluble and BaSO₄ is insoluble. The sulphate ion, SO₄²⁻, is large. As the cation grows, the lattice enthalpy changes only a little, because the large anion dominates the spacing. The hydration enthalpy of the cation falls more. So less energy is released on dissolving, and solubility drops.
Table 17.4 values (solubility at 298 K, mol per 100 g of water).
| Hydroxide | Solubility | Sulphate | Solubility |
|---|---|---|---|
| Mg(OH)₂ | 2 × 10⁻⁵ | MgSO₄ | 1.83 × 10⁻¹ |
| Ca(OH)₂ | 1.5 × 10⁻³ | CaSO₄ | 4.66 × 10⁻³ |
| Sr(OH)₂ | 3.4 × 10⁻³ | SrSO₄ | 7.11 × 10⁻⁵ |
| Ba(OH)₂ | 1.5 × 10⁻² | BaSO₄ | 9.43 × 10⁻⁷ |
Read the hydroxide column downwards and the numbers rise. Read the sulphate column downwards and they fall.
Key points
- Same two energies, opposite results. The deciding question is: which falls more, lattice enthalpy or hydration enthalpy?
- Hydroxides: lattice enthalpy falls more, so solubility rises.
- Sulphates: hydration enthalpy falls more, so solubility falls.
- Figure 17.4 in the book shows both trends side by side as two arrows pointing opposite ways. Sketch it in your answer: hydroxides from less soluble at Mg to more soluble at Ba, sulphates from more soluble at Mg to less soluble at Ba.
Worked examples
Example 1. Using Table 17.4, how many times more soluble is Ba(OH)₂ than Mg(OH)₂?
- Ratio = (1.5 × 10⁻²) / (2 × 10⁻⁵).
- = 0.75 × 10³ = 750 times.
Example 2. How many times more soluble is MgSO₄ than BaSO₄?
- Ratio = (1.83 × 10⁻¹) / (9.43 × 10⁻⁷).
- ≈ 1.94 × 10⁵, so close to two hundred thousand times.
Example 3. Explain why BaSO₄ is insoluble while MgSO₄ dissolves.
- SO₄²⁻ is a large anion, so changing the cation from Mg²⁺ to Ba²⁺ changes the lattice enthalpy only slightly.
- The hydration enthalpy of Ba²⁺ is much less negative than that of Mg²⁺, because Ba²⁺ is larger.
- Less energy is released on hydration, so BaSO₄ does not dissolve.
Example 4. Arrange Ca(OH)₂, Ba(OH)₂ and Sr(OH)₂ in order of increasing solubility.
- Solubility rises down the group: Ca, then Sr, then Ba.
- Order: Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂. Table 17.4 agrees: 1.5 × 10⁻³ < 3.4 × 10⁻³ < 1.5 × 10⁻².
Common mistakes in the exam
- Swapping the trends. Hydroxides rise, sulphates fall, going down the group.
- Calling hydration enthalpy endothermic. It is always exothermic. Lattice enthalpy is the endothermic one.
- Saying only one energy changes. Both decrease down the group. The trend depends on which decreases more.
- Misreading powers of ten. 1.5 × 10⁻² is bigger than 3.4 × 10⁻³. Compare exponents first.
- Explaining with size alone. Give the energy reason, not just "the ion is bigger".
Exam technique
- Define both energies with their signs before explaining any trend.
- Use one sentence per group: "Lattice enthalpy falls more than hydration enthalpy, so solubility increases."
- Quote one example pair from the book: Mg(OH)₂ and Ba(OH)₂, or MgSO₄ and BaSO₄.
- Check your board's notice for the topics in your paper.
Practice questions
Which is more soluble, CaSO₄ or SrSO₄? Answer: CaSO₄ (4.66 × 10⁻³ against 7.11 × 10⁻⁵).
Is lattice enthalpy exothermic or endothermic? Answer: Endothermic. Energy is needed to break the lattice.
Name the sparingly soluble Group 2 hydroxide given in the book. Answer: Mg(OH)₂.
What sign of ΔH_soln favours solubility? Answer: More negative, or less positive.
Why does solubility of hydroxides rise down the group? Answer: Both energies fall, but lattice enthalpy falls more, so the solid breaks up more easily.
Arrange MgSO₄, CaSO₄ and BaSO₄ in order of decreasing solubility. Answer: MgSO₄ > CaSO₄ > BaSO₄, following Table 17.4.
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