Explain the increasing solubility of Group 2 hydroxides
Compare lattice and hydration changes to account for the hydroxide solubility trend down Group 2.
A Level; Cambridge International Chemistry 9701, 2028–2030 syllabus, Topic 27. Compare Group 2 ions at the same charge and distinguish lattice formation, lattice dissociation, hydration and solution enthalpy. Trend explanations are qualitative unless data are supplied. Written practice is self-checked
Before you begin
- Use AS Group 2 reactivity trends, ionic charge and radius, Hess cycles, and balanced equations.
Work on paper. Try each question before revealing help, and explain why your method works.
Quick readiness check
What is the Group 2 hydroxide solubility trend down the group?
What you will learn
- State how Group 2 hydroxide solubility changes down the group
- Compare changes in lattice and hydration enthalpy
- Explain why the net dissolution becomes more favourable qualitatively
Hydroxide solubility increases from magnesium to barium
The hydroxides become more soluble down Group 2. Magnesium hydroxide is sparingly soluble, while the lower-group hydroxides dissolve more readily; use the qualitative order unless a question supplies solubility data.
The formula M(OH)₂ stays in a 1:2 ratio. Compare the same type of solid across the group so that the cation size is the changing factor.
Balance lattice and hydration changes
As M²⁺ becomes larger, the electrostatic attraction to OH⁻ in the lattice weakens, so lattice dissociation requires less energy. Hydration of the larger cation is also less exothermic, which opposes dissolution.
For hydroxides, the lattice-dissociation enthalpy becomes less endothermic by more than the hydration enthalpy becomes less exothermic. The signed sum, ΔHsol, therefore becomes more negative (more exothermic or less endothermic) down the group, consistent with increasing solubility. State both changes; do not claim hydration becomes more exothermic for larger ions.
Pause and explain
What happens to hydration enthalpy of M²⁺ as the cation gets larger?
Worked example
Explain why Ba(OH)₂ is more soluble than Mg(OH)₂.
Show the worked solution
- Ba²⁺ has a larger radius than Mg²⁺.
- The hydroxide lattice is easier to separate for the larger cation.
- Although Ba²⁺ hydration is less exothermic, the lattice change dominates the qualitative balance.
Answer Ba(OH)₂ is more soluble. Down Group 2, the larger cation weakens the lattice attraction so less energy is needed to separate the ions. Hydration of Ba²⁺ releases less energy than hydration of Mg²⁺, but the reduction in lattice-separation cost has the greater effect for the hydroxides, making dissolution more favourable.
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Write your answer and reasoning first. Use a hint only if you are stuck. The model solution is for self-checking; your written work is not automatically marked.
Order Mg(OH)₂, Ca(OH)₂ and Ba(OH)₂ by increasing solubility.
Give me a hint
Apply the hydroxide trend from the top to the bottom of Group 2.
Compare my reasoning
- Solubility increases down the group.
- Place magnesium first and barium last.
- Keep the compound formula consistent.
Mg(OH)₂ < Ca(OH)₂ < Ba(OH)₂ in solubility.
Look for these in your work
- Group order correct
- Trend applied
- Inequality points in the right direction
A learner explains increasing hydroxide solubility by saying both lattice dissociation and hydration become less endothermic down the group. Correct the explanation.
Give me a hint
Check the process direction and the sign of hydration for gaseous ions becoming aqueous ions.
Compare my reasoning
- Hydration is exothermic in the stated direction.
- For a larger cation it becomes less exothermic.
- Lattice dissociation becomes less endothermic; compare the net balance.
Lattice dissociation becomes less endothermic as the larger cation weakens the lattice. Hydration is exothermic but becomes less exothermic for a larger cation. For the hydroxides, the fall in lattice-separation cost has the greater qualitative effect, so solubility increases.
Look for these in your work
- Two processes distinguished
- Hydration sign and trend corrected
- Dominant balance linked to solubility
A data set shows a hydroxide dissolving with ΔHsol changing from +15 kJ mol⁻¹ for Mg(OH)₂ to −5 kJ mol⁻¹ for Ba(OH)₂. What does this support, and what does it not prove?
Give me a hint
Interpret the enthalpy trend and keep it separate from a direct numerical solubility measurement.
Compare my reasoning
- The solution enthalpy becomes more favourable down the group.
- This is consistent with increasing hydroxide solubility.
- ΔHsol values alone do not give exact solubilities.
The increasingly negative ΔHsol supports a more favourable enthalpy contribution to dissolution and is consistent with hydroxide solubility increasing down the group. It does not by itself give the exact solubility or full equilibrium constant.
Look for these in your work
- Data trend interpreted
- Consistency stated
- Limit of inference stated
Common mistakes
- Mentioning only lattice energy and ignoring hydration
- Saying hydration of a larger ion is more exothermic
- Reversing the hydroxide solubility trend
- Treating a qualitative explanation as a numerical ΔHsol calculation
What can you explain now?
What is the Group 2 hydroxide solubility trend down the group?
Compare with the explanation
It increases
The larger cations make dissolution more favourable overall.
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Make it stick
Why does increasing hydroxide solubility require comparing both lattice separation and ion hydration?
If you needed the solution, close it and solve the problem again from a blank page. A correct answer is stronger when you can explain the reason for each step.