AS Chemistry: Group 2 — Revision Notes

Condensed for the final weeks. For the full explanation, use the Group 2: The Alkaline Earth Metals study guide.

Reactivity down the group

Increases down the group: atomic radius grows and shielding increases, so the two outer electrons are lost more readily despite the greater nuclear charge.

metal + oxygen  ->  metal oxide          2Mg + O2 -> 2MgO
metal + water   ->  hydroxide + hydrogen Ca + 2H2O -> Ca(OH)2 + H2
metal + steam   ->  oxide + hydrogen     Mg + H2O -> MgO + H2   (Mg only)

Magnesium reacts only very slowly with cold water but vigorously with steam — and note it gives the oxide with steam, the hydroxide with cold water.

The special case of barium and dilute sulfuric acid

Ba + H2SO4  ->  BaSO4 + H2      but the reaction STOPS almost immediately

Barium sulfate is insoluble and forms a coating on the metal surface, physically blocking further acid contact — this is a solubility effect, not a reactivity effect, and is a common source of confusion since barium is the most reactive Group 2 metal, yet appears to react the least with this specific acid.

Reactions of oxides, hydroxides and carbonates with dilute acid

The oxide, hydroxide and carbonate of any Group 2 metal all follow the same base/carbonate-plus-acid logic met elsewhere in the course:

MO      + 2HCl  ->  MCl2 + H2O
M(OH)2  + 2HCl  ->  MCl2 + 2H2O
MCO3    + 2HCl  ->  MCl2 + H2O + CO2

(and equivalently with dilute H₂SO₄ — again limited for barium compounds by the insolubility of the BaSO₄ produced).

The sulfate-ion test

Adding aqueous barium chloride and then dilute hydrochloric acid to an unknown solution is the standard qualitative test for sulfate ions:

Ba2+(aq) + SO4 2-(aq)  ->  BaSO4(s)     dense white precipitate

The dilute HCl is added first to remove any carbonate ions present (which would otherwise also give a white precipitate with Ba²⁺ and cause a false positive) — only a genuine sulfate gives a precipitate that persists once the acid has been added.

The two opposing solubility trends

This is the pair candidates most often reverse.

Compound	Trend down the group	Example
Hydroxides	Solubility INCREASES	Mg(OH)₂ sparingly soluble → Ba(OH)₂ soluble
Sulfates	Solubility DECREASES	MgSO₄ soluble → BaSO₄ insoluble

Consequences worth knowing:

• Mg(OH)₂ — milk of magnesia, an antacid, precisely because it is only sparingly soluble.
• BaSO₄ — barium meal for X-rays; safe despite barium’s toxicity because it is insoluble.
• Ca(OH)₂ — limewater; the CO₂ test relies on insoluble CaCO₃ forming.

Thermal stability

Increases down the group for both nitrates and carbonates.

carbonate:  MCO3  ->  MO  + CO2
nitrate:    2M(NO3)2 -> 2MO + 4NO2 + O2      (brown NO2 gas)

Why (A Level explanation, useful context but not required at AS — AS 10.1.3 only requires describing the trend with equations): the cation becomes larger down the group, so its charge density falls. A smaller, more polarising cation distorts the anion’s electron cloud more, weakening the C–O or N–O bond and lowering the decomposition temperature.

So magnesium compounds decompose most easily; barium compounds are the most stable.

At AS, describing the trend correctly with the equations above is sufficient; the polarising-power mechanism is A Level Topic 27.1.1 material — “because it is bigger” alone would still be an incomplete explanation if an explanation is asked for, but AS itself does not require the mechanism.

Uses

• Ca(OH)₂ — neutralises acidic soils.
• CaCO₃ — removes SO₂ from flue gases; makes cement and iron.
• MgO — refractory lining for furnaces (very high melting point).
• BaSO₄ — barium meal.

Exam traps

• Reversing the hydroxide and sulfate solubility trends.
• Saying magnesium reacts vigorously with cold water — it does not; only with steam.
• Explaining thermal stability without mentioning charge density or polarising power.
• Forgetting nitrate decomposition gives brown NO₂.
• Writing Group 2 ions as 1+ — they are always 2+.
• Concluding barium is unreactive towards dilute sulfuric acid because the reaction stops quickly — it stops due to the insoluble BaSO₄ coating, not because barium itself is unreactive (it is the most reactive Group 2 metal).
• Forgetting to acidify with dilute HCl before the sulfate test, risking a false positive from carbonate ions.

Self-test

1. Why does reactivity increase down Group 2?
2. State the solubility trends for the hydroxides and the sulfates.
3. Why is barium sulfate safe for medical use?
4. Explain why BaCO₃ is more thermally stable than MgCO₃.
5. What products form when calcium nitrate is heated?
6. Explain why the reaction between barium and dilute sulfuric acid stops almost immediately, despite barium being the most reactive Group 2 metal.
7. Describe the test for sulfate ions in solution, including why dilute HCl is added first.

Answers: 1. The outer electrons are further from the nucleus and better shielded, so they are more easily lost despite the increased nuclear charge. 2. Hydroxide solubility increases down the group; sulfate solubility decreases. 3. It is insoluble, so the toxic Ba²⁺ ions are not absorbed by the body. 4. Ba²⁺ is larger with lower charge density and therefore less polarising power, so it distorts the carbonate ion’s electron cloud less, weakening the C–O bond less and requiring a higher temperature to decompose. 5. Calcium oxide, brown nitrogen dioxide and oxygen. 6. The insoluble BaSO₄ product coats the metal surface, physically blocking further contact between the acid and the unreacted barium beneath — a solubility effect, not a reactivity one. 7. Add aqueous barium chloride and dilute hydrochloric acid; a dense white precipitate (BaSO₄) confirms sulfate ions are present. The HCl is added to remove carbonate ions first, since they would otherwise also give a white precipitate with Ba²⁺ and produce a false positive.

For the full reaction equations, including the elements’ reactions with oxygen and acids, see the Group 2: The Alkaline Earth Metals study guide.