3.2.2 Group 2

When you move from Mg to Ba, keep the electronic trends clear: atomic and ionic radii increase and first ionisation energy decreases. The melting points are less tidy, so I would not try to force them into a perfectly smooth sequence.

For the size/IE trends, each step down adds a shell and increases shielding. For melting point, stay with metallic structure and the changing attraction between M²⁺ ions and delocalised electrons. There is no useful simple monotonic rule to invent.

If you explain the Group 2 radius and ionisation-energy trends, start with the extra occupied shell each step down. The outer electron is farther from the nucleus and more shielded, so the attraction it feels is weaker and first ionisation energy falls.

The same shell addition makes atomic and M²⁺ ionic radii larger. Increased nuclear charge does not reverse the trend because distance and shielding also increase.

For Group 2 melting points, keep the explanation tied to metallic structure. All the metals have lattices of 2+ ions and delocalised electrons, but changing ionic size and crystal structure means the measured melting points do not follow a perfectly smooth trend.

The measured melting points are not a neat straight trend. If data are supplied, describe what the data actually do and then relate the changes to metallic attraction; do not force a memorised “increases” or “decreases” onto every pair.

Compare the Group 2 metals with water: the reactivity increases down the group. Ca, Sr and Ba form M(OH)₂ and H₂ with cold water increasingly vigorously, while Mg reacts only slowly with cold water and much more readily with steam.

The trend follows easier loss of the two outer electrons as radius and shielding increase. For steam, magnesium forms MgO + H₂, so keep the cold-water and steam products distinct.

In titanium extraction, magnesium reduces TiCl₄ to titanium: TiCl₄ + 2Mg → Ti + 2MgCl₂. Mg goes from oxidation state 0 to +2, so it is oxidised and acts as the reducing agent.

If you are asked to justify the term “reducing agent”, follow the electrons: magnesium supplies the electrons that reduce Ti(IV) to Ti(0).

Group 2 hydroxides become more soluble down the group. Mg(OH)₂ is only sparingly soluble, while Ba(OH)₂ is much more soluble, so saturated solutions become more alkaline down the group as more OH⁻ enters solution.

Keep “more soluble” separate from “stronger base”. The important course trend here is how much hydroxide dissolves and therefore the concentration of OH⁻ that can be produced.

Mg(OH)₂ is sparingly soluble: most remains as solid, but a small amount dissolves to give Mg²⁺ and OH⁻ ions. That is enough for it to behave as a base even though you would not describe it as freely soluble.

This wording matters because “insoluble” would hide the small equilibrium concentration of ions that gives the suspension its alkaline behaviour.

When a question asks why these hydroxides are useful, connect the use directly to neutralisation. Mg(OH)₂ neutralises excess stomach acid as an antacid, while Ca(OH)₂ neutralises acidic soils and raises their pH.

In both cases the useful chemistry is acid–base neutralisation by OH⁻. If an explanation is requested, connect the application directly to that neutralisation reaction. Naming the product alone is incomplete.

Basic CaO or CaCO₃ can remove acidic SO₂ from flue gases. The sulfur dioxide is converted into a calcium sulfite/sulfate-containing solid, so less SO₂ is released to the atmosphere.

With CaCO₃, neutralisation also releases CO₂. You do not need a vague pollution statement here: identify the acid–base reaction that traps the sulfur in a solid calcium compound.

Group 2 sulfate solubility decreases down the group: MgSO₄ is soluble and BaSO₄ is extremely insoluble. This is the opposite direction to the Group 2 hydroxide solubility trend.

I’d learn those two trends together because swapping them is an easy mistake: hydroxides become more soluble down the group; sulfates become less soluble.

BaSO₄ is effectively insoluble in water. That very low solubility is central both to the white precipitate in the sulfate test and to its use as a barium-containing contrast material.

The safety point is specifically the low concentration of free Ba²⁺ released from BaSO₄. Soluble barium compounds are toxic, so “contains barium” by itself does not tell you the risk.

For sulfate ions, acidify the sample and add barium chloride solution. A white precipitate of BaSO₄ is a positive result: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).

Write the ionic equation if asked and make the observation explicit. “It goes cloudy” is less precise than identifying the white precipitate.

BaSO₄ is insoluble and has a medical use as a barium contrast material for imaging the digestive tract. Its very low solubility means very little Ba²⁺ is released while it passes through the gut.

I’d learn the pairing simply: insoluble BaSO₄ — medical contrast material. You do not need a separate account of X-ray physics for this chemistry point.

For the sulfate test, give the two reagents separate jobs. Ba²⁺ is there because sulfate gives the very insoluble white precipitate BaSO₄; acidifying first removes ions such as carbonate that could otherwise give a misleading barium precipitate.

The acid therefore improves selectivity; the barium ion supplies the precipitating ion. Keep those jobs separate when the question asks why each reagent is used.