Group 2: Alkaline Earth Metals

Group 2 of the Periodic Table — beryllium, magnesium, calcium, strontium, barium and radium — provides the cleanest demonstration of a vertical reactivity trend in the s-block. Down the group, the outer ns² electrons sit further from the nucleus, ionisation energies fall, and metals lose their two valence electrons more readily. The visible consequences are dramatic: magnesium fizzes feebly in cold water but burns with a blinding white flame in steam, while calcium, strontium and barium evolve hydrogen vigorously at room temperature, leaving alkaline hydroxide solutions of increasing pH. Two solubility patterns run in opposite directions — hydroxides become more soluble, sulfates become less soluble — and both can be rationalised by the relative sensitivities of lattice and hydration enthalpies to cation size. Industrially the group sustains construction (CaCO₃, CaO), agriculture (Ca(OH)₂ liming), flue-gas desulfurisation, magnesium alloys, titanium extraction (Kroll process) and medical imaging (BaSO₄ barium meals). This lesson anchors Required Practical 10 — investigating the reactions of Group 2 elements and the relative solubilities of their compounds.

Spec mapping (AQA 7405): This lesson maps to §3.2.2 (Group 2, the alkaline earth metals). It builds on L0 (periodicity and Period 3 trends) and contrasts with L2 (Group 7 oxidising-ability trend, which runs opposite in direction). The quantitative rationale for the sulfate-solubility trend forward-links to §3.1.4 (energetics — lattice and hydration enthalpies in Born–Haber and dissolution cycles). The practical anchor is Required Practical 10. Refer to the official AQA specification document for the exact wording of each subsection.

Assessment objectives: AO1 recalls general formula [Ar/Kr/Xe]ns², electron configurations of Be–Ba, and the directions of the reactivity, hydroxide-solubility and sulfate-solubility trends. AO2 dominates Paper 1: predicting and balancing reactions of Group 2 metals with water, oxygen and dilute acid, and explaining the solubility trends in terms of lattice vs hydration enthalpy. AO3 — the A⁎ discriminator — asks students to rationalise why lattice enthalpy stays roughly constant for sulfates while hydration enthalpy falls sharply with cation size, to evaluate industrial uses (e.g. why BaSO₄ is safe to swallow despite Ba²⁺ toxicity), and to integrate Born–Haber thinking with periodic-trend prediction.

General Formula and Electron Configurations

Every Group 2 element has the outer configuration ns² and forms a dipositive M²⁺ ion with a noble-gas core:

Element	Z	Full configuration	Short form	Ion (M²⁺)
Be	4	1s² 2s²	[He] 2s²	[He]
Mg	12	1s² 2s² 2p⁶ 3s²	[Ne] 3s²	[Ne]
Ca	20	1s² 2s² 2p⁶ 3s² 3p⁶ 4s²	[Ar] 4s²	[Ar]
Sr	38	[Ar] 3d¹⁰ 4s² 4p⁶ 5s²	[Kr] 5s²	[Kr]
Ba	56	[Kr] 4d¹⁰ 5s² 5p⁶ 6s²	[Xe] 6s²	[Xe]
Ra	88	[Rn] 7s²	[Rn] 7s²	[Rn]

Loss of both s electrons generates a closed-shell cation, which is why the +2 oxidation state dominates Group 2 chemistry and Be⁺ or Ca⁺ species are never observed in stable compounds.

Physical Properties and Trends

Property	Trend down Group 2	Underlying reason
Atomic radius	Increases	Each successive element occupies a higher principal-quantum-number shell
First ionisation energy	Decreases	Outer electrons further from nucleus and better shielded; this outweighs the rise in nuclear charge
Second ionisation energy	Decreases (parallels IE₁)	Same shielding-and-distance argument applies
Electronegativity	Decreases	Bonding pair held less tightly as atomic radius grows
Melting point	Generally decreases (Mg → Sr; Ba slightly higher than Sr)	Metallic-bond strength weakens as cation–delocalised-electron distance increases
Density	Generally increases	Atomic mass grows faster than atomic volume

The melting-point pattern is not perfectly monotonic — beryllium melts unusually high (1287 °C) because of its tiny radius, and magnesium and strontium adopt different crystal packings — but the broad downward drift from Mg to Ba is what AQA examines.

Reactivity Trend Down the Group

Reactivity increases going down Group 2.

The first and second ionisation energies both fall down the group: IE₁ + IE₂ for Mg is 2189 kJ mol⁻¹, for Ca 1735, for Sr 1614, for Ba 1467.
Although nuclear charge rises (Mg: +12, Ba: +56), the outer electrons sit in successively larger shells with more shielding by full inner shells; the effective nuclear charge experienced by the ns² electrons changes only modestly.
The net result: it becomes progressively easier to remove the two outer s electrons, so the metals are more readily oxidised, M(s) → M²⁺(aq) + 2e⁻, going down the group.
Beryllium is anomalous: its very small radius and high charge density confer a passivating oxide layer that resists attack by water and dilute acid, so Be is functionally inert in most of the reactions discussed below.

Key Point: "Reactivity increases down the group" applies to oxidation by water, oxygen and acid. It is the opposite of Group 7, where the most reactive halogen sits at the top (F₂) because halogens gain electrons and the smallest atom does so most readily.

Reactions with Water

Group 2 metals (excluding Be, which does not react under standard conditions) reduce water to hydrogen, generating the metal hydroxide:

General equation: M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)

Specific Reactions

Beryllium: No reaction with water at any temperature under standard A-Level conditions. The thin BeO/Be(OH)₂ layer passivates the surface.

Magnesium (cold water): Very slow; only a faint stream of bubbles after several days. The solution becomes weakly alkaline as small amounts of Mg(OH)₂ dissolve.

Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g)

Magnesium (steam): Vigorous; a brilliant white flame is observed and the oxide, not the hydroxide, is the solid product because steam temperatures (≥ 100 °C) decompose any Mg(OH)₂ formed back to MgO.

Mg(s) + H₂O(g) → MgO(s) + H₂(g)

Calcium: Steady reaction with cold water. Bubbles of hydrogen rise smoothly; the solution turns cloudy because Ca(OH)₂ is only sparingly soluble (limewater); the solid Ca shrinks and may float on the gas. pH of the saturated solution ≈ 12.4.

Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

Strontium: More vigorous than calcium — rapid effervescence, the metal may glow faintly, and Sr(OH)₂ is appreciably soluble so the solution stays clearer than for Ca.

Sr(s) + 2H₂O(l) → Sr(OH)₂(aq) + H₂(g)

Barium: Vigorous; bubbles stream off immediately and the strongly alkaline Ba(OH)₂ solution can reach pH ≈ 13–14.

Ba(s) + 2H₂O(l) → Ba(OH)₂(aq) + H₂(g)

Exam Tip: The trend in vigour with cold water is Mg (very slow) → Ca (steady) → Sr (vigorous) → Ba (very vigorous), tracking the fall in IE₁ + IE₂. To test the gas, use a lit splint (squeaky pop → H₂). To confirm a CO₂ evolution from carbonate decomposition, use limewater (turns milky) — different gas, different test.

Reactions with Oxygen

All Group 2 metals burn in oxygen to give white ionic oxides:

General equation: 2M(s) + O₂(g) → 2MO(s)

Metal	Equation	Flame colour	Solid product
Mg	2Mg + O₂ → 2MgO	Brilliant white	White MgO
Ca	2Ca + O₂ → 2CaO	Brick-red / orange-red	White CaO
Sr	2Sr + O₂ → 2SrO	Crimson	White SrO
Ba	2Ba + O₂ → 2BaO	Pale green / apple-green	White BaO

The characteristic flame colours arise from outer-electron excitation followed by photon emission as electrons fall back to lower energy levels — the same principle as flame tests in qualitative analysis.

Peroxides and Superoxides

At elevated temperatures or under excess oxygen, the heavier Group 2 metals form higher-oxygen species — the AQA syllabus expects awareness rather than full mechanism:

Strontium and barium can form peroxides (containing O₂²⁻): for example BaO + ½O₂ → BaO₂.
Barium in excess oxygen can form a small amount of superoxide Ba(O₂)₂ at very high temperatures; this is much less stable than the alkali-metal analogues (KO₂, RbO₂).
Mg and Ca form only the simple oxide under normal burning conditions; their cations are too polarising for the larger, more diffuse O₂²⁻ and O₂⁻ anions to be stable.

The Group 2 oxides are all basic and react with water to form alkaline hydroxides:

MO(s) + H₂O(l) → M(OH)₂(aq)

CaO + H₂O → Ca(OH)₂ is a strongly exothermic reaction (the basis of "slaking" lime industrially).

Reactions with Dilute Acid

Group 2 metals reduce dilute mineral acids to hydrogen:

General equation: M(s) + 2HCl(aq) → MCl₂(aq) + H₂(g)

Observation: Bubbles of colourless gas effervesce vigorously, the metal disappears progressively, and the solution warms (the reaction is exothermic). Confirm the gas with a lit splint — a "squeaky pop" indicates hydrogen. (A glowing splint that relights would indicate O₂; limewater that turns milky would indicate CO₂ — different tests for different gases, and worth distinguishing in the exam.)

Specific examples:

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Ca(s) + 2HCl(aq) → CaCl₂(aq) + H₂(g)
Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g)

With dilute sulfuric acid, magnesium reacts smoothly but calcium reaction quickly slows because CaSO₄ is only sparingly soluble — a CaSO₄ coating forms over the metal surface and passivates it. This is a useful synoptic example of how solubility (the next section) modifies kinetics.

Common Misconception: Concentrated HNO₃ oxidises most Group 2 metals to give NO or NO₂ rather than H₂. This is beyond AQA's required-equation list but is a reasonable AO3 extension: at A⁎ level, recognise that "dilute acid → H₂" is the rule and concentrated oxidising acids behave differently.

Solubility of Group 2 Hydroxides — Increases Down the Group

Hydroxide	Solubility (mol dm⁻³, 25 °C)	Description	pH of saturated solution
Mg(OH)₂	≈ 2 × 10⁻⁴	Insoluble	≈ 10
Ca(OH)₂	≈ 1.5 × 10⁻²	Sparingly soluble ("limewater")	≈ 12.4
Sr(OH)₂	≈ 6.6 × 10⁻²	Moderately soluble	≈ 13
Ba(OH)₂	≈ 0.24	Soluble	≈ 13–14

Industrial / everyday uses tied to solubility

Mg(OH)₂ — "milk of magnesia": Used as an antacid and mild laxative. Its very low solubility means the suspension neutralises stomach HCl gently without raising blood pH systemically.
Ca(OH)₂ — "limewater": The traditional CO₂ test (Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l); milky precipitate). Also dressed onto acidic agricultural soils to raise pH ("liming"), and added to drinking water to adjust pH and reduce pipe corrosion.
Ca(OH)₂ — flue-gas desulfurisation: Power-station flue gas is sprayed with a Ca(OH)₂ slurry to remove SO₂: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O (further oxidised to gypsum CaSO₄·2H₂O, sold for plasterboard).
Ba(OH)₂: Strong soluble base used in analytical titrations of weak acids.

Solubility of Group 2 Sulfates — Decreases Down the Group

Sulfate	Solubility (g per 100 g H₂O, 20 °C)	Description
MgSO₄	≈ 35	Soluble ("Epsom salt", MgSO₄·7H₂O)
CaSO₄	≈ 0.24	Sparingly soluble (gypsum, plaster of Paris)
SrSO₄	≈ 0.011	Almost insoluble
BaSO₄	≈ 0.00024	Insoluble

The qualitative test for sulfate ions

Acidified BaCl₂(aq) (the acid is dilute HCl, to remove any carbonate that would also precipitate) gives a thick white precipitate with any soluble sulfate:

Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

Why is BaSO₄ safe to swallow despite Ba²⁺ being toxic?

BaSO₄'s solubility in water and stomach acid is so low (≈ 2 × 10⁻³ g L⁻¹; Ksp ≈ 1 × 10⁻¹⁰) that the equilibrium Ba²⁺ concentration in body fluids is negligible — far below the toxic threshold (≈ 20 mg Ba²⁺ causes acute poisoning). Patients swallow a thick BaSO₄ suspension ("barium meal" or "barium swallow") before X-ray imaging; barium's high atomic number gives strong X-ray absorption, opacifying the gut for diagnosis. Soluble barium salts (BaCl₂, Ba(NO₃)₂) are highly toxic because they liberate free Ba²⁺, which blocks K⁺ channels in cardiac muscle.

Quantitative Rationale: Lattice Enthalpy vs Hydration Enthalpy

Whether an ionic salt MX dissolves is decided by the sign and magnitude of the enthalpy of solution ΔH_sol:

ΔH_sol = −ΔH_LE(MX) + ΔH_hyd(M²⁺) + ΔH_hyd(Xⁿ⁻)

(ΔH_LE here written as a positive number for the lattice-dissociation enthalpy.)

Both ΔH_LE and ΔH_hyd depend on ion radius and charge, but the dependence is different for hydroxide and sulfate salts:

Hydroxides — increasing solubility down the group

The anion OH⁻ is small. Lattice enthalpy is therefore sensitive to both ion sizes; as the cation grows (Mg²⁺ → Ba²⁺), 1/(r₊ + r₋) falls noticeably and ΔH_LE becomes less exothermic.
ΔH_hyd of the cation also falls (Mg²⁺ is strongly hydrated, Ba²⁺ weakly), but more slowly than ΔH_LE because hydration depends on 1/r₊ while lattice depends on 1/(r₊ + r₋) where r₋ is fixed and small.
Net: ΔH_sol becomes more exothermic going down the group → hydroxide solubility increases.

Sulfates — decreasing solubility down the group

The anion SO₄²⁻ is large. Lattice enthalpy is dominated by the large anion radius and is almost insensitive to which Group 2 cation is paired with it: ΔH_LE for MgSO₄, CaSO₄, SrSO₄ and BaSO₄ differ by only a few percent.
ΔH_hyd of the cation still falls sharply with cation size (Mg²⁺ is small and highly charge-dense; Ba²⁺ is large and weakly hydrated).
Net: ΔH_LE is roughly constant while the (exothermic) hydration contribution shrinks → ΔH_sol becomes less exothermic (or more endothermic) going down → sulfate solubility decreases.

This argument — "small fixed anion vs large fixed anion" — is exactly the AO3 explanation expected at A⁎. It also forward-links to L4 of the energetics course (§3.1.4), where Born–Haber cycles for Group 2 halides quantify these enthalpies explicitly.

Exam Tip: Mnemonic — "Hydroxides Hate Smallness; Sulfates Shrink." Or remember that the larger of the two ions controls which trend wins: when the anion is small (OH⁻), cation size matters most for the lattice; when the anion is large (SO₄²⁻), cation size matters most for the hydration.

Thermal Decomposition of Carbonates and Nitrates

Group 2 carbonates and nitrates decompose on strong heating; the temperature required increases down the group, because the polarising power of the cation falls.

Carbonates: MCO₃(s) → MO(s) + CO₂(g)

Carbonate	Decomposition T / °C
MgCO₃	≈ 540
CaCO₃	≈ 840
SrCO₃	≈ 1290
BaCO₃	≈ 1360

Nitrates: 2M(NO₃)₂(s) → 2MO(s) + 4NO₂(g) + O₂(g) — brown NO₂ fumes are diagnostic.

The smaller, more highly charge-dense Mg²⁺ polarises the carbonate or nitrate anion more, distorting its electron cloud and weakening a C–O or N–O bond; decomposition therefore occurs at a lower temperature. Larger Ba²⁺ polarises the anion less, so a higher temperature is needed.

Industrial Uses Anchored in the Trends

Magnesium alloys (Al–Mg, Mg–Zn): Light-weight, high strength-to-mass ratio; used in aerospace, automotive wheels, laptop chassis. Pure Mg also burns brilliantly white in flares and pyrotechnics.
Titanium extraction (Kroll process): TiCl₄(l) + 2Mg(l) → Ti(s) + 2MgCl₂(l), at ≈ 850 °C under an argon atmosphere. Magnesium is preferred because it is reactive enough to reduce Ti⁴⁺ but its chloride MgCl₂ is easily electrolysed back to Mg for recycling. This is a flagship A-Level industrial example of Group 2 reactivity used for economic benefit.
CaO and Ca(OH)₂ — steel-making: CaO ("quicklime") is added to a blast-furnace charge to remove silicate impurities as a CaSiO₃ slag. Ca(OH)₂ is used to neutralise acidic mine drainage and acidified soil.
CaCO₃ — construction: Limestone is the raw material for cement (heated with clay) and a primary aggregate for concrete.
Ca(OH)₂ — flue-gas desulfurisation: Ca(OH)₂ + SO₂ → CaSO₃ + H₂O; further oxidation gives marketable gypsum.
BaSO₄ — medical radiography: Insoluble, high-Z contrast agent for upper-GI imaging.
Ca²⁺ in biology: Bone (hydroxyapatite, Ca₅(PO₄)₃OH), enzyme cofactor, neurotransmitter trigger — synoptic link to A-Level Biology.

RP10 Anchor — Investigating Group 2 Reactions

Required Practical 10 in AQA 7405 examines the relative reactivity of Group 2 metals and the relative solubilities of their compounds. Outline:

Group 2: The Alkaline Earth Metals