Lesson 8: Group 2 — Alkaline Earth Metals

Group 2, the alkaline earth metals — beryllium to barium — show some of the cleanest trends in the periodic table, making them ideal for learning how electronic structure explains chemistry. Their compounds are everywhere: limestone buildings, antacid tablets, the plaster in hospital casts and the contrast agent in X-ray imaging. This lesson covers trends in atomic radius, ionisation energy and reactivity, the reactions with oxygen and water, solubility trends of hydroxides and sulphates, thermal decomposition of carbonates and nitrates, and the everyday uses of Group 2 compounds.

Trends in properties down Group 2

Atomic radius increases down the group: each successive element adds a new electron shell, so the outer electrons are further from the nucleus despite the increasing nuclear charge. First ionisation energy decreases: the outer electron is further away and more shielded by inner shells, outweighing the extra protons. Explain every trend with all three factors — nuclear charge, distance, shielding; examiners penalise answers that mention only one.

Melting points generally decrease down the group (Mg is anomalously low because of its crystal structure — a detail worth knowing for "explain the anomaly" questions). Reactivity increases down the group: the metals lose their two outer electrons more easily as ionisation energies fall. This mirrors Group 1, but Group 2 reactions are less vigorous because removing two electrons costs more energy than removing one.

Beryllium is the exception that proves the rules: its small, highly charged Be²⁺ ion polarises anions so strongly that its compounds have significant covalent character (Fajans' rules), and beryllium itself is toxic. A-level questions focus on Mg to Ba.

Reactions with oxygen and water

Group 2 metals burn in oxygen to form white solid oxides: 2Mg(s) + O₂(g) → 2MgO(s), with the characteristic brilliant white flame of magnesium. Barium also forms the peroxide BaO₂. The oxides are basic: MgO reacts with dilute acids to form salts, and the more soluble oxides and hydroxides give alkaline solutions — hence "alkaline earth metals".

With water, the metals form hydroxides and hydrogen: Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g). Reactivity increases down the group: magnesium reacts only very slowly with cold water (but vigorously with steam), calcium reacts steadily with cold water, and barium reacts rapidly. Strontium and barium hydroxides are soluble enough to give distinctly alkaline solutions.

The trend is explained by ionisation energies: forming M²⁺ requires removing two electrons, which becomes easier down the group. Practical work may involve measuring the pH of Group 2 hydroxide solutions: pH increases down the group as hydroxide solubility increases, giving higher OH⁻ concentrations. Link observation (pH) to trend (solubility) to cause (ionisation energy and lattice effects).

Solubility trends — hydroxides and sulphates

Solubility of Group 2 hydroxides increases down the group: Mg(OH)₂ is sparingly soluble, Ba(OH)₂ much more soluble. Adding NaOH to solutions of Mg²⁺, Ca²⁺, Sr²⁺ and Ba²⁺ gives white hydroxide precipitates, with magnesium producing the most precipitate. Conversely, solubility of Group 2 sulphates decreases down the group: BaSO₄ is famously insoluble while MgSO₄ (Epsom salts) dissolves readily.

The sulphate trend gives the classic test for sulphate ions: add dilute HCl (to remove carbonate interference), then BaCl₂ solution — a white precipitate of BaSO₄ confirms sulphate: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s). The explanation lies in the balance of lattice and hydration enthalpies (see the worked examples below). You need the trend and the test; the enthalpy explanation earns the extra marks.

Barium sulphate's insolubility makes it safe to swallow as a barium meal for X-ray imaging of the gut — even though soluble barium compounds are highly toxic, BaSO₄ passes through unabsorbed. This is a favourite "explain using solubility" exam question linking chemistry to medicine.

Thermal decomposition of carbonates and nitrates

Group 2 carbonates decompose on heating: MCO₃(s) → MO(s) + CO₂(g). Thermal stability increases down the group — magnesium carbonate decomposes readily while barium carbonate needs much stronger heating. Explanation: the small, highly charged Mg²⁺ ion polarises the carbonate ion, distorting it toward CO₂ and O²⁻ and weakening it; larger cations polarise less, so the carbonate survives higher temperatures. This is Fajans' polarisation applied to stability.

Group 2 nitrates decompose similarly: 2M(NO₃)₂(s) → 2MO(s) + 4NO₂(g) + O₂(g), giving brown NO₂ gas. Stability again increases down the group for the same polarisation reason. Practical questions may ask you to heat carbonates and test for CO₂ with limewater (which turns milky), comparing decomposition temperatures.

Limestone (CaCO₃) decomposition is industrially vital: heating in lime kilns produces quicklime, CaO, for cement and steelmaking. The CO₂ released is a major industrial emission — an environmental point examiners like to attach to the chemistry. The stability trend also explains why Group 1 carbonates (except Li₂CO₃) do not decompose: their singly charged cations polarise too weakly.

Uses of Group 2 compounds

Magnesium is used in the extraction of titanium (the Kroll process: TiCl₄ + 2Mg → Ti + 2MgCl₂) because it is a strong enough reducing agent to displace titanium, and in lightweight alloys for aircraft and cars. Magnesium hydroxide (milk of magnesia) and calcium carbonate are antacids: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O.

Calcium hydroxide (slaked lime) is spread on fields to neutralise acidic soil and is used in water treatment. Calcium oxide (quicklime) makes cement and mortar. Strontium compounds give the red colour in fireworks and flares. Every use traces to a property: basicity for antacids and agriculture, reducing power for titanium extraction, insolubility for barium meals, flame colour for fireworks.

Exam questions often ask you to "suggest a use" from properties, or vice versa. Structure your answer as property → explanation → use. Example: "Barium sulphate is insoluble (property) because the Ba²⁺ and SO₄²⁻ ions form a very stable lattice (explanation), so it can be swallowed safely as an X-ray contrast medium without releasing toxic Ba²⁺ ions (use)."

Data treatment — thermal decomposition and solubility tests

Worked example: comparing carbonate stability quantitatively. A student heats 2.10 g of MgCO₃ (Mᵣ 84.3) strongly in a fume cupboard and collects the gas over water. Moles of MgCO₃ = 2.10 / 84.3 = 0.0249 mol, so complete decomposition should release 0.0249 × 24.0 = 0.598 dm³ of CO₂ at RTP. Repeating with 2.10 g of CaCO₃ (Mᵣ 100.1) gives only 0.0210 mol — and, more importantly, CaCO₃ needs a much hotter flame to decompose at all. Reporting both the gas volumes and the heating required turns a qualitative trend into evidence: the smaller Mg²⁺ polarises carbonate more strongly, so less thermal energy is needed to break it down.

Practical treatment: the hydroxide-precipitation trend. Add NaOH(aq) dropwise to separate tubes of 0.1 mol dm⁻³ MgCl₂, CaCl₂, SrCl₂ and BaCl₂. All give white hydroxide precipitates, but their amounts fall visibly down the group: Mg(OH)₂ dense and persistent, Ba(OH)₂ barely visible. Test the pH of each supernatant: it rises down the group, because more hydroxide dissolves → more OH⁻(aq). The two observations (less precipitate, higher pH) are the experimental signature of "hydroxide solubility increases down Group 2" — quote both in exam answers.

Going deeper — charge density ties the whole group together

One quantity, charge density (charge divided by ionic size, roughly), organises almost every Group 2 trend. It falls down the group as Mg²⁺ → Ba²⁺ grow larger at constant +2 charge. Falling charge density means: weaker polarisation of anions (carbonates and nitrates become thermally more stable down the group); less exothermic hydration enthalpies (the ion attracts water dipoles less fiercely); and lattice enthalpies that also become less exothermic, but by different amounts for different anions — which is why the two solubility trends run in opposite directions. For hydroxides, the small OH⁻ ion's lattice enthalpy falls steeply as the cation grows, while hydration falls less steeply overall, so dissolving wins down the group. For sulphates, the huge SO₄²⁻ ion keeps lattice enthalpies almost constant (the anion dominates the ion separation) while hydration falls, so dissolving loses down the group. This "which term changes faster" argument is the rigorous version of the trend and converts recall into reasoning.

The same idea predicts unfamiliar cases: Group 2 fluorides are sparingly soluble (small F⁻, strong lattices), and the trend of carbonate stability continues into Group 1 — Li₂CO₃ decomposes on heating like a Group 2 carbonate because Li⁺ is small with high charge density, while Na₂CO₃ and K₂CO₃ do not. Diagonal relationships (Li resembling Mg) are examiners' favourite way to test whether you understand charge density rather than merely the group's facts.

The lime cycle packages several of these ideas into one industrial story. Limestone, CaCO₃, is heated strongly to give quicklime, CaO, plus CO₂ (thermal decomposition — feasible above about 1110 K, Lesson 4). Quicklime reacts vigorously and exothermically with water to give slaked lime, Ca(OH)₂. Slaked lime solution (limewater) is the CO₂ test reagent: bubbled CO₂ first turns it milky (a precipitate of CaCO₃) and, in excess, the precipitate redissolves as soluble calcium hydrogencarbonate forms. The cycle closes because the same CaCO₃ re-forms. Group 2 flame tests complete the identification toolkit: calcium brick-red, strontium crimson red, barium apple-green (magnesium has no visible flame colour — its emission lies largely outside the visible range). A question that gives a white solid, a flame colour, effervescence with acid and a milky limewater test is asking you to walk this cycle step by step.

Common mistakes

• Reversing the solubility trends. Hydroxide solubility increases down Group 2; sulphate solubility decreases. Learn them as a pair and check which compound the question names.
• Explaining thermal stability by "bond strength" alone. The mark is for polarisation: the small, highly charged cation distorts the carbonate/nitrate ion; larger cations polarise less, so stability increases down the group.
• Omitting the HCl in the sulphate test. Acid is added before BaCl₂ to destroy carbonate, which would otherwise give a white BaCO₃ precipitate and a false positive.
• Saying Mg "does not react with water". It reacts very slowly with cold water and vigorously with steam — the trend is graded, not on/off.

Exam Experience

Q1. (a) Write an equation for the reaction of barium with water, and state two observations. [3] (b) A student adds barium chloride solution to a solution containing sulphate ions. Write the ionic equation for the reaction and state what is observed. [2] (c) Explain why magnesium carbonate decomposes at a much lower temperature than calcium carbonate. [3]

Mark scheme. Q1(a): Ba(s) + 2H₂O(l) → Ba(OH)₂(aq) + H₂(g) [1]; observations: effervescence/fizzing [1] and the barium sinks/dissolves (the solution becomes alkaline) [1]. (b) Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) [1]; a white precipitate forms [1]. (c) Mg²⁺ is smaller than Ca²⁺ with the same charge, so it has a higher charge density [1]; it polarises (distorts) the carbonate ion's electron cloud more strongly [1]; the distorted carbonate breaks down to CO₂ and O²⁻ more easily, so decomposition occurs at a lower temperature [1].