Group 2 — trends in reactivity

Down Group 2 (Be → Ba) the atomic radius and shielding increase, so the first and second ionisation energies fall. The metal loses its two outer electrons more easily, so reactivity increases down the group — each metal is a stronger reducing agent.

• With oxygen: 2Mg(s) + O₂(g) → 2MgO(s) — a brilliant white flame.
• With chlorine: Mg(s) + Cl₂(g) → MgCl₂(s).
• With water: Mg reacts only very slowly with cold water, but Ca → Ba react readily: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g). Effervescence gets more vigorous down the group.
• Mg with steam is different: Mg(s) + H₂O(g) → MgO(s) + H₂(g).
• Oxides with water give alkaline hydroxides: CaO(s) + H₂O(l) → Ca(OH)₂(aq). Oxides with dilute acid give a salt and water: MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l).

Solubility and thermal stability trends

Two trends that run in opposite directions — learn them separately:

• Hydroxides become MORE soluble down Group 2. Mg(OH)₂ is almost insoluble (milk of magnesia); Ba(OH)₂ dissolves well and gives a strongly alkaline solution. So the pH of the resulting solution rises down the group.
• Sulfates become LESS soluble down Group 2. BaSO₄ is essentially insoluble — that is exactly why acidified BaCl₂ is the test for sulfate ions, and why the (toxic) Ba²⁺ ion can be swallowed safely as a "barium meal" for X-rays.

Thermal stability of the nitrates and carbonates increases down the group. A small, highly charged cation has a large polarising power: it distorts (polarises) the large anion's electron cloud, weakening a bond within the anion, so the compound decomposes at a lower temperature. Down the group the cation gets bigger, so polarising power falls and the compound becomes more stable.

Group 1 exception: lithium is the odd one out — Li₂CO₃ and LiNO₃ decompose like Group 2 compounds (Li⁺ is small and polarising). Other Group 1 nitrates give only the nitrite and oxygen: 2NaNO₃ → 2NaNO₂ + O₂.

Thermal stability

Your turn — decomposing a nitrate

Flame tests

Dip a clean nichrome or platinum wire in concentrated HCl, then into the solid, then hold it in a hot Bunsen flame.

• Heat energy excites an electron to a higher energy level. When the electron drops back down, the energy is emitted as light of a definite frequency — the colour is characteristic of the element because its energy levels are unique.
• Li red · Na yellow · K lilac · Ca brick-red · Sr crimson · Ba pale green.
• Mg²⁺ gives no colour: its electrons are held so tightly that the energy of the flame is not enough to excite them (and any emission would be in the UV).

Practical detail: the HCl converts the compound to the more volatile chloride, and a contaminated wire (sodium is everywhere) will drown the true colour in yellow.

Group 7 — physical trends and oxidising power

Down Group 7 (F₂ → I₂) the molecules contain more electrons, so the London forces between them get stronger and the melting and boiling temperatures rise: F₂ pale yellow gas, Cl₂ green gas, Br₂ red-brown liquid, I₂ grey-black solid that sublimes to a purple vapour.

Chemically the halogens are oxidising agents — they gain one electron each: X₂ + 2e⁻ → 2X⁻. Down the group the atomic radius and shielding increase, so the incoming electron is attracted less strongly: oxidising power decreases down the group.

• Cl₂ displaces Br⁻ and I⁻ · Br₂ displaces only I⁻ · I₂ displaces neither.
• Halide ions are reducing agents, and reducing power increases down the group — I⁻ is the best reducing agent (the outer electron is furthest from the nucleus and most shielded).

Trend in oxidising power

Halides with concentrated sulfuric acid

This reaction is a superb test of reducing power. Concentrated H₂SO₄ is added to the solid potassium halide.

• KF, KCl — no redox, just an acid–base reaction: KCl + H₂SO₄ → KHSO₄ + HCl (misty white fumes). Cl⁻ is too weak a reducing agent to reduce sulfur.
• KBr — some HBr fumes, plus Br⁻ reduces S from +6 to +4: 2Br⁻ + 2H₂SO₄ → Br₂ + SO₂ + 2H₂O + SO₄²⁻. Observations: brown fumes of Br₂ and a choking gas (SO₂).
• KI — I⁻ is the strongest reducing agent, taking sulfur all the way from +6 down to −2: 8I⁻ + 9H₂SO₄ → 4I₂ + H₂S + 4H₂O + 8HSO₄⁻. Observations: purple iodine vapour / black solid, and the rotten-egg smell of H₂S. SO₂ and S (a yellow solid) also form.

The pattern: the further down the group, the more the halide reduces the sulfur — H₂SO₄ (+6) → SO₂ (+4) → S (0) → H₂S (−2).

Your turn — oxidation number of sulfur

The qualitative tests you must know

Order matters: test for carbonate first (it fizzes with acid), then sulfate, then halide — because carbonate and sulfate both give white precipitates that would confuse the silver nitrate test.

• Carbonate: add dilute HCl → effervescence; the gas turns limewater milky. CO₃²⁻ + 2H⁺ → H₂O + CO₂
• Sulfate: add dilute HCl (to remove carbonate) then BaCl₂(aq) → white precipitate. Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)
• Halide: add dilute HNO₃ then AgNO₃(aq): Ag⁺(aq) + Cl⁻(aq) → AgCl(s) white · AgBr cream · AgI yellow. Confirm with ammonia: AgCl dissolves in dilute NH₃; AgBr dissolves only in concentrated NH₃; AgI is insoluble even in concentrated NH₃.
• Ammonium: warm with NaOH(aq) → a gas that turns damp red litmus blue. NH₄⁺ + OH⁻ → NH₃ + H₂O

Why acidify first?

Your turn — neutralising a Group 2 hydroxide

Which test result is it?

Tap an observation, then tap the ion it identifies.

Silver halides and ammonia

Your turn — chlorine in chlorate(V)

Predicting for astatine

Match the halogen reaction to its observation

Tap an item on the left, then its partner on the right.

The big ideas to know

You've now covered Topic 4: Inorganic Chemistry and the Periodic Table of the Edexcel A-level Chemistry (9CH0) specification. Press Finish to see your score.