What you'll learn
Group II (the alkaline earth metals) and Group VII (the halogens) show how chemical behaviour changes systematically down a group. In Group II, metals become more reactive down the group; in Group VII, the halogens become less reactive as oxidising agents. This guide covers the physical and chemical trends in Group II — reactions with water and oxygen, the solubility of hydroxides and sulfates, and the thermal stability of carbonates and nitrates — and in Group VII — physical properties, oxidising power, displacement reactions, reactions of halide ions with silver nitrate and concentrated sulfuric acid, the thermal stability of hydrogen halides, and the reactions of chlorine with water and sodium hydroxide. Uses include limestone and lime in Caribbean construction and agriculture, and chlorine in water treatment.
CAPE questions frequently require trends to be explained using atomic radius, shielding and ionic charge density, and ask for observations and equations for test-tube reactions. Learn both the patterns and the reasons.
Key terms and definitions
Alkaline earth metals — The Group II elements: Be, Mg, Ca, Sr, Ba.
Halogens — The Group VII elements: F, Cl, Br, I.
Thermal stability — Resistance to decomposition on heating.
Charge density — The charge of an ion divided by its size; small, highly charged ions have high charge density.
Polarisation — Distortion of an anion's electron cloud by a nearby cation.
Oxidising agent — A species that accepts electrons and is itself reduced.
Displacement reaction — A more reactive halogen displaces a less reactive halide from solution.
Disproportionation — A reaction in which the same element is simultaneously oxidised and reduced.
Core concepts
Group II: physical trends
Down the group:
- atomic radius increases (more shells);
- first and second ionisation energies decrease (outer electrons further from the nucleus and more shielded);
- reactivity increases, because the two outer electrons are lost more easily to form M²⁺ ions;
- melting points generally decrease (with irregularities), because metallic bonding weakens as ions get larger.
Group II: reactions
With oxygen: all burn to form oxides, MO: 2Mg + O₂ → 2MgO. Barium can also form a peroxide.
With water:
- Magnesium reacts very slowly with cold water, but burns in steam: Mg + H₂O(g) → MgO + H₂.
- Calcium, strontium and barium react with cold water with increasing vigour: Ca + 2H₂O → Ca(OH)₂ + H₂.
Solubility trends
- Hydroxides become more soluble down the group. Mg(OH)₂ is sparingly soluble (used as "milk of magnesia" to treat indigestion); Ba(OH)₂ is quite soluble, giving a strongly alkaline solution.
- Sulfates become less soluble down the group. MgSO₄ is soluble; BaSO₄ is very insoluble — used in "barium meals" for X-ray imaging, and as a test for sulfate ions (a white precipitate with barium chloride).
Thermal decomposition of carbonates and nitrates
Carbonates decompose on heating to form the oxide and carbon dioxide: MCO₃ → MO + CO₂. Nitrates decompose to form the oxide, nitrogen dioxide and oxygen: 2M(NO₃)₂ → 2MO + 4NO₂ + O₂.
Both become more thermally stable down the group — higher temperatures are needed. Explanation: smaller cations such as Mg²⁺ have a higher charge density, so they polarise the large carbonate or nitrate ion more strongly, distorting its electron cloud and weakening its bonds, making decomposition easier. Larger cations further down the group polarise less, so the anions are more stable.
Caribbean application: limestone (calcium carbonate) is heated in kilns to make quicklime (CaO), which is then slaked with water to make slaked lime (Ca(OH)₂). Lime is used in cement and mortar, to neutralise acidic soils for agriculture, and in water treatment and sugar refining.
Group VII: physical trends
| Halogen | State at room temperature | Colour |
|---|---|---|
| F₂ | gas | pale yellow |
| Cl₂ | gas | yellow-green |
| Br₂ | liquid | red-brown |
| I₂ | solid | grey-black (purple vapour) |
Boiling points increase down the group because molecules have more electrons, so London forces become stronger. Electronegativity decreases down the group.
Group VII: oxidising power
Halogens act as oxidising agents by gaining electrons: X₂ + 2e⁻ → 2X⁻. Oxidising power decreases down the group because the atoms get larger and the incoming electron is further from the nucleus and more shielded, so it is attracted less strongly.
Displacement reactions demonstrate this:
- Cl₂ + 2KBr → 2KCl + Br₂ (solution turns orange)
- Cl₂ + 2KI → 2KCl + I₂ (solution turns brown)
- Br₂ + 2KI → 2KBr + I₂
- I₂ does not displace Cl⁻ or Br⁻.
Adding a non-polar solvent such as cyclohexane makes colours clearer: bromine is orange, iodine is purple in the organic layer.
Reactions of halide ions
With silver nitrate (acidified with dilute nitric acid):
- Cl⁻ → white precipitate of AgCl, soluble in dilute ammonia;
- Br⁻ → cream precipitate of AgBr, soluble in concentrated ammonia;
- I⁻ → yellow precipitate of AgI, insoluble in concentrated ammonia.
Ag⁺(aq) + X⁻(aq) → AgX(s)
With concentrated sulfuric acid: reducing power of halide ions increases down the group.
- NaCl + H₂SO₄ → NaHSO₄ + HCl (steamy fumes; no redox).
- NaBr gives HBr, but some Br⁻ is oxidised to Br₂ (orange fumes) while sulfuric acid is reduced to SO₂.
- NaI gives HI, but I⁻ is readily oxidised to I₂ (purple fumes / black solid), and H₂SO₄ is reduced further to S (yellow) and H₂S (rotten egg smell).
Iodide is the strongest reducing agent because its outer electrons are furthest from the nucleus and most easily lost.
Thermal stability of hydrogen halides
Stability decreases from HF to HI because the H–X bond becomes longer and weaker as the halogen atom gets larger. A hot wire placed into HI causes purple iodine vapour to form, whereas HCl does not decompose.
Reactions of chlorine
With water: Cl₂ + H₂O ⇌ HCl + HClO. Chloric(I) acid (HClO) kills bacteria, which is why chlorine is used to treat drinking water and swimming pools across the Caribbean. This is a disproportionation: chlorine (0) is reduced to −1 in HCl and oxidised to +1 in HClO.
With cold dilute sodium hydroxide: Cl₂ + 2NaOH → NaCl + NaClO + H₂O. Sodium chlorate(I) is the active ingredient in household bleach.
With hot concentrated sodium hydroxide: 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O (chlorine is oxidised to +5).
Worked examples
Example 1: Thermal stability
Explain why calcium carbonate decomposes at a lower temperature than barium carbonate. (3 marks)
The Ca²⁺ ion is smaller than Ba²⁺, so it has a higher charge density (1). It polarises the carbonate ion more, distorting its electron cloud and weakening the C–O bonds (1). Less energy is therefore needed to decompose CaCO₃ (1).
Example 2: Displacement
Predict and explain what happens when chlorine water is added to potassium iodide solution.
The solution turns brown because iodine is formed: Cl₂ + 2I⁻ → 2Cl⁻ + I₂. Chlorine is a stronger oxidising agent than iodine, so it removes electrons from iodide ions.
Example 3: Identifying a halide
An unknown sodium halide gives a cream precipitate with acidified silver nitrate that dissolves in concentrated but not dilute ammonia. Identify the halide.
The halide is bromide.
Example 4: Disproportionation
Use oxidation numbers to show that the reaction of chlorine with cold NaOH is a disproportionation.
In Cl₂, chlorine is 0. In NaCl it is −1 (reduced); in NaClO it is +1 (oxidised). The same element is both oxidised and reduced.
Example 5: Volume of gas from a carbonate
2.50 g of calcium carbonate is heated until it has completely decomposed. Calculate the volume of carbon dioxide produced at room temperature and pressure. (Mᵣ CaCO₃ = 100; molar gas volume = 24 dm³ at r.t.p.)
CaCO₃ → CaO + CO₂. Moles CaCO₃ = 2.50 ÷ 100 = 0.0250 mol = moles CO₂. Volume = 0.0250 × 24 = 0.600 dm³ (600 cm³).
Example 6: Concentrated sulfuric acid and sodium bromide
Write equations for the reactions of sodium bromide with concentrated sulfuric acid and identify the redox step.
Acid–base step: NaBr + H₂SO₄ → NaHSO₄ + HBr (steamy fumes). Redox step: 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O (orange-brown fumes of bromine).
In the redox step, bromine is oxidised from −1 to 0 and sulfur is reduced from +6 in H₂SO₄ to +4 in SO₂. Bromide ions are strong enough reducing agents to reduce sulfuric acid to SO₂, but not as far as sulfur or H₂S, which iodide achieves.
Example 7: Uses linked to properties
Explain why magnesium hydroxide is used in antacids and why barium sulfate is safe to swallow for X-rays, even though barium ions are toxic.
Magnesium hydroxide is a weak base that is only sparingly soluble, so it neutralises excess stomach acid gently without making the stomach strongly alkaline: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O. Barium sulfate is extremely insoluble, so almost no Ba²⁺ ions dissolve and enter the bloodstream; it absorbs X-rays and outlines the digestive tract.
Example 8: Predicting a reaction
Predict whether bromine will react with aqueous sodium chloride, and explain.
No reaction occurs. Bromine is a weaker oxidising agent than chlorine, so it cannot remove electrons from chloride ions to form chlorine.
Trends summary
| Property | Group II (down the group) | Group VII (down the group) |
|---|---|---|
| Atomic radius | increases | increases |
| Reactivity | increases (metals lose electrons more easily) | decreases (non-metals gain electrons less easily) |
| Ionisation / electronegativity | ionisation energy decreases | electronegativity decreases |
| Solubility | hydroxides ↑, sulfates ↓ | — |
| Thermal stability | carbonates and nitrates ↑ | hydrogen halides ↓ |
| Redox | stronger reducing agents | weaker oxidising agents; halide ions stronger reducing agents |
Seeing the two groups side by side highlights that reactivity trends depend on whether atoms lose or gain electrons.
Common mistakes and how to avoid them
- Mixing up the solubility trends — hydroxides become more soluble, sulfates less soluble, down Group II.
- Explaining thermal stability without polarisation.
- Saying halogens become more reactive down the group — their oxidising power decreases.
- Forgetting to acidify silver nitrate with nitric acid (to remove carbonate ions that would also give a precipitate).
- Wrong precipitate colours — white, cream, yellow for Cl⁻, Br⁻, I⁻.
- Forgetting that HClO is the bactericidal species in chlorinated water.
Self-test questions
- Which Group II hydroxide is the most soluble, and which sulfate is the least soluble? Answer: Barium hydroxide is the most soluble hydroxide; barium sulfate is the least soluble sulfate.
- Write the equation for the thermal decomposition of magnesium nitrate. Answer: 2Mg(NO₃)₂ → 2MgO + 4NO₂ + O₂.
- What colour is seen in the cyclohexane layer when bromine water is added to aqueous potassium iodide and shaken with cyclohexane? Answer: Purple, because iodine is displaced and dissolves in the organic layer.
- Explain why HI decomposes more readily than HCl when heated. Answer: The H–I bond is longer and weaker than the H–Cl bond because iodine is a larger atom, so less energy is needed to break it.
How this links to your Internal Assessment
Group II and VII chemistry offers many SBA opportunities: comparing the thermal decomposition of Group II carbonates by timing how long limewater takes to turn cloudy; investigating the solubility of hydroxides by titration; carrying out halogen displacement reactions; or identifying unknown halides using silver nitrate and ammonia. A Caribbean-relevant project could measure the available chlorine in local bleach brands by titration. Record observations precisely (colours, precipitates, gas tests), write equations, and evaluate safety and accuracy.
Exam technique for Group II and Group VII chemistry
- Explain trends with size, shielding and charge density, not just "reactivity".
- Give observations (colours, fumes, precipitates) as well as equations.
- Use oxidation numbers to explain redox and disproportionation.
- Learn the silver halide and concentrated sulfuric acid reactions as a set.
Quick revision summary
- Group II: reactivity increases down; hydroxides more soluble, sulfates less soluble; carbonates and nitrates more thermally stable down the group (less polarisation).
- Lime from limestone: CaCO₃ → CaO → Ca(OH)₂; uses in construction, agriculture and water treatment.
- Group VII: boiling points increase down; oxidising power decreases; displacement shows Cl₂ > Br₂ > I₂.
- AgX precipitates: white AgCl, cream AgBr, yellow AgI with distinct ammonia solubility.
- Reducing power of halides increases down the group (reactions with concentrated H₂SO₄).
- HX thermal stability decreases from HF to HI.
- Chlorine disproportionates with water and NaOH; HClO and NaClO disinfect and bleach.