amazon.co.ukPeriodicity and Group Chemistry
A* revision for periodic trends, Group 2 and Group 7 chemistry, halogen displacement and Period 3 oxides, with exam-style questions for AQA, Edexcel, OCR, WJEC and CCEA A-Level Chemistry.
🎯 Learning Objectives
- Explain trends in atomic radius, ionisation energy and melting point across Period 3, linking melting points to structure and bonding
- Describe Group 2 reactions with water, oxygen and acids, and explain the reactivity trend using ionisation energies
- Explain the opposing solubility trends of Group 2 hydroxides and sulfates and their uses (antacids, barium meals, lime cycle)
- Describe Group 7 trends, displacement reactions and the halide tests with AgNO₃ and NH₃
- Explain chlorine's disproportionation in water and NaOH, and evaluate water chlorination benefits and risks
- Classify Period 3 oxides as basic, acidic or amphoteric and write equations for their reactions with water and acids/bases
📋 Key Definitions and Core Concepts
Periodicity: The repeating pattern of physical and chemical properties of the elements across periods, arising from repeating outer electron configurations.
Displacement reaction: A more reactive halogen displaces the halide ion of a less reactive halogen from solution; the halogen is reduced and the halide oxidised.
Disproportionation: A redox reaction in which one element is simultaneously oxidised and reduced, e.g. Cl₂ → Cl⁻ + ClO⁻ in NaOH.
Amphoteric: A substance that reacts with both acids and bases, e.g. Al₂O₃.
Basic oxide: An oxide that reacts with water to form an alkaline solution or neutralises acids, e.g. Na₂O, MgO.
Acidic oxide: An oxide that reacts with water to form an acidic solution or reacts with bases, e.g. SO₂, P₄O₁₀, SiO₂.
Oxidising power: The tendency of a species to gain electrons (be reduced); decreases down Group 7 as atomic radius and shielding increase.
Disproportionation product: For chlorine in cold dilute NaOH: NaCl (−1) and NaClO (+1) — bleach.
Periodic Trends: Radius, Ionisation Energy and Melting Points
Periodicity is the repeating pattern of physical and chemical properties across periods. Atomic radius decreases across a period: nuclear charge increases while electrons enter the same outer shell, so attraction on the outer electrons strengthens and pulls the shell in. Radius increases down a group: each element has an extra full inner shell and more shielding. Ionic radius mirrors this: cations are smaller than their atoms (one shell lost), anions larger (extra repulsion in the filled outer shell); ions with the same electron configuration shrink as nuclear charge rises (N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺).
First ionisation energy increases across a period (rising Zeff, constant shielding) with two dips: aluminium below magnesium (3p electron in a higher-energy sub-shell) and sulfur below phosphorus (paired 3p electrons repel). It decreases down a group (extra shell: more distance and shielding). Full explanations with the three factors (nuclear charge, distance, shielding) are the standard 4-marker.
Melting points across Period 3 explain structure types: Na, Mg, Al — metallic, MP rises left to right (ion charge +1 → +3 and more delocalised electrons per atom → stronger metallic bonding; Al melts highest of the three). Si — giant covalent, highest MP of the period (strong covalent bonds must break). P₄, S₈, Cl₂ — simple molecular, low MPs depending on molecules per mole/electrons (S₈ > P₄ > Cl₂: more electrons, stronger London forces). Ar — monatomic, lowest. The Na→Ar MP graph is the standard interpretation question: identify the structure type from the shape of the graph and justify each by the forces broken on melting.
Group 2: Trends, Reactions and Solubility
Group 2 (alkaline earth) metals form 2+ ions by losing two s electrons; reactivity increases down the group (ionisation energies fall: extra shell, more shielding — outer electrons lost more easily). Reactions: with water — Mg reacts very slowly with cold water (faster with steam: Mg + H₂O(g) → MgO + H₂), Ca reacts steadily with cold water (Ca + 2H₂O → Ca(OH)₂ + H₂), Ba vigorously; with oxygen — 2Mg + O₂ → 2MgO (bright white flame); with dilute acids — Mg + 2HCl → MgCl₂ + H₂, rate increasing down the group.
Group 2 oxides are basic, reacting with water to form alkaline hydroxides: CaO + H₂O → Ca(OH)₂. Hydroxide solubility increases down the group (Mg(OH)₂ sparingly soluble, Ba(OH)₂ very soluble) while sulfate solubility decreases down the group (MgSO₄ soluble, BaSO₄ insoluble). The solubility trends link to hydration/lattice enthalpies: down the group the ions get larger, the hydration enthalpy becomes less exothermic — the balance with lattice enthalpy differs for hydroxides and sulfates, giving opposite trends. BaSO₄'s insolubility is the basis of the sulfate test (white precipitate with Ba²⁺) and of barium meals (BaSO₄ too insoluble to be toxic, X-ray contrast agent).
Uses to quote: Mg(OH)₂ as an antacid (milk of magnesia — neutralises excess stomach acid); Ca(OH)₂ in agriculture (lime, neutralises acidic soil); CaCO₃ in the lime cycle (CaCO₃ → CaO + CO₂ on heating; CaO + H₂O → Ca(OH)₂; Ca(OH)₂ + CO₂ → CaCO₃ + H₂O); BaSO₄ in medicine. Trend explanations must always invoke ionisation energy, atomic radius and shielding — 'reacts more vigorously because outer electrons are lost more easily down the group'.
Writing trend explanations for full marks: every Group 2/Group 7 trend answer needs the causal chain — state the trend, name the factor that changes (ionic/atomic radius, shielding, nuclear charge), link to the chemistry (electrons lost more easily, attraction for incoming electron weaker), conclude. 'Barium reacts faster because it is more reactive' scores nothing; 'barium's outer electrons are lost more easily because its larger atomic radius and greater shielding reduce the attraction of the nucleus, so the first two ionisation energies are lower' scores fully. Displacement answers must name the redox change: the displacing halogen gains electrons (reduced, oxidation number 0 → −1) and the halide loses them (oxidised, −1 → 0) — writing only the colour change loses the chemistry marks. Where equations are requested, include state symbols and the correct balancing (Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂) — ionic equations are expected, not full formula equations with spectator ions.
Group 7: Trends, Displacement, Tests and Disproportionation
Group 7 (halogens) exist as diatomic molecules (F₂, Cl₂, Br₂, I₂) with increasing melting/boiling points down the group (more electrons, stronger London forces: Cl₂ gas, Br₂ liquid, I₂ solid at room temperature). Reactivity and oxidising power decrease down the group: atomic radius and shielding increase, so the attraction for the incoming electron falls — the ability to gain an electron (be reduced) weakens. Electronegativity likewise falls down the group.
Displacement reactions: a more reactive halogen displaces a less reactive halide. Cl₂ displaces Br⁻ and I⁻ (Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, solution turns orange; Cl₂ + 2I⁻ → 2Cl⁻ + I₂, brown); Br₂ displaces I⁻ only; I₂ displaces none. This is redox — the displacing halogen is reduced, the halide oxidised. The trend is confirmed by E° values: E°(F₂/F⁻) = +2.87 V down to E°(I₂/I⁻) = +0.54 V.
Disproportionation — the same element simultaneously oxidised and reduced — is the key reaction of chlorine: in water, Cl₂ + H₂O ⇌ HCl + HClO (chlorine 0 → −1 and +1); with cold dilute NaOH, Cl₂ + 2NaOH → NaCl + NaClO + H₂O (also 0 → −1/+1), making household bleach (NaClO). Benefits of chlorinating water: kills bacteria, preventing waterborne disease; risks: HClO can react with organic matter forming potentially carcinogenic chlorinated hydrocarbons — the benefit/risk evaluation is a standard 4-6 marker. Adding a small amount of NaOH reduces the risk while keeping disinfection.
Halide tests: add dilute HNO₃ (removes interfering CO₃²⁻/OH⁻), then AgNO₃: AgCl white, AgBr cream, AgI yellow, all insoluble; solubility decreases down the group. Confirm with dilute NH₃ (dissolves AgCl only) then conc. NH₃ (dissolves AgBr, not AgI). The nitric acid step matters — using HCl would add Cl⁻ and give a false white precipitate. Halide ion reducing power increases down the group (larger ions lose electrons more easily): concentrated H₂SO₄ tests show NaCl gives HCl only (no redox), NaBr gives HBr + Br₂ (SO₂ produced), NaI gives HI + I₂ + H₂S/H₂SO₄ reduced further — fumes and colour changes identify the halide.
Period 3 Oxides: Bonding, Structure and Acid-Base Behaviour
Period 3 oxides show the periodic shift from ionic to covalent bonding: Na₂O and MgO are ionic (giant lattices, high melting points: 1275 °C and 2852 °C); Al₂O₃ is ionic but with significant covalent character (polarisation of O²⁻ by the small, highly charged Al³⁺ — amphoteric); SiO₂ is giant covalent (very high MP, 1610 °C); P₄O₁₀, SO₂, SO₃ are simple molecular (low MPs: 340 °C, sublimes, 17 °C). Melting point falls sharply after Al₂O₃ as bonding changes from giant to molecular.
Acid-base behaviour: Na₂O and MgO are basic — react with water giving alkaline solutions (Na₂O + H₂O → 2NaOH, pH ~14; MgO + H₂O → Mg(OH)₂, pH ~9 sparingly soluble) and neutralise acids (Na₂O + 2HCl → 2NaCl + H₂O). SiO₂, P₄O₁₀, SO₂ and SO₃ are acidic — react with water/alkalis giving acidic solutions (SO₂ + H₂O → H₂SO₃, pH ~1; SO₃ + H₂O → H₂SO₄, pH ~0-1; P₄O₁₀ + 6H₂O → 4H₃PO₄; SiO₂ reacts with hot concentrated NaOH: SiO₂ + 2NaOH → Na₂SiO₃ + H₂O). Al₂O₃ is amphoteric — reacts with both acids (Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O) and bases (Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄) but is insoluble in water. Acidic character increases across the period because electronegativity difference between element and oxygen falls — bonding shifts covalent.
Explanations earn marks by linking bonding to behaviour: ionic oxides contain O²⁻ which accepts protons (basic); covalent oxides of non-metals form oxyacids because the E–O bond polarises water. The classic exam answer: explain why Na₂O has a higher melting point than SO₃ — giant ionic lattice with strong electrostatic forces vs simple molecular with weak London forces. Also quote environmental chemistry: SO₂ from burning fossil fuels causes acid rain (dissolves in atmospheric water as H₂SO₃/H₂SO₄); MgO/CaO used in flue gas desulfurisation to remove SO₂.
Linking oxide behaviour to structure earns the evaluation marks: ionic oxides (Na₂O, MgO) contain O²⁻ ions which are strong proton acceptors — hence basic; covalent oxides of non-metals (SO₃, P₄O₁₀) form oxyacids because the polarised E–O bonds transfer proton character to water — hence acidic. The trend across the period mirrors electronegativity: as the element becomes more electronegative, the difference with oxygen falls, bonding becomes covalent and the oxide becomes acidic. Al₂O₃ sits at the crossover — its small, highly charged Al³⁺ polarises the oxide ion, giving covalent character and amphoteric behaviour (reacts with both acids and bases). Environmental applications complete the standard answers: SO₂/SO₃ from fossil-fuel combustion dissolve in atmospheric moisture forming acid rain (H₂SO₃/H₂SO₄); MgO and CaO remove SO₂ from flue gases (flue gas desulfurisation: CaO + SO₂ → CaSO₃); CO₂ from combustion slightly acidifies rainwater (carbonic acid) — quote these links because application marks are embedded in most Period 3 oxide questions.
Required Practical: Test-Tube Reactions and Ion Identification
Required practical (AQA RP4, OCR PAG4): use test-tube reactions to identify cations and anions and to compare trends. Standard identification scheme: cations — add NaOH(aq) dropwise then in excess: Mg²⁺ gives a white precipitate insoluble in excess; Ca²⁺ slight white precipitate (often none at low concentration); Al³⁺ white precipitate dissolving in excess NaOH (amphoteric hydroxide: Al(OH)₃ + OH⁻ → [Al(OH)₄]⁻); NH₄⁺ warm with NaOH and test the gas with damp red litmus (turns blue — ammonia released: NH₄⁺ + OH⁻ → NH₃ + H₂O); Group 1 cations flame tests (Li crimson, Na yellow, K lilac). Anions — sulfate: acidify with HNO₃ then BaCl₂/Ba(NO₃)₂, white precipitate BaSO₄; halides: HNO₃ then AgNO₃ (white/cream/yellow); carbonate: dilute acid fizzing (CO₂ turns limewater milky); hydroxide: pH paper/red litmus blue.
Trend experiments to describe: comparing Group 2 metals with water (rate of fizzing increases Mg → Ca → Ba) and measuring the pH of the resulting hydroxide solutions (rises down the group as hydroxides become more soluble); halogen displacement in solution (Cl₂ displaces Br⁻ and I⁻; observe orange/brown colour changes) — the observations confirm the oxidising-power trend. Writing up: state the reagent, the observation with colour, and the conclusion (ion identified) — 'white precipitate insoluble in excess NaOH' identifies Mg²⁺, but 'white precipitate' alone is ambiguous with Ca²⁺/Al³⁺ before the excess test.
Exam technique: safety points earn marks — gloves/goggles for concentrated NaOH; fume cupboard for halogens (Cl₂ toxic); never mix AgNO₃ with ammonia carelessly. Precision points: use fresh samples, control volumes, repeat observations. The identification scheme rewards knowing the 'excess' step — dissolving the Al(OH)₃ precipitate is the discriminating observation between Al³⁺ and Mg²⁺, and it links directly to the amphoteric behaviour of Al₂O₃/Al(OH)₃ covered in Period 3 oxides.
💡 Worked Example
Exam-style: State and explain the trend in reactivity of the Group 2 metals down the group, and predict what you would observe when barium is added to cold water.
Mark-scheme logic: Reactivity increases down the group. Down the group each element has an extra full inner shell, so atomic radius and shielding both increase and the first (and second) ionisation energies fall — the two outer s electrons are lost more easily, so the metals react more readily. With barium in cold water: reacts vigorously, fizzing as hydrogen gas is released, and the solution becomes alkaline (Ba(OH)₂ is soluble): Ba + 2H₂O → Ba(OH)₂ + H₂. Mark points: trend stated; extra shell/shielding; ionisation energy decrease; observation with equation. (AQA 3.2.2.1-3.2.2.2; OCR 3.1.2.1)
❓ Practice Questions & Mark Schemes
Write equations for the reactions of (a) Cl₂ with water, (b) Cl₂ with cold dilute NaOH, and state the oxidation state changes.
Answer: (a) Cl₂ + H₂O ⇌ HCl + HClO: chlorine goes 0 → −1 and +1 (disproportionation). (b) Cl₂ + 2NaOH → NaCl + NaClO + H₂O: also 0 → −1 and +1. Both are disproportionation reactions; (b) makes bleach (NaClO).
Describe tests to distinguish NaCl, NaBr and NaI solutions.
Answer: Add dilute HNO₃ then AgNO₃: AgCl white, AgBr cream, AgI yellow (all insoluble). Confirm with dilute NH₃ (dissolves AgCl only) then concentrated NH₃ (dissolves AgBr, not AgI). HNO₃ removes interfering CO₃²⁻/OH⁻; using HCl would give a false positive with Cl⁻.
Explain why Group 7 oxidising power decreases down the group, and predict whether Br₂ displaces I⁻ from solution.
Answer: Down the group atomic radius and shielding increase, so the attraction of the halogen atom for an incoming electron falls — the ability to be reduced (oxidise others) weakens. E° falls from +2.87 V (F₂/F⁻) to +0.54 V (I₂/I⁻). Br₂ is a stronger oxidant than I₂, so yes — Br₂ displaces I⁻: Br₂ + 2I⁻ → 2Br⁻ + I₂ (solution turns brown).
Explain why Al₂O₃ is described as amphoteric and write equations for its reactions with HCl and NaOH.
Answer: Al₂O₃ reacts with both acids and bases: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O (basic behaviour) and Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄ (acidic behaviour). It is insoluble in water. Its small, highly charged Al³⁺ polarises O²⁻, giving significant covalent character to the ionic lattice.
Explain the opposing trends in solubility of Group 2 hydroxides and sulfates.
Answer: Hydroxide solubility increases down the group (Mg(OH)₂ sparingly soluble, Ba(OH)₂ very soluble); sulfate solubility decreases (MgSO₄ soluble, BaSO₄ insoluble). Down the group, ions get larger so hydration enthalpy becomes less exothermic; the balance between hydration and lattice enthalpy differs for the two anions, producing opposite trends. BaSO₄'s insolubility underlies the sulfate test and barium meals.
Explain why Na₂O has a much higher melting point than SO₃.
Answer: Na₂O is a giant ionic lattice: strong electrostatic forces between Na⁺ and O²⁻ must be overcome to melt it (1275 °C). SO₃ is simple molecular: only weak London forces between molecules, so it melts at 17 °C. The bonding changes from ionic to covalent across the period as the electronegativity difference with oxygen falls.