# Redox Chemistry and Groups 1, 2 and 7

> Edexcel International A-Level Chemistry · Edexcel IAL Chemistry Unit 2 (WCH12)
> Source: https://www.owlsprep.com/study/edexcel-ial-chemistry-u2-redox-chemistry-and-groups-1/

This guide covers all Edexcel IAL Unit 2 content for redox chemistry, Groups 1, 2 and 7, including oxidation number calculations, group trend explanations, core titration practicals, and common ion identification tests required for the WCH12 exam.

**Prerequisites:** [Ionic compound formula writing and balancing equations](https://www.owlsprep.com/study/edexcel-ial-chemistry-u1-ionic-equations/); [Atomic structure and periodic trends (ionisation energy, electronegativity)](https://www.owlsprep.com/study/edexcel-ial-chemistry-u1-periodic-trends/)

## Learning objectives

- Apply oxidation number rules to calculate values for peroxides, metal hydrides and polyatomic ions
- Write balanced ionic half-equations and full redox equations for s/p block reactions
- Explain Group 1/2 trends in reactivity, solubility of hydroxides/sulfates, and thermal stability of carbonates/nitrates
- Perform acid-base titration concentration calculations and interpret volumetric analysis uncertainty
- Recall Group 7 trends in physical properties, reactivity, and disproportionation reactions
- Identify common ions (carbonate, sulfate, ammonium, halides) using standard test procedures

## Redox Chemistry Fundamentals

**Oxidation Number** — A number assigned to an atom in a compound to represent the number of electrons it has lost or gained relative to its neutral state. Core rules: F=-1, O=-2 (except peroxides = -1, OF₂ = +2), H=+1 (except metal hydrides = -1), sum of oxidation numbers equals the charge on the species.

*Example:* Oxidation number of S in H₂SO₄ = +6

> **mnemonic**
>
> OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain of electrons

Redox reactions involve a transfer of electrons: the reducing agent donates electrons (is oxidised, oxidation number increases), while the oxidising agent accepts electrons (is reduced, oxidation number decreases).

**Disproportionation** — A redox reaction where the same element is both oxidised and reduced in the same reaction.

*Example:* $Cl_2 + H_2O \rightarrow HCl + HClO$: Cl changes from 0 to -1 (reduction) and +1 (oxidation)

**Worked example:** Calculate the oxidation number of chlorine in $ClO_3^-$, and write the half-equation for its reduction to $Cl^-$ in acidic conditions.

1. Let $x$ = oxidation number of Cl. Sum of oxidation numbers = charge on ion: $x + 3(-2) = -1$ so $x = +5$.
2. Balance Cl atoms: $ClO_3^- \rightarrow Cl^-$
3. Balance O atoms by adding $H_2O$: $ClO_3^- \rightarrow Cl^- + 3H_2O$
4. Balance H atoms by adding $H^+$: $ClO_3^- + 6H^+ \rightarrow Cl^- + 3H_2O$
5. Balance charge by adding electrons: +5 charge on left, -1 on right, so add 6$e^-$ to left: $ClO_3^- + 6H^+ + 6e^- \rightarrow Cl^- + 3H_2O$

> **Exam tip:** Always show full working for oxidation number calculations, as method marks are awarded even if your final answer is incorrect.

## Group 1 and 2 Chemistry

Reactivity increases down Groups 1 and 2 because first ionisation energy decreases: larger atomic radius and more electron shielding reduce the attraction between the nucleus and outer electrons, making electrons easier to lose during reactions.

**Thermal Stability** — The ability of a compound to resist decomposition when heated. For Group 1/2 carbonates and nitrates, stability increases down the group: larger cations have lower polarising power, so they cause less distortion of the anion electron cloud, making decomposition require more energy.

> **Flame Test Procedure**
>
> Clean nichrome wire with concentrated HCl and heat to red hot before dipping in the sample to avoid contamination from previous tests. Recall required flame colours: Li = crimson red, Na = yellow, K = lilac, Ca = brick red, Sr = red, Ba = green.

**Worked example:** Write balanced equations for (a) the reaction of strontium (Group 2) with water, and (b) the thermal decomposition of strontium nitrate.

1. Group 2 metal + water → metal hydroxide + hydrogen: $Sr + 2H_2O \rightarrow Sr(OH)_2 + H_2$
2. Group 2 nitrate decomposition produces metal oxide, nitrogen dioxide and oxygen: $2Sr(NO_3)_2 \rightarrow 2SrO + 4NO_2 + O_2$

Key solubility trends: Group 2 hydroxides increase in solubility down the group (Mg(OH)₂ is almost insoluble, Ba(OH)₂ is highly soluble). Group 2 sulfates decrease in solubility down the group (MgSO₄ is soluble, BaSO₄ is almost insoluble).

> **Exam tip:** Solubility trends are often tested with sulfate ion test questions: the ionic equation for a positive sulfate test is $Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)$ (white precipitate).

## Group 7 (Halogen) Chemistry

Key trends down Group 7: melting/boiling point increases (stronger London dispersion forces as molecule size increases), electronegativity decreases (larger atomic radius reduces attraction for bonding electrons), and reactivity decreases (harder to gain an outer electron).

**Halogen Displacement Reaction** — A redox reaction where a more reactive halogen (higher up Group 7) displaces a less reactive halide ion from its compound, producing a characteristic colour change in aqueous or non-polar solvent.

| Halogen | Aqueous solution colour | Hexane (non-polar) layer colour |
| --- | --- | --- |
| $Cl_2$ | Pale green | Pale green |
| $Br_2$ | Orange/yellow | Orange/red |
| $I_2$ | Brown | Purple/violet |

**Worked example:** Write the full ionic equation for the reaction between chlorine water and potassium iodide solution, and state the colour change in the hexane layer after shaking.

1. Cl₂ is more reactive than I₂, so it oxidises I⁻ to I₂ and is reduced to Cl⁻. Half equations: $Cl_2 + 2e^- \rightarrow 2Cl^-$ and $2I^- \rightarrow I_2 + 2e^-$
2. Combine half equations to cancel electrons: $Cl_2 + 2I^- \rightarrow 2Cl^- + I_2$
3. Colour change in hexane: from colourless to purple/violet

Common disproportionation reactions of chlorine: (1) Cold dilute NaOH: $Cl_2 + 2NaOH \rightarrow NaCl + NaClO + H_2O$ (bleach production), (2) Water: $Cl_2 + H_2O \rightleftharpoons HCl + HClO$ (water treatment, HClO kills bacteria), (3) Hot concentrated NaOH: $3Cl_2 + 6NaOH \rightarrow 5NaCl + NaClO_3 + 3H_2O$.

> **Exam tip:** Always state the oxidation number changes for the element undergoing disproportionation to gain full marks in exam questions.

## Volumetric Analysis and Ion Tests

**Standard Solution** — A solution of accurately known concentration, prepared by dissolving a measured mass of primary standard solid in distilled water and making up to a fixed volume in a volumetric flask.

Core Practical 4 uses a standard solution of a solid acid (e.g. ethanedioic acid) to find the concentration of a NaOH solution via titration. Indicators: phenolphthalein (pink in alkali, colourless in acid) or methyl orange (yellow in alkali, red in acid).

**Worked example:** 25.0 cm³ of NaOH solution is titrated against 0.100 mol dm⁻³ ethanedioic acid ($H_2C_2O_4$, diprotic). The mean titre is 22.4 cm³. Calculate the concentration of NaOH.

1. Balanced equation: $H_2C_2O_4 + 2NaOH \rightarrow Na_2C_2O_4 + 2H_2O$
2. Moles of ethanedioic acid = $0.100 \times 0.0224 = 0.00224$ mol
3. Mole ratio NaOH:$H_2C_2O_4$ = 2:1, so moles NaOH = $0.00224 \times 2 = 0.00448$ mol
4. Concentration of NaOH = $0.00448 / 0.025 = 0.179$ mol dm⁻³ (3 s.f.)

> **Ion Test Protocols**
>
> Always add dilute acid before BaCl₂ for sulfate tests to eliminate carbonate impurities (which also form white precipitates with Ba²⁺). Use dilute HNO₃, not HCl, to acidify samples for halide tests to avoid introducing Cl⁻ ions that cause false positives.

- Carbonate test: Add dilute acid, bubble gas through limewater (turns cloudy). Ionic equation: $CO_3^{2-} + 2H^+ \rightarrow CO_2 + H_2O$
- Ammonium test: Add NaOH, warm, gas turns damp red litmus blue. Ionic equation: $NH_4^+ + OH^- \rightarrow NH_3 + H_2O$

> **Exam tip:** Give titration calculation answers to the same number of significant figures as the least precise value given in the question, usually 3 s.f.

## Common pitfalls

- **Wrong:** Writing oxidation numbers without the sign first, e.g. writing 2 instead of +2 for Mg²⁺
  - Why it fails: Oxidation numbers represent electron gain/loss, so the sign is essential to distinguish oxidation from reduction
  - Correct: Always write the sign before the magnitude for oxidation numbers, e.g. +2, -1
- **Wrong:** Forgetting that Group 2 nitrate decomposition produces O₂ as well as NO₂ and metal oxide
  - Why it fails: Unbalanced equations lose marks, and O₂ is a required product of full Group 2 nitrate decomposition
  - Correct: Use the general equation: $2M(NO_3)_2 \rightarrow 2MO + 4NO_2 + O_2$ for Group 2 nitrates
- **Wrong:** Using HCl instead of HNO₃ before adding AgNO₃ for halide tests
  - Why it fails: HCl contains Cl⁻ ions that produce a false positive white AgCl precipitate
  - Correct: Always use dilute nitric acid to acidify halide test samples to remove impurities without introducing Cl⁻
- **Wrong:** Stating Group 2 sulfate solubility increases down the group
  - Why it fails: This is a common confusion with the Group 2 hydroxide solubility trend, which does increase down the group
  - Correct: Recall: Group 2 hydroxides = more soluble down, sulfates = less soluble down (BaSO₄ is highly insoluble)
- **Wrong:** Rounding titration calculation results to 2 significant figures when given data is to 3 s.f.
  - Why it fails: Mark schemes explicitly penalise incorrect significant figures for quantitative calculations
  - Correct: Give answers to the same number of significant figures as the least precise value provided in the question

## Cheatsheet

| Concept | Key Exam Fact |
| --- | --- |
| Oxidation number rules | F=-1, O=-2 (peroxides=-1, OF₂=+2), H=+1 (metal hydrides=-1) |
| Disproportionation | Same element is both oxidised and reduced in one reaction |
| Group 2 solubility | Hydroxides ↑ down group, sulfates ↓ down group |
| Halide test results | AgCl=white (soluble dilute NH₃), AgBr=cream (conc NH₃), AgI=yellow (insoluble) |
| Titration indicators | Phenolphthalein: pink→colourless, Methyl orange: yellow→red for acid→base titrations |
| Flame colours | Li=red, Na=yellow, K=lilac, Ca=brick red, Ba=green |

## What's next

You have now mastered all core content for Redox Chemistry and Groups 1, 2, and 7 for Edexcel IAL Chemistry Unit 2. This topic forms the foundation for more advanced redox content in Unit 4, including electrode potentials and electrochemical cells, as well as periodicity extensions in higher units. To consolidate your learning, practise past paper questions focused on group trend explanations, titration calculations, and ion identification tests, making sure you follow the mark scheme wording for command terms like 'explain' and 'describe'. Pay special attention to core practical write-ups for titrations and flame tests, as these make up ~15% of Unit 2 exam marks.

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