Pearson Edexcel International GCSE in Chemistry (4CH1) Specification Study Notes

Pearson Edexcel International GCSE in Chemistry (4CH1) Specification Overview

• Qualification Title: Pearson Edexcel International GCSE in Chemistry (4CH1).

• First Teaching: September 2017.

• First Examination: June 2019.

• Linear Qualification: All assessments are taken at the end of the course in one examination series.

• Specification Issue: Issue 3 (September 2024).

• Accreditation Status: Not accredited or regulated by any UK regulatory body.

• Content Consistency: Equivalent in standard to Pearson’s regulated GCSE in Chemistry, but designed for an international context.

Qualification Structure and Assessment at a Glance

The qualification consists of two externally-assessed papers: Chemistry Paper 1 and Chemistry Paper 2.

Chemistry Paper 1

• Paper Code: 4CH1/1C and 4SD0/1C.

• Topic Coverage: Assesses core content only (content without bold text or a ‘C’ reference).

• Weighting: $61.1\%$ of the total qualification.

• Duration: 2 hours.

• Total Marks: 110.

• Availability: November and June.

• Assessment Style: A mixture of multiple-choice, short-answer, calculations, and extended open-response questions.

• Calculator Use: Permitted.

Chemistry Paper 2

• Paper Code: 4CH1/2C.

• Topic Coverage: Assesses all content, including core and depth content (designated by bold text and ‘C’ references).

• Weighting: $38.9\%$ of the total qualification.

• Duration: 1 hour and 15 minutes.

• Total Marks: 70.

• Assessment Style: Similar to Paper 1, including multiple-choice, short-answer, calculations, and extended open-response questions.

• Depth: Bold statements in the specification indicate sub-topics addressed in greater depth.

Topic 1: Principles of Chemistry

(a) States of Matter

• Three States: Understand arrangements, movement, and energy of particles in solids, liquids, and gases.

• Interconversions: Changes between states, including names of processes, required energy changes, and particle rearrangement.

• Evidence for Particles:

  • Dilution of coloured solutions.

  • Diffusion of gases.

• Terminology:

  • Solvent: The liquid in which a solute dissolves.

  • Solute: The substance that dissolves in a liquid.

  • Solution: The mixture formed by a solute and solvent.

  • Saturated Solution: A solution where no more solute can be dissolved at a specific temperature.

• Solubility (Depth Content - 1.5C, 1.6C, 1.7C):

  • Measured in units of $g \, \text{per} \, 100\,g \, \text{of solvent}$.

  • Construction and interpretation of solubility curves.

  • Practical: Investigate the solubility of a solid in water at a specific temperature.

(b) Elements, Compounds, and Mixtures

• Classification: Differentiating between elements, compounds, and mixtures.

• Purity: Pure substances have fixed melting and boiling points; mixtures melt or boil over a range of temperatures.

• Separation Techniques:

  • Simple distillation.

  • Fractional distillation.

  • Filtration.

  • Crystallisation.

  • Paper chromatography.

• Chromatography: Interpretation of chromatograms and calculation of $R_f$ values to identify mixture components.

  • $R_f = \frac{\text{distance moved by the substance}}{\text{distance moved by the solvent}}$

• Practical: Investigate paper chromatography using inks/food colourings.

(c) Atomic Structure

• Definitions: Differentiation between atoms and molecules.

• Sub-atomic Particles:

  • Protons: Mass $1$, Charge $+1$, Position (nucleus).

  • Neutrons: Mass $1$, Charge $0$, Position (nucleus).

  • Electrons: Mass $\frac{1}{1836}$, Charge $-1$, Position (shells surrounding nucleus).

• Nuclear Terms:

  • Atomic Number: Number of protons in the nucleus.

  • Mass Number: Sum of protons and neutrons.

  • Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons.

  • Relative Atomic Mass ($A_r$): The weighted mean mass of an atom of an element relative to $\frac{1}{12}$ the mass of an atom of carbon-12.

• Calculation: Calculate $A_r$ from isotopic abundances.

(d) The Periodic Table

• Arrangement: Organized by atomic number, groups (vertical columns), and periods (horizontal rows).

• Electronic Configuration: Deducing configuration for the first 20 elements (e.g., Sodium is $2, 8, 1$).

• Classification: Identifying metals vs. non-metals based on electrical conductivity, acid-base character of oxides, and position relative to the “staircase.”

• Group Characteristics: Similar chemical properties within a group due to the same number of electrons in the outer shell.

• Group 0 (Noble Gases): Unreactive due to having a full outer shell (stable electronic configuration).

(e) Chemical Formulae, Equations, and Calculations

• Equations: Writing balanced chemical equations and word equations, including state symbols: $(s)$, $(l)$, $(g)$, $(aq)$.

• Relative Formula Mass ($M_r$): Calculated from $A_r$ values.

• The Mole (mol): The unit for the amount of a substance.

  • $\text{Mass} = \text{moles} \times M_r$

• Yield: Calculation of percentage yield.

  • $\text{Percentage Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$

• Empirical vs. Molecular Formulae:

  • Empirical Formula: Simplest whole-number ratio of atoms of each element in a compound.

  • Molecular Formula: Actual number of atoms of each element in a molecule.

• Depth Content (1.34C – 1.35C):

  • Calculations involving concentration: $\text{moles} = \text{volume} (dm^3) \times \text{concentration} (mol/dm^3)$.

  • Gas volumes: $1 \, \text{mole}$ of gas occupies $24\,dm^3$ ($24,000\,cm^3$) at room temperature and pressure (rtp).

• Practical: Determine the formula of a metal oxide (e.g., magnesium oxide) via combustion or reduction.

(f) Ionic Bonding

• Formation: Ions formed by loss or gain of electrons to achieve full outer shells.

• Specific Ion Charges:

  • Group 1 ($+1$), Group 2 ($+2$), Group 3 ($+3$).

  • Group 5 ($-3$), Group 6 ($-2$), Group 7 ($-1$).

  • Others: $Ag^+$, $Cu^{2+}$, $Fe^{2+}$, $Fe^{3+}$, $Pb^{2+}$, $Zn^{2+}$, $H^+$, $\text{OH}^-$, $\text{NH}_4^+$, $\text{CO}_3^{2-}$, $\text{NO}_3^-$, $\text{SO}_4^{2-}$.

• Bonding Nature: Electrostatic attraction between oppositely charged ions.

• Properties: Giant ionic lattices result in high melting/boiling points. Brittle. Conduct electricity only when molten or in aqueous solution.

(g) Covalent Bonding

• Formation: Sharing a pair of electrons between atoms.

• Bonding Nature: Electrostatic attraction between the shared pair of electrons and the nuclei of the involved atoms.

• Simple Molecular Structures: Gases, liquids, or low-melting solids due to weak intermolecular forces. Boiling points generally increase with increased relative molecular mass.

• Giant Covalent Structures: High melting/boiling points (e.g., diamond, graphite, $C_{60}$ fullerene).

  • Diamond: Hard, non-conductive.

  • Graphite: Soft, conductive (delocalised electrons).

  • $C_{60}$ Fullerene: Simple molecular, non-conductive.

(h) Metallic Bonding (Depth Content)

• 1.52C/1.53C/1.54C: Metallic lattice represented as a 2-D diagram of positive ions in a “sea” of delocalised electrons.

• Properties: Good electrical conductors (mobile electrons), malleable (layers of ions can slide).

(i) Electrolysis (Depth Content)

• Fundamental Terms:

  • Anion: Negative ion.

  • Cation: Positive ion.

• Process: Breakdown of an electrolyte (molten or aqueous ionic compound) using electricity.

• Experiments: Electrolysis of molten lead(II) bromide, aqueous sodium chloride, dilute sulfuric acid, and copper(II) sulfate.

• Ionic Half-equations: Representing oxidation (loss of electrons) and reduction (gain of electrons) at electrodes.

Topic 2: Inorganic Chemistry

(a-b) Group 1 and Group 7

• Group 1 (Alkali Metals): Lithium, Sodium, Potassium. Reactivity increases down the group. Explained by the increasing distance of the outer electron from the nucleus (1.24C).

• Group 7 (Halogens): Chlorine (gas, green), Bromine (liquid, red-brown), Iodine (solid, grey/purple vapor). Reactivity decreases down the group.

• Displacement: A more reactive halogen will displace a less reactive halogen from its halide solution.

(c) Gases in the Atmosphere

• Composition of Dry Air: Approximately $78\%$ Nitrogen, $21\%$ Oxygen, $0.9\%$ Argon, and $0.04\%$ Carbon Dioxide.

• Experiments: Determining Oxygen percentage using phosphorus or iron.

• Greenhouse Gases: $CO_2$ contributes to climate change.

(d) Reactivity Series

• Order: Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, Zinc, Iron, Copper, Silver, Gold.

• Redox Definitions:

  • Oxidation: Gain of oxygen or loss of electrons.

  • Reduction: Loss of oxygen or gain of electrons.

• Rusting of Iron: Requires both water and oxygen. Prevention through barriers, galvanising (zinc coating), or sacrificial protection (more reactive metal).

(e) Extraction and Uses of Metals (Depth Content)

• Extraction: Metals below carbon are extracted by reduction with carbon (e.g., iron). Metals above carbon require electrolysis (e.g., aluminium).

• Alloys: Mixtures of metals that are harder than pure metals because differently sized atoms disrupt the regular lattice, preventing layers from sliding.

(f-g) Acids, Bases, and Salts

• Indicators: Litmus (red in acid, blue in alkali), Phenolphthalein (colourless in acid, pink in alkali), Methyl orange (red in acid, yellow in alkali).

• pH Scale: $0-3$ (strong acid), $4-6$ (weak acid), $7$ (neutral), $8-10$ (weak alkali), $11-14$ (strong alkali).

• Proton Transfer: Acids are proton ($H^+$) donors; bases are proton acceptors.

• Solubility Rules (2.34):

  • All Sodium, Potassium, Ammonium compounds and all Nitrates are soluble.

  • Chlorides are soluble (except Silver and Lead).

  • Sulfates are soluble (except Barium, Calcium, and Lead).

  • Carbonates and Hydroxides are generally insoluble (except Sodium, Potassium, and Ammonium).

(h) Chemical Tests

• Gases:

  • $H_2$: Squeaky pop with lighted splint.

  • $O_2$: Relights a glowing splint.

  • $CO_2$: Turns limewater cloudy.

  • $NH_3$: Turns damp red litmus paper blue.

  • $Cl_2$: Bleaches damp blue litmus paper.

• Flame Tests: $Li^+$ (red), $Na^+$ (yellow), $K^+$ (lilac), Ca^{2+ (orange-red), Cu^{2+ (blue-green).

• Cation Tests: $Cu^{2+}$ (blue ppt), $Fe^{2+}$ (green ppt), $Fe^{3+}$ (brown ppt) with sodium hydroxide.

• Anion Tests: Halides with acidified silver nitrate ($Cl^-$ white, $Br^-$ cream, $I^-$ yellow). Sulfate with acidified barium chloride (white ppt).

Topic 3: Physical Chemistry

(a) Energetics

• Exothermic: Heat energy given out ($\Delta H$ is negative).

• Endothermic: Heat energy taken in ($\Delta H$ is positive).

• Calorimetry: $Q = mc\Delta T$.

  • $Q$ = heat energy change ($J$).

  • $m$ = mass of substance being heated ($g$).

  • $c$ = specific heat capacity ($4.18 \, J/g/^{\circ}C$ for water).

  • $\Delta T$ = change in temperature.

• Bond Energies (Depth Content): Bond-breaking is endothermic; bond-making is exothermic.

(b) Rates of Reaction

• Collision Theory: Rate depends on frequency of collisions and the energy of collisions.

• Factors: Surface area, concentration, temperature, pressure (gases), and catalysts.

• Catalyst: Increases rate by providing an alternative pathway with lower activation energy.

(c) Equilibria (Depth Content)

• Dynamic Equilibrium: Forward and reverse reactions occur at the same rate; concentrations remain constant (closed system).

• Le Chatelier's Principles (Applications):

  • Temperature increase shifts towards endothermic direction.

  • Pressure increase shifts towards the side with fewer gas moles.

Topic 4: Organic Chemistry

(a) Introduction

• Hydrocarbon: Compound of hydrogen and carbon only.

• Nomenclature: IUPAC rules for naming up to six carbon atoms.

• Isomerism: Same molecular formula, different structural formula.

(b) Crude Oil

• Fractional Distillation: Process separating crude oil into:

  • Refinery Gases: Bottled gas.

  • Gasoline: Petrol for cars.

  • Kerosene: Jet fuel.

  • Diesel: Fuel for trucks/trains.

  • Fuel Oil: Fuel for ships/heating.

  • Bitumen: Road surfacing.

• Cracking: Converts long-chain alkanes into shorter-chain alkanes and alkenes ($600-700\,^{\circ}C$, silica/alumina catalyst).

(c-d) Alkanes and Alkenes

• Alkanes: General formula $C_nH_{2n+2}$. Saturated. Substitution reaction with halogens (requires UV light).

• Alkenes: General formula $C_nH_{2n}$. Unsaturated ($C=C$ bond). Addition reactions with bromine water (decolourises from orange to colourless).

(e-g) Alcohols, Carboxylic Acids, and Esters (Depth Content)

• Alcohols: Functional group $-OH$. Ethanol produced by fermentation or reaction of ethene with steam ($300\,^{\circ}C$, $60-70 \, atm$, phosphoric acid).

• Carboxylic Acids: Functional group $-COOH$ (e.g., ethanoic acid in vinegar).

• Esters: Functional group $-COO-$. Formed from alcohol + carboxylic acid. Ethyl ethanoate is a key example. Volatile with sweet smells.

(h) Synthetic Polymers

• Addition Polymers: Formed from monomers (alkenes). Difficult to dispose of because they are inert (non-biodegradable).

• Condensation Polymers (Depth Content): Dicarboxylic acid + diol $\rightarrow$ polyester + water. Some are biodegradable (biopolyesters).

Assessment Objectives (AOs)

• AO1 ($38-42\%$): Knowledge and understanding of chemistry.

• AO2 ($38-42\%$): Application of knowledge, analysis, and evaluation.

• AO3 ($19-21\%$): Experimental skills, analysis, and evaluation of data/methods.

Experimental Skills and Mathematical Requirements

Experimental Skills assessed in written exams:

• Devise and plan investigations.

• Identify variables (independent, dependent, control).

• Assess reliability, accuracy, and validity.

• Safe and skillful practical techniques.

Mathematical Skills:

• Standard form and decimal form usage.

• Calculating arithmetic mean.

• Plotting and interpreting frequency tables, bar charts, and histograms.

• Using $\text{y} = \text{mx} + \text{c}$ for linear graphs.

• Determining slopes of tangents to curves for rates of change.

Command Word Taxonomy

• Calculate: Obtain a numerical answer, showing working.

• Describe: Give an account of (linked statements, no justification needed).

• Explain: Requires a justification/reasoning for a point.

• Evaluate: Review information to form a conclusion/judgment.

• Suggest: Propose a solution to a problem in a novel context.

• Deduce: Reach a conclusion from provided information.

Suggested Practical Investigations (Appendix 6)

• Investigate thermal decomposition of carbonates (Calcium, Zinc, Copper).

• Compare temperature rise of different fuels.

• Cracking of paraffin oil.

• Reactions of Group 2 metals.

• Precipitate formation tests.

• Building models of covalent molecules.

• Electroplating metal objects.

• Determine the molar volume of hydrogen gas via Magnesium + HCl.