CHEMISTRY Double science syllabus

1. Principles of Chemistry

(a) States of matter

• Understand the three states of matter in terms of the arrangement, movement, and energy of the particles.

• Interconversions between states:

  • Names of interconversions (melting, freezing, evaporation, condensation, sublimation, deposition).

  • Achieved through changes in temperature or pressure.

  • Changes in arrangement, movement, and energy of particles are factors.

• Explanation of experiments with dilution of colored solutions and diffusion of gases.

• Key terms:

  • Solvent: substance that dissolves a solute.

  • Solute: substance dissolved in a solvent.

  • Solution: mixture formed by a solute and solvent.

  • Saturated solution: solution containing the maximum amount of solute that can dissolve.

(b) Elements, Compounds, and Mixtures

• Classification of a substance as an element, compound, or mixture.

• Pure substances have fixed melting and boiling points; mixtures can melt or boil over a range.

• Techniques for separation:

  • Simple distillation.

  • Fractional distillation.

  • Filtration.

  • Crystallization.

  • Paper chromatography.

• Chromatograms provide information about the composition of a mixture.

• Calculation of R values to identify mixture components.

• Practical: Use paper chromatography with inks/food colorings.

(c) Atomic Structure

• Definitions:

  • Atom: smallest unit of an element.

  • Molecule: two or more atoms bonded together.

• Structure of an atom:

  • Protons (positive charge), neutrons (neutral), electrons (negative charge).

• Key terms:

  • Atomic number: number of protons.

  • Mass number: total number of protons and neutrons.

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

  • Relative atomic mass (Ar): weighted average mass of an atom compared to carbon-12.

• Calculation of Ar from isotopic abundances.

(d) The Periodic Table

• Arrangement of elements by:

  • Atomic number.

  • Groups (columns) and periods (rows).

• Deduction of electronic configurations for the first 20 elements.

• Classification of elements based on:

  • Electrical conductivity.

  • Acid-base character of oxides to categorize as metals or non-metals.

• Electronic configuration relates to position in the Periodic Table.

• Similar chemical properties of elements within the same group.

• Noble gases (Group 0) are chemically inert.

(e) Chemical Formulae, Equations, and Calculations

• Writing word equations and balanced chemical equations (including state symbols).

• Calculating relative formula masses (Mr) from Ar.

• Understanding moles (mol) as a unit for the amount of substance.

• Conducting calculations involving amount of substance, Ar, and Mr.

• Calculation of reacting masses from experimental data and chemical equations.

• Percentage yield calculation.

• Experimental determination of formulae of simple compounds (e.g., metal oxides, water, salts with water of crystallization).

• Definitions:

  • Empirical formula: simplest whole-number ratio of atoms in a compound.

  • Molecular formula: actual number of atoms of each element in a molecule.

• Calculation of empirical and molecular formulae from data.

• Practical: Determining the formula of a metal oxide via combustion/reduction (e.g., magnesium oxide, copper(II) oxide).

(f) Ionic Bonding

• Formation of ions through electron loss/gain.

• Charges of common ions:

  • Metals from Groups 1, 2, and 3.

  • Non-metals from Groups 5, 6, and 7.

  • Common ions: Ag+, Cu2+, Fe2+, Fe3+, Pb2+, Zn2+, H+, OH-, NH4+, CO3²-, NO3-, SO4²-.

• Writing formulae for ionic compounds using these ions.

• Dot-and-cross diagrams for ionic compounds (limited to Groups 1, 2, 3, 5, 6, 7).

• Understanding ionic bonding and electrostatic attractions.

• Properties: giant ionic lattices have high melting/boiling points; ionic compounds conduct electricity when molten or in solution, not as solids.

(g) Covalent Bonding

• Covalent bonds involve sharing of electron pairs.

• Understanding covalent bonds through electrostatic attractions.

• Dot-and-cross diagrams for:

  • Diatomic molecules (H2, O2, N2, halogens).

  • Inorganic molecules (water, ammonia, carbon dioxide).

  • Organic molecules (methane, ethane, ethene, halogen-containing compounds).

• Explanation of states of substances based on molecular structure.

• Factors affecting melting and boiling points, including intermolecular forces.

• Understanding properties of giant covalent structures (diamond, graphite, C60 fullerene) related to conductivity and hardness.

2. Inorganic Chemistry

(a) Group 1 (Alkali Metals)

• Elements: lithium, sodium, potassium.

• Recognize family properties via reactions with water.

• Differences in reactions with air and water indicate reactivity trends.

• Predict properties of other alkali metals based on trends.

(b) Group 7 (Halogens)

• Elements: chlorine, bromine, iodine.

• Understand colors, states (at room temperature), and trends in properties.

• Use trends in Group 7 for predicting properties of halogens.

• Displacement reactions involving halogens and halides as evidence for reactivity trends.

(c) Gases in the atmosphere

• Approximate volume percentages of four main gases in dry air.

• Experiments to determine oxygen percentage in air via metal/non-metal reactions.

• Combustion of elements in oxygen: magnesium, hydrogen, sulfur.

• Formation of carbon dioxide from thermal decomposition of metal carbonates.

• Carbon dioxide as a greenhouse gas contributing to climate change.

• Practical: measure oxygen percentage in air using metals/non-metals.

(d) Reactivity Series

• Arranging metals based on reaction with:

  • Water.

  • Dilute hydrochloric/sulfuric acid.

• Displacement reactions order of reactivity: K, Na, Li, Ca, Mg, Al, Zn, Fe, Cu, Ag, Au.

• Conditions leading to iron rusting and prevention methods: barrier, galvanizing, sacrificial protection.

• Key terms: oxidation, reduction, redox, oxidizing agent, reducing agent.

(e) Acids, Alkalis, and Titrations

• Use of litmus, phenolphthalein, and methyl orange to distinguish acidic/alkaline solutions.

• pH scale classification:

  • Strongly acidic (0-3), weakly acidic (4-6), neutral (7), weakly alkaline (8-10), strongly alkaline (11-14).

• Universal indicator for measuring approximate pH values.

• Acids provide hydrogen ions; alkalis provide hydroxide ions.

• Alkalis neutralize acids.

(f) Acids, Bases, and Salt Preparations

• General rules for ionic compound solubility:

  • Solubility of common sodium, potassium, ammonium compounds.

  • All nitrates are soluble.

  • Common chlorides (except Ag, Pb) are soluble.

  • Common sulfates (except Ba, Ca, Pb) are soluble.

  • Common carbonates (except Na, K, NH4) are insoluble.

  • Common hydroxides (except Na, K, Ca) are insoluble.

• Proton transfer definition of acids and bases:

  • Acid = proton donor; Base = proton acceptor.

• Reactions of acids with metals, bases, and carbonates to form salts.

• Describing preparation of pure, dry soluble salts from insoluble reactants.

• Practical: prepare hydrated copper(II) sulfate crystals from copper(II) oxide.

(g) Chemical Tests

• Tests for gases:

  • Hydrogen, oxygen, carbon dioxide, ammonia, chlorine.

• Flame test procedure and colors for cations:

  • Li+ (red), Na+ (yellow), K+ (lilac), Ca2+ (orange-red), Cu2+ (blue-green).

• Tests for cations:

  • NH4+ using NaOH;

  • Cu2+, Fe2+, Fe3+ using NaOH.

• Tests for anions:

  • Cl-, Br-, I- using acidified AgNO3;

  • SO4²- using acidified BaCl2;

  • CO3²- using HCl to identify gas evolved.

• Test for water purity: using anhydrous CuSO4 and a physical test for purity.

3. Physical Chemistry

(a) Energetics

• Exothermic reactions release heat; endothermic reactions absorb heat.

• Calorimetry experiments: combustion, displacement, dissolving, neutralization.

• Calculation of heat energy change (Q = mcΔT) and molar enthalpy change (ΔH).

• Practical: investigate temperature changes in dissolving salts, neutralizations, and combustion.

(b) Rates of Reaction

• Investigate effects of changes in surface area, concentration, temperature, and catalysts on reaction rates.

• Particle collision theory explained: changes in solid surface area, solution concentration, gas pressure, temperature.

• Definition of catalysts: increase reaction rate without being consumed; lower activation energy.

• Practical: Investigate effects of different solids on reactions/catalysis (e.g., hydrogen peroxide).

(c) Reversible Reactions and Equilibria

• Reversible reactions are indicated by the symbol = in equations.

• Examples: dehydration of hydrated copper(II) sulfate, heating ammonium chloride.

4. Organic Chemistry

(a) Introduction

• Hydrocarbons are composed of hydrogen and carbon.

• Represent organic molecules using empirical, molecular, general, structural, and displayed formulae.

• Terms: homologous series, functional group, isomerism explained.

• Naming rules follow IUPAC nomenclature for compounds up to six carbons.

• Structural and displayed formulae from molecular formula analyses.

• Classification of reactions: substitution, addition, combustion (no mechanisms required).

(b) Crude Oil

• Crude oil as a mixture of hydrocarbons.

• Fractional distillation for separation into fractions.

• Fractions: refinery gases, gasoline, kerosene, diesel, fuel oil, bitumen.

• Trends in color, boiling point, viscosity of fractions.

• Fuels release heat when burned; products of combustion include complete/incomplete combustion gases.

• Carbon monoxide's toxicity explained (not needing hemoglobin details).

• NOx formation in car engines and sulfur dioxide formation from fuel impurities.

• Environmental impact: acid rain contributors.

• Catalytic cracking explained (conditions and purpose).

(c) Alkanes

• General formula for alkanes presented.

• Classification as saturated hydrocarbons.

• Structural and displayed formulae for alkanes (up to C5) and naming unbranched isomers.

• Alkane reactions with halogens (mono-substitution; no mechanisms required).

(d) Alkenes

• Alkenes defined by functional group >C=C<.

• General formula for alkenes specified.

• Classification as unsaturated hydrocarbons.

• Structural and displayed formulae for alkenes (up to C4) and naming unbranched isomers.

• Alkenes react with bromine producing dibromoalkanes.

• Bromine water use for alkane vs. alkene differentiation.

(e) Synthetic Polymers

• Addition polymers created from many monomers.

• Repeat unit structure development for common addition polymers (e.g., poly(ethene), poly(propene), etc.).

• Deduction of monomer structure from repeat unit, and vice versa.

• Issues around addition polymer disposal: inertness, toxic gas production upon burning.