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WASSCE · 27 topics

Chemistry

G3N tutors you through the full WASSCE Chemistry syllabus offline — from Introduction to Chemistry, Scientific Method and Atoms, The Concept of the Moles, Mole Ratios, Chemical Formulae and Chemical Equations and more — with adaptive lessons, instant quizzes and exam-ready summaries.

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Syllabus

What you’ll cover in Chemistry.

The complete topic outline G3N teaches, mapped to the WASSCE curriculum.

Year 1

9 topics
Introduction to Chemistry, Scientific Method and Atoms
  • Describe chemical processes around us and their applications in everyday life
    • Define chemistry as a scientific discipline that studies matter, its composition, structure, properties, and the principles governing its behaviour
    • Identify the branches of pure chemistry: physical chemistry, organic chemistry, and inorganic chemistry
    • Describe applied chemistry and its practical applications in food science, medicine, pharmaceuticals, materials science, agriculture, and environmental science
    • Explain why chemistry is considered the 'central science' due to its connections with physics, biology, and environmental science
    • Describe the impact of chemistry on daily life: food and nutrition, agriculture, medicine, transportation, and energy production
    • Give examples of Ghanaian chemical industries: Tema Oil Refinery (TOR), Ghana Standards Authority, Glofert fertiliser company, Tobinco Pharmaceutical Ltd, Ernest Chemist
    • Identify career opportunities in chemistry and related fields: pharmacist, medical doctor, biochemist, chemical engineer, chemistry teacher, nurse, laboratory technician
    • Describe educational pathways for chemistry careers: pharmaceuticals, environmental science, industrial chemistry, forensic science, chemical engineering, food science
  • Discuss and explain safety rules and hazard symbols in the laboratory
    • State and explain rules and regulations in the chemistry laboratory (no eating, no tasting chemicals, wipe spills, add acid to water, wear PPE, report accidents)
    • Identify and interpret the nine chemical hazard symbols: harmful, irritant, corrosive, toxic, oxidising, flammable, explosive, radioactive, biohazard
    • Describe types of chemical hazards with examples: explosives (dynamite), flammable liquids (ethanol), corrosive substances (HCl, NaOH), toxic substances (arsenic, mercury), oxidising substances (H₂O₂)
    • Identify prohibition signs used in laboratories: no naked flame, danger, no smoking
    • Identify first aid signs: first aid, safety shower, eye wash station
    • Describe types of personal protective equipment (PPE): respirators/gas masks, hand gloves, eye protectors/safety goggles, protective clothing (lab coats, aprons)
    • Describe laboratory safety equipment: eye shower station and fume chamber/hood
    • Analyse lab images to identify good and wrong practices, PPE, glassware, and potential hazards
  • Explain why chemicals should be stored by compatibility and not alphabetically in the laboratory
    • Explain the risks of alphabetical chemical storage: incompatible substances stored adjacent can cause fires, explosions, or toxic fume release
    • Explain the benefits of compatibility-based chemical storage: prevents accidents, organises storage, protects laboratory workers
    • Use a chemical incompatibility chart to identify chemicals that must not be stored together
    • Describe steps to put out a small fire using a fire blanket (stop, drop and roll; cover base of fire; seal edges)
    • Describe steps to use a fire extinguisher (PASS method: Pull alarm, Aim at base, Squeeze handle, Sweep side to side)
  • Investigate the scientific method of inquiry
    • Define the scientific method as a systematic, logical, and repeatable approach to investigating natural phenomena
    • List and describe the seven steps of the scientific method: observation, formulate a question, develop a hypothesis, conduct experiments, collect and analyse data, draw conclusions, communicate results
    • Distinguish between a hypothesis (testable guess), scientific theory (well-established explanation), and scientific law (mathematical relationship between observables)
    • Apply the scientific method to identify and investigate real-world problems in the school environment or community
    • Design an experiment and formulate a hypothesis for an identified problem
  • Identify the main postulates of Dalton's atomic theory and explain the weaknesses of the theory
    • State the four postulates of Dalton's atomic theory: atoms are small indivisible particles; atoms cannot be created or destroyed; atoms of the same element are identical in mass and size; atoms combine in simple whole-number ratios to form compounds
    • Explain the modifications to Dalton's atomic theory: atoms are divisible into subatomic particles; atoms can be created/destroyed in nuclear reactions; isotopes exist with different masses; large organic molecules contain thousands of atoms
    • Define isotopes as atoms of the same element with the same proton number but different numbers of neutrons
    • Define relative atomic mass as the average mass of an atom of an element relative to 1/12 the mass of carbon-12
    • Define relative molecular mass as the sum of relative atomic masses of all atoms in a molecule
    • Describe the mass spectrometer: instrument that measures mass-to-charge ratio (m/z) of ions through vaporisation, ionisation, acceleration, deflection, and detection
    • Identify the five processes in a mass spectrometer: vaporisation, ionisation, acceleration by electric field, deflection by magnetic field, detection and signal generation
    • Interpret a mass spectrum (graph of percentage abundance vs. relative atomic mass) to identify isotopes and calculate relative atomic mass
  • Describe the cathode ray experiment and alpha particle scattering experiment and identify the weaknesses of J. J. Thomson and Rutherford's models of the atom
    • Describe J. J. Thomson's cathode ray experiment: electric current through vacuum tube produces negatively charged particles (cathode rays/electrons) deflected by electric and magnetic fields
    • Describe J. J. Thomson's plum pudding model: atom is a positively charged sphere with electrons embedded throughout like plums in a pudding
    • State the weaknesses of Thomson's model: cannot explain the nucleus, gives no information on electron arrangement, does not explain atomic mass distribution
    • Describe Rutherford's alpha scattering experiment: alpha particles fired at thin gold foil; most pass through, some deflect at large angles, some scatter backward
    • Describe Rutherford's nuclear model: small dense positively charged nucleus at the centre surrounded by orbiting negatively charged electrons
    • State the weaknesses of Rutherford's model: cannot explain atomic stability (electrons should spiral into nucleus), cannot account for emission spectra, cannot explain isotopes
    • Identify the properties of subatomic particles: proton (positive charge, mass 1 amu, in nucleus), neutron (neutral, mass 1 amu, in nucleus), electron (negative charge, mass 1/1840 amu, outside nucleus)
  • State the main postulates of Bohr's planetary theory and explain the importance of the quantum numbers to the electron structure of the atom
    • State the six postulates of Bohr's planetary model: electrons move in fixed circular orbits; only certain orbits are allowed; electrons in an orbit do not radiate energy; energy emitted equals the difference between initial and final energy levels; larger orbits have more energy; transitions between levels produce or absorb photons
    • Define continuous spectrum as electromagnetic radiation containing photons of all energy levels within a range; examples include light bulb filament and stars
    • Define line spectrum as discrete wavelengths of radiation produced when excited atoms emit light; each element has a unique line spectrum
    • Explain the relationship between hydrogen's emission spectrum and electron energy levels: each spectral line corresponds to a specific electron energy level transition
    • Describe contributions of quantum theory: wave-particle duality, discrete energy levels, uncertainty principle, quantum numbers, and electron spin
    • Identify the four quantum numbers: principal quantum number (n, energy level size), angular momentum/azimuthal quantum number (l, subshell shape, 0 to n-1), magnetic quantum number (m, orientation, -l to +l), spin quantum number (s, +½ or -½)
    • Describe the shapes of orbitals: s orbital (spherical, holds 2 electrons), p orbitals (dumbbell-shaped, px py pz, holds 6 electrons), d orbitals (complex shapes, holds 10 electrons)
  • Apply Aufbau's principle, Pauli's exclusion principle and Hund's rule of maximum multiplicity to write the electron configuration of the first thirty elements of the periodic table
    • State Aufbau's principle: electrons fill atomic orbitals in order of increasing energy (1s, 2s, 2p, 3s, 3p, 4s, 3d)
    • State Pauli's exclusion principle: no two electrons in an atom can have the same set of four quantum numbers; two electrons in the same orbital must have opposite spins
    • State Hund's rule of maximum multiplicity: within a subshell, electrons occupy orbitals singly with parallel spins before pairing up
    • Write electron configurations using s, p, d notation (e.g., Carbon: 1s² 2s² 2p²)
    • Write electron configurations of ions (e.g., Al³⁺: 1s² 2s² 2p⁶) and elements with irregular configurations (e.g., Cr: [Ar] 3d⁵ 4s¹, Cu: [Ar] 3d¹⁰ 4s¹)
    • Write electron configurations using the electrons-in-boxes method with arrows indicating spin direction
    • Explain the relative stability of fully filled (e.g., s², p⁶, d¹⁰), half-filled (e.g., p³, d⁵), and partially filled orbitals
  • Describe radioactivity and the properties of radiations and compare isotopes based on their stability as well as their applications in everyday life
    • Define radioactivity as the emission of energy (particles or electromagnetic radiation) from unstable atomic nuclei
    • Distinguish between nuclear reactions and chemical reactions: nature of particles, energy changes, rate of reaction, triggering factors, product stability, and emission of radiation
    • Describe properties of alpha radiation: helium nucleus (2 protons, 2 neutrons), charge +2, mass 4 amu, penetrates only a few centimetres in air, blocked by paper, high ionisation potential
    • Describe properties of beta radiation: high-energy electron, charge -1, negligible mass, penetrates several millimetres of aluminium, lower ionisation potential than alpha
    • Describe properties of gamma radiation: high-energy photons, no charge or mass, penetrates several centimetres of lead, lowest ionisation energy of the three radiations
    • Explain factors affecting nuclear stability: neutron-to-proton ratio and binding energy per nucleon
    • Define half-life as the time for half the undecayed atoms in a sample to decay; calculate half-life from experimental data using the halving method and logarithmic formula (t½ = 0.693/λ)
    • Describe uses of radioisotopes: medical imaging (technetium-99, iodine-131), cancer treatment (iodine-131, strontium-89), industrial sterilisation (cobalt-60), agricultural applications, scientific research (carbon-14 dating), energy production
The Concept of the Moles
  • Explain relative atomic mass and relative molecular mass
    • Define relative atomic mass (Ar) as the average mass of one atom of an element compared to 1/12 the mass of one atom of carbon-12, expressed in atomic mass units (amu)
    • Define atomic mass unit (amu) as equal to 1/12th the mass of a carbon-12 atom (approximately 1.6603 × 10⁻²⁴ g)
    • Describe relative atomic mass as an average that accounts for the natural abundance of all isotopes of an element
    • Define relative molecular mass (Mr) as the sum of the relative atomic masses of all atoms in a molecule
    • Calculate relative molecular masses of compounds such as formic acid (CH₂O₂), acetic acid (C₂H₄O₂), ethanol (C₂H₆O), sucrose (C₁₂H₂₂O₁₁)
    • Calculate relative molecular mass for ions and compounds including NH₃, CH₄, SO₄²⁻, and Na₂CO₃·10H₂O
  • Describe the mole as a unit of the amount of substance
    • Define the mole (mol) as the SI unit of amount of substance
    • State Avogadro's constant as 6.022 × 10²³ particles (atoms, molecules, or ions) per mole
    • Define entities as any distinct atom, molecule, or ion
    • Calculate the number of moles from number of entities using: n = N ÷ L, where N is number of particles and L is Avogadro's constant
    • Calculate the number of entities from a given number of moles: N = n × L
  • Calculate different physical quantities (number of entities, mass and volume) based on the amount of substance
    • Calculate the number of moles from mass using: n = m ÷ M, where m is mass in grams and M is molar mass in g/mol
    • Calculate mass from moles using: m = n × M
    • Calculate molar mass as the mass of one mole of a substance in g/mol (numerically equal to the relative molecular mass)
    • Calculate the volume occupied by a gas at STP (standard temperature and pressure) using: V = n × 22.4 dm³
    • Calculate the number of moles from volume of gas at STP using: n = V ÷ 22.4
    • Perform multi-step calculations combining entities, mass, and volume (e.g., calculate number of molecules and volume from a given mass)
  • Explain the mole concept and its relevance in preparation of standard solutions
    • Define concentration in mol/dm³ (molarity) using: c = n ÷ V, where V is volume in dm³
    • Define concentration in g/dm³ using: c = m ÷ V
    • Calculate the mass of solute dissolved in a given volume to achieve a target concentration
    • Calculate the volume of solution needed to obtain a specific number of moles
    • Convert between concentration units (mol/dm³ and g/dm³) using molar mass
    • List apparatus used to prepare a standard solution from a solid solute: balance, beaker, glass rod, volumetric flask, wash bottle
    • Describe the procedure for preparing a standard solution: weigh solute, dissolve in minimum water, transfer to volumetric flask, make up to the mark, mix thoroughly
    • Identify potential sources of error when preparing standard solutions and how to minimise them
Mole Ratios, Chemical Formulae and Chemical Equations
  • Use IUPAC nomenclature to name inorganic compounds and write their formulae based on the laws of chemical combination
    • Define oxidation number as the number of electrons gained or lost by an atom when forming a compound
    • State the rules for assigning oxidation numbers: zero for elemental state; -2 for oxygen (except peroxide -1 and superoxide -½); +1 for hydrogen bonded to non-metal, -1 bonded to metal; ion oxidation state equals the charge; sum equals total charge on molecule or ion
    • Apply rules for naming binary ionic compounds: name cation first, anion second with '-ide' suffix; indicate oxidation state in Roman numerals for metals with atomic number above 20
    • Apply rules for naming simple acids: use prefix 'hydro', root name of central atom, suffix 'ic', and word 'acid' (e.g., HCl = hydrochloric acid)
    • Apply rules for naming oxoacids: use prefixes 'oxo', 'dioxo', 'trioxo', 'tetraoxo' to indicate oxygen atoms; add root name of central atom; add suffix 'ate' with oxidation state in Roman numerals; add 'acid' (e.g., H₂SO₄ = tetraoxosulphate (VI) acid)
    • Apply rules for naming acid salts: name metal cation first, indicate oxidation state if atomic number above 20, add 'hydrogen', then name the oxoanion (e.g., NaHSO₄ = sodium hydrogen tetraoxosulphate (VI))
    • Apply rules for naming simple non-ionic compounds: name electropositive element first, root name of anion, suffix '-ide' (e.g., HCl = hydrogen chloride)
    • Apply rules for naming molecular compounds where elements form multiple compounds with different oxygen atoms: name electropositive element, indicate oxidation state, name electronegative element (e.g., CO₂ = Carbon (IV) oxide)
  • Determine empirical and molecular formulae and calculate percentage composition of elements in compounds
    • Define empirical formula as the simplest whole number ratio of atoms present in a compound
    • Determine empirical formula by: stating mass or percentage of each element; dividing by relative atomic mass; finding simplest whole number ratio
    • Define molecular formula as the actual number of atoms of each element in the simplest unit of a substance
    • Derive molecular formula from empirical formula using: n = molecular mass ÷ empirical formula mass; then multiply empirical formula subscripts by n
    • Define structural formula as the formula showing how atoms in a molecule are bonded to each other
    • Calculate percentage by mass composition of an element in a compound using: % = (relative atomic mass × number of atoms × 100) ÷ relative molecular mass
    • State the Law of Conservation of Mass: total mass of reactants equals total mass of products in a chemical reaction
    • State the Law of Definite Proportion: a given compound always contains its component elements in fixed mass ratios
  • Write and balance chemical equations for different types of reactions
    • Classify and write combustion reactions: substance reacts with oxygen to give oxides (e.g., CH₄ + 2O₂ → CO₂ + H₂O)
    • Classify and write synthesis reactions: two or more substances combine to form a single product (A + B → AB)
    • Classify and write displacement reactions: one atom or ion replaces another in a reactant (e.g., Mg + CuSO₄ → MgSO₄ + Cu)
    • Classify and write decomposition reactions: a compound breaks into simpler substances under heat, light or catalyst (AB → A + B)
    • Write ionic equations by separating dissolved ionic compounds into free ions; identify and cancel spectator ions to yield the net ionic equation
    • Balance chemical equations using conservation of atoms: count atoms on each side, adjust stoichiometric coefficients, verify balance for all elements
  • Perform calculations involving stoichiometric relationships
    • Define stoichiometry as the relationship between quantities of reactants and products in a chemical reaction
    • Define mole ratio as the ratio of moles of any two substances in a balanced equation, taken from stoichiometric coefficients
    • Calculate the number of entities produced using mole ratio: find moles of known substance (n = N ÷ L), apply mole ratio, then calculate N = n × L
    • Calculate the mass of a product or reactant using mole ratio: find moles of known substance (n = m ÷ M), apply mole ratio, then calculate m = n × M
    • Calculate the concentration of a substance using mole ratio in a titration: find moles of known substance (n = c × V), apply mole ratio, then calculate c = n ÷ V
    • Calculate the volume of a gas produced using mole ratio: find moles of known substance, apply mole ratio, then calculate V = n × Vₘ (where Vₘ = 22.4 dm³/mol at STP)
  • Identify limiting and excess reagents and calculate percentage yield
    • Define limiting reagent as the reactant completely consumed in a reaction that limits the amount of product formed
    • Define excess reagent as the reactant remaining after the limiting reagent is consumed
    • Determine the limiting reagent by: calculating initial moles of each reactant; using mole ratios to find required moles; comparing with available moles — the reactant requiring more moles than available is the limiting reagent
    • Calculate the mass of product formed using the moles of the limiting reagent and the mole ratio between the limiting reagent and the product
    • Define actual yield as the amount of product actually obtained from a reaction in practice
    • Define theoretical yield as the maximum amount of product calculated from the limiting reagent using stoichiometry
    • Calculate percentage yield using: percentage yield = (actual yield ÷ theoretical yield) × 100
    • Explain reasons for percentage yield being less than 100%: reversible reactions, side reactions, incomplete separation of products
Kinetic Theory and the States of Matter
  • Describe the characteristics of solids, relate them to bonding type, and identify uses of diamond, graphite and iodine
    • Describe the characteristics of solids: definite shape and volume, closely packed particles in orderly arrangement, high density, strong forces of attraction/repulsion, low compressibility, generally high melting points
    • Classify types of solids: ionic (e.g. NaCl), molecular (e.g. I2), metallic (e.g. Fe), covalent network/atomic (e.g. SiO2, diamond, graphite)
    • Relate bonding type to physical properties (melting point, hardness, electrical conductivity) for Fe, NaCl, SiO2, I2, diamond, and graphite
    • Describe the molecular structure of diamond (tetrahedral 3D covalent network, very hard, electrical insulator) and graphite (layered hexagonal sheets with delocalised electrons, soft, electrical conductor); relate each structure to its uses
    • State uses of diamond: cutting tools, jewellery, abrasives, drill bits; state uses of graphite: electrodes, pencil leads, lubricant, moderator in nuclear reactors
    • State uses of iodine in everyday life: antiseptic for wounds, treatment of thyroid conditions, water purification
    • Determine the melting point of covalent solids (e.g. benzoic acid, oxalic acid, ethanamide): seal substance in capillary tube, attach to thermometer, heat in oil bath with stirring, record temperature range at which solid melts
  • Describe the characteristics of liquids, distinguish between vapour and gas, and explain the effect of external pressure on boiling point
    • Describe the characteristics and nature of liquids: particles in close random arrangement, fixed volume but takes shape of container, moderate compressibility, moderate density, capable of flow (fluidity) and viscosity
    • Distinguish between a vapour (substance in gaseous state below its critical temperature; can be liquefied by pressure alone) and a gas (above its critical temperature; cannot be liquefied by pressure alone)
    • Explain the concepts of external pressure, saturated vapour pressure, boiling, and evaporation
    • Describe the relationship between vapour pressure and boiling point: a liquid boils when its saturated vapour pressure equals the external (atmospheric) pressure
    • Explain the effect of external pressure on boiling point: increased external pressure raises boiling point (pressure cooker cooks food faster); reduced external pressure lowers boiling point (water boils at lower temperature at high altitude)
    • Perform an experiment to make a liquid boil below its normal boiling point by reducing external pressure (e.g. using a syringe or vacuum pump over warm water)
    • Define viscosity as the resistance of a liquid to flow; compare the viscosity of common liquids (e.g., water, oil, syrup, honey) and explain that higher viscosity is caused by stronger intermolecular forces between molecules
  • Explain the kinetic theory of matter and apply it to distinguish between the properties of solids, liquids and gases
    • State the five postulates of the kinetic theory of matter: (1) matter consists of tiny particles (atoms, ions, or molecules); (2) particles are in constant random motion; (3) particle kinetic energy is proportional to absolute temperature; (4) intermolecular forces between particles are negligible in an ideal gas; (5) collisions between particles and container walls are perfectly elastic
    • Describe properties of solids using the kinetic model: fixed shape and volume, closely packed particles in orderly arrangement, strong intermolecular forces, particles vibrate about fixed positions, expand on heating
    • Describe properties of liquids using the kinetic model: fixed volume but takes shape of container, particles close together in random arrangement, particles move rapidly in all directions, intermolecular forces weaker than solids, expand on heating
    • Describe properties of gases using the kinetic model: no fixed shape or volume, negligible intermolecular forces, particles widely spaced in rapid random motion, easily compressed, exert pressure through particle collisions with container walls
    • Explain change of state processes using kinetic theory: melting (solids gain kinetic energy, attractive forces weaken, particles move freely), freezing (particles lose energy, forces restrict to fixed positions), evaporation (surface particles escape intermolecular forces), boiling (rapid vaporisation throughout bulk liquid at fixed temperature), condensation (gas particles lose energy, form liquid bonds)
    • Describe a method to determine melting point: seal substance in capillary tube, attach to thermometer, heat in oil bath with stirring, record temperature at which first and last crystal melts
    • Define Brownian motion as the random movement of particles suspended in a fluid resulting from collisions with fast-moving fluid molecules
    • Perform an experiment to demonstrate diffusion of solute particles in liquids: place drops of KI and Pb(NO3)2 at opposite ends of a water-filled trough; observe yellow precipitate of PbI2 forming as the ions diffuse towards each other
  • State and perform calculations involving various Gas Laws and analyse graphs based on the laws
    • State Boyle's Law: the volume of a fixed mass of gas at constant temperature is inversely proportional to its pressure; P₁V₁ = P₂V₂
    • Sketch and interpret graphs for Boyle's Law: P vs V gives a hyperbolic curve; P vs 1/V gives a straight line through the origin
    • State Charles' Law: the volume of a fixed mass of gas at constant pressure is directly proportional to its absolute temperature; V₁/T₁ = V₂/T₂
    • Sketch and interpret graphs for Charles' Law: V vs T (in Kelvin) gives a straight line through the origin
    • State Gay-Lussac's Law: for a fixed mass of gas at constant volume, pressure is directly proportional to absolute temperature; P₁/T₁ = P₂/T₂
    • State the Combined Gas Law combining Boyle's, Charles', and Gay-Lussac's laws: P₁V₁/T₁ = P₂V₂/T₂
    • State Avogadro's Law: equal volumes of gases at the same temperature and pressure contain the same number of moles; V/n = constant
    • Perform calculations using each gas law, always converting temperature to Kelvin (K = °C + 273) and ensuring consistent units
  • State Graham's Law of Diffusion/Effusion and Dalton's Law of Partial Pressure and apply them to perform calculations
    • Define diffusion as the net flow of matter from a region of high concentration to low concentration due to random molecular motion
    • Define effusion as the escape of gaseous molecules through a tiny hole into a vacuum or region of lower pressure
    • State Graham's Law of Diffusion: the rate of diffusion of a gas is inversely proportional to the square root of its density or molar mass at constant temperature and pressure; R ∝ 1/√M
    • Apply Graham's Law to two gases: R₁/R₂ = √(M₂/M₁); express in terms of time: t₁/t₂ = √(M₁/M₂); express in terms of volume: V₁/V₂ = √(M₂/M₁)
    • Explain that lighter gases (lower molar mass) diffuse faster than heavier gases; calculate relative rates of diffusion for pairs of gases
    • Describe the experimental demonstration of Graham's Law using NH₃ and HCl in a glass tube: white ring of NH₄Cl forms closer to HCl end because NH₃ (lighter) diffuses faster
    • State Dalton's Law of Partial Pressures: in a mixture of non-reacting gases, the total pressure equals the sum of partial pressures; P_T = P₁ + P₂ + ...
    • Define mole fraction (X) of a component as: X₁ = n₁ ÷ (n₁ + n₂ + ...)
  • Write the ideal gas equation and apply it in simple calculations using different numerical values of R and units of Pressure and Volume
    • State the ideal gas equation: PV = nRT, derived by combining Boyle's, Charles', and Avogadro's laws
    • Identify variables: P = pressure (Pa), V = volume (m³), n = moles (mol), R = gas constant, T = temperature (Kelvin)
    • State the values of R: 8.314 J mol⁻¹ K⁻¹ (SI units); 0.082057 L atm mol⁻¹ K⁻¹; 62.364 L Torr mol⁻¹ K⁻¹
    • Match units of P, V, and T to the chosen value of R (e.g., if R = 0.082057 L atm mol⁻¹ K⁻¹, pressure must be in atm and volume in litres)
    • Apply PV = nRT to calculate unknown quantities: rearrange for P = nRT/V, V = nRT/P, n = PV/RT, T = PV/nR
    • Perform multi-step ideal gas calculations involving molar mass: find moles from mass (n = m/M), substitute into PV = nRT to find the unknown
  • Explain why gases show deviation from ideal behaviour and suggest how the ideal gas equation could be modified
    • State the six assumptions of the ideal gas law: particles have zero volume; no intermolecular forces; large number of molecules in constant random motion; perfectly elastic collisions; collision time is negligible; motion follows Newton's laws
    • Explain that real gases deviate from ideal behaviour at high pressure (particle volumes become significant, PV exceeds ideal value) and low temperature (intermolecular forces increase as kinetic energy decreases, PV falls below ideal value)
    • Explain that deviation increases with: higher relative molecular mass (larger molecules have more significant volume) and greater polarity (stronger intermolecular forces, e.g., NH₃ deviates more than O₂ which deviates more than Ne)
    • State the Van der Waals equation for n moles: (P + an²/V²)(V - nb) = nRT, where a measures strength of intermolecular forces and b is the excluded molar volume
    • Explain the correction terms in the Van der Waals equation: the an²/V² term corrects for intermolecular attraction reducing pressure; the nb term corrects for the actual volume occupied by gas molecules
  • Design and perform experiments to prepare and test for gases (hydrogen, ammonia and carbon dioxide gases)
    • Describe preparation of hydrogen gas: react zinc granules or magnesium ribbon with dilute HCl or H₂SO₄; equation: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g); collect by displacement of water
    • State the test for hydrogen gas: a burning/lighted splint produces a 'pop' sound; dry hydrogen using anhydrous CaCl₂
    • Describe physical and chemical properties of hydrogen: lightest gas, colourless, odourless, insoluble in water; neutral to litmus; burns in oxygen with pop sound to form water; reacts with halogens
    • State uses of hydrogen gas: manufacture of ammonia (Haber process), hydrogenation of fats to make margarine, oxy-hydrogen flame for welding and cutting metals, fuel cells
    • Describe preparation of carbon dioxide gas: react marble chips (CaCO₃) with dilute HCl; equation: CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g); collect by downward displacement of air
    • State the test for carbon dioxide: pass through limewater (saturated Ca(OH)₂ solution) — turns milky; excess CO₂ clears the milkiness; dry using concentrated H₂SO₄ or anhydrous CaCl₂
    • Describe physical and chemical properties of carbon dioxide: colourless, odourless, denser than air, soluble in water, condenses at low temperatures into dry ice (solid CO₂); turns moist blue litmus red; does not support combustion; undergoes photosynthesis
    • State uses of carbon dioxide: fire extinguishers, fizzy drinks, photosynthesis, refrigerants
Inter Atomic Bonding
  • Explain ionic bonding and its formation and state the properties of ionic compounds
    • Define ionic bond as an electrostatic force of attraction between a positively charged cation and a negatively charged anion, formed by complete electron transfer from a less electronegative atom to a more electronegative atom
    • Use Lewis dot symbols (element symbol with dots representing valence electrons) to represent atoms and explain ionic bond formation
    • Explain cation formation (atom loses electrons, typically a metal, becomes positively charged) and anion formation (atom gains electrons, typically a non-metal, becomes negatively charged)
    • Describe ionic bond formation with examples: Na donates one electron to F to form NaF; Li donates one electron to F to form LiF
    • Identify three factors that affect ionic bond formation: (1) low ionisation energy of the metallic element forming the cation; (2) large electron affinity of the non-metallic element forming the anion; (3) large lattice energy (related to smaller ion size and higher ion charge)
    • State the properties of ionic compounds: high melting and boiling points; some dissolve in polar solvents like water; conduct electricity in aqueous form or molten state (ions free to move); are good insulators in solid form (ions cannot move); are hard and brittle
    • Explain polarisation of an anion: distortion of the symmetrical electron cloud of an anion by a nearby cation; the ability of a cation to polarise an anion is called its polarising power
    • Identify factors affecting the polarising power of a cation: magnitude of positive charge (directly proportional, e.g., Na+ < Mg2+ < Al3+) and size of cation (inversely proportional, e.g., Li+ > Na+ > K+ > Rb+)
  • Explain covalent bonding and its formation and state the properties of covalent compounds
    • Define covalent bond as a chemical bond formed when two atoms mutually share a pair of electrons, typically between non-metal atoms, to achieve a stable duplet or octet electron configuration
    • Explain that covalent bonding involves overlapping of atomic orbitals each containing one electron from each atom, resulting in a shared pair of electrons
    • Describe pure (non-polar) covalent bonds: electrons shared equally between identical atoms (e.g., H2, Cl2, F2); electrons have an equal probability of being near each nucleus
    • Explain lone pairs as pairs of valence electrons not involved in bonding that remain on individual atoms (e.g., F2 has three lone pairs on each fluorine atom)
    • Describe dative (coordinate) covalent bonds: both electrons in the shared pair are donated by one atom; distinguish from typical covalent bonds where each atom contributes one electron
    • Describe formation of the ammonium ion (NH4+): NH3 + H+ → NH4+; nitrogen's lone pair is donated to a proton (H+) which carries its positive charge to give NH4+
    • Describe formation of the hydroxonium ion (H3O+): H2O + H+ → H3O+; oxygen's lone pair is donated to a proton to produce the positively charged hydroxonium ion
    • Describe polar covalent bonds: electrons shared unequally due to electronegativity difference; the more electronegative atom acquires a partial negative charge (delta-) and the less electronegative atom acquires a partial positive charge (delta+); e.g., H-Cl bond in hydrogen chloride
  • Explain metallic bonding and state the properties of metals
    • Define metallic bond as an electrostatic force of attraction between fixed positive metal ions and a sea of delocalised (mobile) electrons surrounding them in a regular lattice structure
    • Explain metallic bond formation: metal atoms with low electronegativity release their outermost electrons to form a delocalised electron sea; positive metal ions are held together in the lattice by attraction to the surrounding electron cloud
    • Identify factors that determine metallic bond strength: number of valence electrons (more electrons → stronger bond, e.g., Al > Mg > Na), size of atoms (smaller atoms → closer proximity → stronger bond), and proximity of atoms in the lattice
    • Explain malleability: metals can be hammered or pressed into thin sheets without breaking because layers of ions can slide past each other while the delocalised electrons maintain the bonding; examples include aluminium foil and gold leaf
    • Explain ductility: metals can be stretched into thin wires because the electron sea holds the lattice together as ions are rearranged; example includes copper used for electrical wiring
    • Explain that metals are good conductors of heat: delocalised electrons can transfer kinetic energy quickly through the lattice, causing the metal to heat up rapidly and evenly
    • Explain that metals are good conductors of electricity: delocalised electrons move freely throughout the lattice when a potential difference is applied, allowing current to flow; silver and copper are excellent electrical conductors
Intermolecular Bonding
  • Describe the different types of intermolecular forces and explain how they arise from the structural features of molecules
    • Define intermolecular forces (IMF) as the attractive and repulsive forces that arise between molecules of a substance, controlling how molecules interact with each other
    • Distinguish intermolecular forces (between molecules) from intramolecular forces (forces holding atoms together within a molecule, e.g. covalent bonds); covalent bonds (50-200 kJ/mol) are much stronger than IMFs (1-12 kJ/mol)
    • Identify the three main types of Van der Waals forces: dipole-dipole interactions, hydrogen bonds, and induced dipole-induced dipole (London dispersion) forces
    • Explain dipole-dipole interactions: occur between polar molecules that have a permanent separation of positive and negative charges; the positive end of one polar molecule is attracted to the negative end of another polar molecule
    • Explain hydrogen bonds: a special, stronger type of dipole-dipole interaction that occurs when a hydrogen atom is covalently bonded to a small, highly electronegative atom (O, N, or F); the partially positive H is attracted to a partially negative electronegative atom on a nearby molecule
    • Rank the relative strength of intermolecular forces: London dispersion < dipole-dipole < hydrogen bonds; all are weaker than covalent bonds
    • List examples of compounds that form hydrogen bonds: water (H2O), ammonia (NH3), hydrogen fluoride (HF), methanol (CH3OH), ethanol (C2H5OH), acetic acid (CH3COOH), and biomolecules such as DNA and proteins
    • Explain London dispersion (induced dipole-induced dipole) forces: occur between non-polar covalent molecules; temporary uneven electron distribution creates a transient dipole in one molecule that induces a dipole in a neighbouring molecule, causing weak mutual attraction
  • Explain how intermolecular forces affect the physical properties of compounds
    • Explain how IMF strength determines boiling and melting points: substances with stronger IMFs require more energy to overcome attractions and therefore have higher melting and boiling points (e.g., water with hydrogen bonds has a high boiling point); substances with weaker IMFs have lower boiling and melting points (e.g., methane with only London dispersion forces has a very low boiling point)
    • Apply the rule 'like dissolves like' to solubility: polar substances dissolve in polar solvents because similar IMFs operate between solute and solvent; non-polar substances dissolve in non-polar solvents; hydrogen bonding allows substances like ethanol to dissolve readily in water
    • Note that temperature and pressure also affect solubility (e.g., gases dissolve better at higher pressures)
    • Explain how IMF strength determines surface tension: stronger IMFs produce higher surface tension; the surface of a liquid behaves like a stretched elastic membrane because cohesive IMFs pull surface molecules inward; increasing temperature lowers surface tension by giving molecules more kinetic energy
    • Explain how IMF strength determines enthalpy of vaporisation: substances with stronger IMFs need more energy to vaporise and therefore have a higher enthalpy of vaporisation; noble gases with weak London dispersion forces have very low enthalpies of vaporisation
    • Explain how IMF strength determines viscosity: liquids with stronger IMFs have higher viscosity (thicker, flow less easily) because molecules are held together more strongly (e.g., honey flows slowly due to extensive hydrogen bonding); liquids with weaker IMFs flow easily
    • Explain how IMF strength determines volatility: substances with stronger IMFs have lower volatility (do not evaporate easily) because molecules are tightly held; substances with weaker IMFs have high volatility and evaporate easily
Periodic Properties
  • Use the electron configuration of elements to determine their position on the periodic table
    • Define periodicity as the repeating trends in chemical and physical properties of elements arranged in order of increasing atomic number in the periodic table
    • Classify elements into s-block (Groups 1 and 2, valence electrons in s-subshell), p-block (Groups 13-18, valence electrons in p-subshell), and d-block (Groups 3-12, valence electrons in d-subshell) based on the subshell containing their valence electrons
    • Classify elements using the IUPAC group numbering system: Groups 1-18 numbered according to the number of valence electrons (Group 1 = 1 valence electron, Group 2 = 2, Groups 13-18 have 3-8 valence electrons)
    • Classify elements using the Roman numeral system: Groups I-VIII with A/B designations indicating the subshell of the last electron; use electron configuration to assign group and period
    • Determine the period of an element from its electron configuration: the period number equals the highest principal quantum number present; the group number is determined by the number of valence electrons
    • Classify elements as metals (shiny, malleable, ductile, good conductors, form cations), non-metals (dull, brittle, poor conductors, form anions), or semi-metals/metalloids (intermediate properties, e.g., boron, silicon, germanium) based on properties and periodic table position
    • Describe physical properties of Group 1 alkali metals: relatively soft, low density, low melting and boiling points compared to most metals, solid at room temperature
    • Describe physical properties of Group 2 alkaline earth metals: harder than alkali metals, higher density and melting/boiling points than Group 1, all solid at room temperature
  • Explain how periodic properties change with atomic number and principal quantum number
    • State the Periodic Law: the properties of elements are periodic functions of their atomic numbers; elements display repeating trends when arranged by increasing atomic number
    • Define the five key periodic properties: atomic radius (half the distance between nuclei of two covalently bonded atoms), ionic radius (size of an ion in a crystal lattice), ionisation energy (energy to remove one mole of electrons from gaseous atoms), electron affinity (energy change when an electron is added to a neutral gaseous atom), and electronegativity (ability of an atom to attract bonding electrons)
    • Explain the four factors affecting periodic properties: nuclear charge (greater nuclear charge pulls electrons closer); distance from nucleus (greater shell number means valence electrons are farther from the nucleus and less attracted); shielding effect (inner electron shells shield valence electrons from full nuclear attraction, reducing effective nuclear charge); electron configuration (filled or half-filled subshells confer extra stability)
    • Describe the trend in atomic radius across a period: decreases from left to right because increasing nuclear charge attracts electrons more strongly, drawing the electron cloud inward and making atoms more compact
    • Describe the trend in atomic radius down a group: increases because each successive period adds a new electron shell (higher principal quantum number), increasing the distance from the nucleus and the shielding effect
    • Describe the trend in ionic radius: cations are smaller than their parent neutral atoms because loss of electrons reduces electron-electron repulsion and contracts the electron cloud; anions are larger than their parent atoms because gain of electrons increases repulsion and expands the electron cloud
    • Describe the trend in ionisation energy across a period: generally increases from left to right due to increasing nuclear charge, which binds electrons more tightly to the nucleus and requires more energy to remove them
    • Describe the trend in ionisation energy down a group: decreases because increasing atomic size and greater electron shielding by inner shells reduces the effective nuclear charge on valence electrons, making them easier to remove
Qualitative and Quantitative Analysis of Organic Compounds
  • Describe the qualitative and quantitative analysis of organic compounds
    • Define qualitative analysis as identifying which elements (C, H, O, N, S, halogens) are present in an organic compound
    • Define quantitative analysis as determining the exact mass composition of elements in an organic compound, typically expressed as percentages
    • Describe distillation as a separation technique for isolating a volatile liquid from non-volatile components: heat mixture until volatile component boils and vaporises, pass vapour through a condenser, collect the liquid distillate; requires significant boiling point differences between components
    • State practical applications of distillation: purifying seawater, separating ethanol from water, purifying essential oils from plants, producing distilled water for laboratory use
    • Describe fractional distillation for separating two or more miscible liquids with similar boiling points: heat mixture so all components vaporise, vapour passes through a fractionating column where components condense at different heights according to their boiling points, each fraction is tapped off separately from most to least volatile
    • State practical applications of fractional distillation: separating methanol-ethanol mixtures, distilling crude oil into useful fractions, isolating oxygen and other gases from liquid air
    • Describe crystallisation as a purification technique for solids: dissolve solid in minimum hot solvent, allow to cool so pure crystals form as the substance becomes supersaturated, filter and dry the crystals in a desiccator; impurities remain in solution at lower concentration
    • State substances purified by crystallisation: benzoic acid, glucose, paracetamol
  • Design an experiment to test for the presence and mass composition of carbon, hydrogen, sulphur, nitrogen and halogens in organic compounds
    • Describe the test for carbon and hydrogen in organic compounds: heat organic sample strongly with copper(II) oxide (oxidising agent); carbon is oxidised to CO2 which turns limewater milky; hydrogen is oxidised to water which turns anhydrous copper(II) sulphate from white to blue
    • Describe Lassaigne's sodium fusion test: fuse organic compound with sodium metal in a fusion tube to convert nitrogen, sulphur, and halogens into water-soluble inorganic ions (sodium cyanide NaCN, sodium sulphide Na2S, and sodium halide NaX); cool and extract with water to form the test filtrate
    • Describe the test for sulphur using the sodium fusion filtrate: add a few drops of freshly prepared sodium nitroprusside solution; formation of a purple or violet colour confirms the presence of sulphur
    • Describe the test for nitrogen using the sodium fusion filtrate: add dilute NaOH, heat and cool the solution, add iron(II) sulphate solution followed by drops of iron(III) chloride solution; formation of a Prussian blue precipitate confirms the presence of nitrogen
    • Describe the test for halogens using the sodium fusion filtrate: acidify with excess concentrated nitric acid and heat to remove sulphides and cyanides; cool and add aqueous silver nitrate (AgNO3); white precipitate (AgCl, soluble in ammonia) confirms chlorine; cream precipitate (AgBr, insoluble in dilute ammonia) confirms bromine; yellow precipitate (AgI, insoluble in concentrated ammonia) confirms iodine
    • Calculate the mass and percentage of carbon from combustion analysis: mass of C = (12/44) × mass of CO2 produced; % C = (mass of C ÷ mass of organic compound) × 100
    • Calculate the mass and percentage of hydrogen from combustion analysis: mass of H = (2/18) × mass of H2O produced; % H = (mass of H ÷ mass of organic compound) × 100
    • Describe the Carius method for estimating halogens: heat a known mass of compound with concentrated HNO3 and AgNO3 in a sealed Carius tube; C and H are oxidised to CO2 and H2O; the halogen forms AgX precipitate which is filtered, dried, and weighed; % X = (atomic mass of X ÷ molecular mass of AgX) × (mass of AgX ÷ mass of compound) × 100
Classifications of Organic Compounds
  • Distinguish between organic and inorganic compounds and classify organic compounds
    • Define organic chemistry as the branch of chemistry that studies the structure and properties of carbon compounds, excluding oxides of carbon (CO, CO2), carbonates, carbides, and cyanides
    • Define organic compounds as compounds containing carbon atoms covalently bonded to other atoms
    • State the properties of carbon that enable formation of many stable compounds: carbon is tetravalent (forms four covalent bonds); carbon atoms can join to form straight chains, branched chains, and ring structures; carbon can form single, double, or triple bonds with itself and other elements; carbon can form isomers (compounds with the same molecular formula but different structural arrangements)
    • Compare organic and inorganic compounds: organic compounds must contain carbon, are usually covalent, have relatively low melting and boiling points, are usually soluble in non-polar solvents, and often exist as liquids or gases at room temperature; inorganic compounds may contain any element except organic carbon, are often ionic, have relatively high melting and boiling points, are often soluble in polar solvents, and usually occur as solids
    • Classify organic compounds into four main types: aliphatic hydrocarbons (open straight or branched chains with C-C, C=C, or C≡C bonds, e.g., alkanes, alkenes, alkynes); alicyclic hydrocarbons (closed rings of carbon atoms with no benzene ring, e.g., cyclopropane); aromatic hydrocarbons (contain one or more benzene rings, e.g., benzene, toluene); heterocyclic compounds (rings containing carbon plus other atoms such as O, N, or S, e.g., pyridine, furan, thiophene)
    • Define hydrocarbons as organic compounds containing only carbon and hydrogen atoms; classify as saturated hydrocarbons (only C-C single bonds, e.g., alkanes) or unsaturated hydrocarbons (containing C=C double bonds or C≡C triple bonds, e.g., alkenes and alkynes)
    • Use IUPAC prefixes to name hydrocarbons by carbon chain length: meth- (1C), eth- (2C), prop- (3C), but- (4C), pent- (5C), hex- (6C), hept- (7C), oct- (8C), non- (9C), dec- (10C) up to eicosan- (20C)
    • Represent organic compounds using three formula types: molecular formula (e.g., C4H10), condensed structural formula (e.g., CH3CH2CH2CH3), and full structural formula showing all bonds and atoms explicitly
  • Explain homologous series and state their properties
    • Define a homologous series as a group of organic compounds having the same general molecular formula, the same functional group, and similar chemical properties, with each successive member differing by a -CH2- unit
    • State the five properties of a homologous series: all members share a general molecular formula; all members have the same general method of preparation; all members exhibit similar chemical properties; all members contain the same functional group; physical properties (e.g., boiling point, melting point, density) change gradually and predictably along the series
    • Write the general formula for alkanes as CnH(2n+2) and list the first five members with molecular and condensed structural formulae: methane CH4, ethane C2H6 (CH3CH3), propane C3H8 (CH3CH2CH3), butane C4H10 (CH3CH2CH2CH3), pentane C5H12
    • Write the general formula for alkenes as CnH(2n) and list members starting from ethene C2H4 (CH2=CH2); note that alkenes cannot have a single-carbon member because the C=C double bond requires at least two carbon atoms; the double bond position can vary in larger alkenes (e.g., but-1-ene and but-2-ene)
    • Explain the trend in physical properties along a homologous series: boiling points and melting points increase with increasing chain length because larger molecules have stronger Van der Waals (dispersion) forces; lower members are gases, middle members are liquids, and higher members are solids at room temperature
    • Explain that members of a homologous series have similar chemical properties because they share the same functional group, while physical properties vary systematically due to the increasing molecular size

Year 2

8 topics
Energy Changes
  • Explain enthalpy change and distinguish between exothermic and endothermic reactions
    • Define enthalpy (H) as a thermodynamic quantity describing the energy of a system: H = U + PV
    • Define enthalpy change (ΔH) as the difference in enthalpy between products and reactants: ΔH = H(products) − H(reactants)
    • Distinguish between open, closed, and isolated systems in terms of energy and matter exchange
    • Explain exothermic reactions as reactions that release heat to surroundings (ΔH negative, products have lower enthalpy than reactants)
    • Explain endothermic reactions as reactions that absorb heat from surroundings (ΔH positive, products have higher enthalpy than reactants)
    • Draw and interpret energy profile diagrams for exothermic and endothermic reactions, labelling reactants, products, activation energy (Ea), and ΔH
  • Describe and calculate the different types of standard enthalpy changes
    • Define standard conditions as 298 K (25°C) and 1 atm (101.3 kPa) with standard concentration of 1.0 mol/dm³
    • Define and calculate standard enthalpy change of reaction (ΔH°rxn): heat change when stoichiometric molar quantities of reactants combine under standard conditions
    • Define and calculate standard enthalpy change of formation (ΔH°f): enthalpy change when one mole of a compound is formed from its elements under standard conditions; ΔH°f of elements in standard states is zero
    • Define and calculate standard enthalpy change of combustion (ΔH°c): enthalpy change when one mole of a substance is completely combusted in oxygen under standard conditions
    • Define and calculate standard enthalpy change of neutralisation (ΔH°n): enthalpy change when one mole of water is formed from acid-base reaction under standard conditions
    • Define and calculate standard enthalpy change of solution (ΔH°soln): enthalpy change when one mole of ionic substance is dissolved in excess water under standard conditions
    • Define standard enthalpy of hydration (ΔH°hyd): enthalpy change when one mole of gaseous ions is dissolved in water to form aqueous ions
    • Use the relationship ΔH°rxn = ΣnΔH°f(products) − ΣmΔH°f(reactants) to calculate standard enthalpy changes
  • Experimentally determine enthalpy changes using calorimetry
    • Describe the experimental determination of enthalpy change of combustion of alcohols using a calorimeter
    • Calculate heat transferred using Q = m × c × ΔT, where c is specific heat capacity of water (4.18 J/g°C)
    • Calculate enthalpy of combustion per gram and per mole from experimental data
    • Describe the experimental determination of standard enthalpy of neutralisation for a strong acid-base reaction
    • Describe the experimental determination of standard enthalpy of solution for a dissolving solute
    • Evaluate sources of experimental error and calculate percentage difference from theoretical values
  • Apply Hess's Law of constant heat summation to calculate enthalpy changes
    • State Hess's Law: the total enthalpy change of a chemical reaction is equal to the sum of all individual enthalpy changes, regardless of the pathway
    • Apply rules for manipulating thermochemical equations: reversing the sign when reversing an equation; multiplying enthalpy by the same factor when scaling coefficients; adding equations and their corresponding ΔH values
    • Use Hess's Law to calculate enthalpy changes for reactions that cannot be measured directly by calorimetry
    • Draw and interpret Born-Haber cycles to calculate lattice energy of ionic compounds
    • Identify the energy steps in a Born-Haber cycle: sublimation energy, bond dissociation energy, ionisation energy, electron affinity, enthalpy of formation, and lattice energy
    • Apply Born-Haber cycles using Hess's Law to calculate unknown energy values for ionic compounds (e.g., NaCl, LiF)
  • Explain bond energy and use it to estimate enthalpy changes
    • Define bond enthalpy as the average energy needed to break one mole of a specific covalent bond in a gaseous molecule
    • Distinguish between bond dissociation energy (energy to break a specific bond) and bond enthalpy (average value across molecules)
    • Define enthalpy of atomisation as the enthalpy change when all bonds in one mole of a substance are broken to produce gaseous atoms
    • Use bond enthalpy data to calculate enthalpy of reaction: ΔH°rxn = ΣΔH°(bonds broken) − ΣΔH°(bonds formed)
    • Explain that bond breaking is endothermic and bond forming is exothermic
    • Compare bond enthalpies of single, double, and triple bonds (e.g., C–C, C=C, C≡C) and relate to bond strength
Chemical Kinetics
  • Measure and express the rate of a chemical reaction
    • Define rate of reaction as the change in concentration (or moles/mass) of a reactant or product per unit time: Rate = Δ[A]/Δt
    • Distinguish between average rate (over a time interval), initial rate (at time t = 0), and instantaneous rate (at a specific moment)
    • Write rate expressions for reactions using stoichiometric coefficients to relate rates of different species
    • Calculate average rates of reaction from concentration-time data and interpret concentration-time graphs
    • Calculate instantaneous rate from the gradient (tangent) of a concentration-time curve
  • Explain the factors that affect the rate of a chemical reaction
    • Explain how temperature affects reaction rate: increased temperature increases kinetic energy of particles, leading to more frequent and energetic collisions
    • Explain how concentration affects reaction rate: higher concentration increases the number of particles per unit volume, increasing collision frequency
    • Explain how surface area affects reaction rate: greater surface area exposes more particles to reactants, increasing collision frequency (relevant when one reactant is solid)
    • Explain how a catalyst increases reaction rate: provides an alternative lower-activation-energy pathway without being consumed in the reaction
    • Explain the role of pressure for gaseous reactions: increased pressure increases concentration of gas molecules, increasing collision frequency
  • Apply collision theory to explain reaction rates
    • State collision theory: reactions occur when reactant particles collide with sufficient energy (activation energy) and correct orientation
    • Define activation energy (Ea) as the minimum energy required for a reaction to occur
    • Use Maxwell-Boltzmann energy distribution curves to explain the effect of temperature on reaction rate
    • Explain how increasing temperature shifts the Maxwell-Boltzmann curve: the peak lowers and shifts right, increasing the proportion of particles with energy greater than or equal to Ea
    • Explain how a catalyst lowers the activation energy, increasing the proportion of particles with sufficient energy for effective collisions
    • Draw Maxwell-Boltzmann distribution curves for reactions with and without a catalyst, showing lower Ea
  • Derive and apply rate equations and determine the order of reactions
    • Write rate equations in the form: Rate = k[A]x[B]y, where k is the rate constant and x and y are orders with respect to each reactant
    • Define order of reaction with respect to a reactant as the exponent to which its concentration term is raised in the rate law
    • Define overall order of reaction as the sum of all individual orders (x + y)
    • Describe and interpret zero-order reactions: rate is constant and independent of reactant concentration; rate = k
    • Describe and interpret first-order reactions: rate is directly proportional to concentration of one reactant; rate = k[A]
    • Describe and interpret second-order reactions: rate is proportional to the square of one reactant's concentration or to the product of two concentrations
    • Determine orders of reaction from experimental data by comparing initial rates when concentrations are varied
    • Calculate the rate constant k (with appropriate units) from rate equation and experimental data
Dynamic Equilibrium
  • Distinguish between reversible and irreversible reactions and describe dynamic equilibrium
    • Distinguish between irreversible reactions (single arrow) and reversible reactions (double arrow)
    • Define dynamic equilibrium as the state where forward and reverse reactions occur at equal rates and concentrations of reactants and products remain constant
    • Explain that at equilibrium the concentrations of reactants and products are constant but not necessarily equal
    • Use the example of nitrogen dioxide and dinitrogen tetroxide (NO2 to N2O4) to illustrate dynamic equilibrium
    • Describe the experiment with anhydrous copper sulphate to demonstrate reversible reactions
  • Write and calculate equilibrium constant expressions (Kc and Kp)
    • Write the equilibrium constant expression Kc for homogeneous reactions in terms of molar concentrations
    • Explain that pure solids and liquids are excluded from the Kc expression (heterogeneous equilibria)
    • Write the equilibrium constant expression Kp for gaseous reactions in terms of partial pressures
    • Derive the relationship between Kc and Kp: Kp = Kc(RT) raised to the power of delta-n, where delta-n is the change in moles of gas
    • Calculate Kc and Kp values from equilibrium concentration and pressure data
    • Calculate unknown equilibrium concentrations given Kc and initial concentrations
    • Interpret Kc values: large Kc (products favoured), small Kc (reactants favoured), Kc = 1 (equal concentrations)
  • Apply Le Chatelier's Principle to predict the effect of changes on equilibrium position
    • State Le Chatelier's Principle: when a system at equilibrium is disturbed, it shifts to counteract the change and re-establish equilibrium
    • Predict the effect of changing concentration on equilibrium position: adding a reactant or removing a product shifts equilibrium forward
    • Predict the effect of changing pressure/volume on gaseous equilibria: increased pressure shifts equilibrium toward the side with fewer moles of gas
    • Predict the effect of changing temperature on equilibrium: increasing temperature favours the endothermic direction
    • Explain that a catalyst does not shift the equilibrium position but speeds up both forward and reverse reactions equally
    • Apply Le Chatelier's Principle to explain why Kc changes with temperature but not with concentration or pressure
  • Apply concepts of equilibrium and rates to industrial chemical processes
    • Explain the conditions used in the Haber process for ammonia synthesis (N2 + 3H2 to 2NH3): temperature ~450 degrees Celsius, pressure ~200 atm, iron catalyst, and the compromise between rate and yield
    • Explain the conditions used in the Contact process for sulphuric acid production (2SO2 + O2 to 2SO3): vanadium(V) oxide catalyst, ~450 degrees Celsius
    • Explain how catalytic converters in cars convert harmful gases (CO, NOx) using Le Chatelier's Principle and high temperature
    • Describe catalytic cracking and reforming in petroleum refining and how conditions are optimised
    • Explain the use of ion exchange in water treatment to soften hard water
Acids, Bases and Salts
  • Explain and compare the Arrhenius, Bronsted-Lowry, and Lewis concepts of acids and bases
    • Define Arrhenius acids as substances that produce hydrogen ions (H+) in aqueous solution and Arrhenius bases as substances that produce hydroxide ions (OH-) in aqueous solution
    • State the limitation of the Arrhenius theory: applies only to aqueous solutions
    • Define Bronsted-Lowry acids as proton (H+) donors and Bronsted-Lowry bases as proton acceptors
    • Identify conjugate acid-base pairs in Bronsted-Lowry reactions (two species that differ by one proton)
    • Explain amphoteric substances (e.g., water, Al(OH)3) that can act as both Bronsted-Lowry acids and bases
    • Define Lewis acids as electron-pair acceptors and Lewis bases as electron-pair donors
    • Compare the three acid-base theories and explain which is the most general (Lewis > Bronsted-Lowry > Arrhenius)
  • Describe the physical and chemical properties of acids and bases
    • State the physical properties of acids: sour taste, turns blue litmus red, turns methyl orange/yellow to pink, pH less than 7, conducts electricity, turns phenolphthalein colourless
    • State the physical properties of bases: bitter taste, turns red litmus blue, slippery/soapy feel, pH greater than 7, conducts electricity in solution, turns phenolphthalein pink
    • Describe chemical properties of acids: reaction with metals above hydrogen in reactivity series to produce hydrogen gas; reaction with bases (neutralisation) to form salts and water; reaction with metal carbonates to produce CO2
    • Describe chemical properties of bases: reaction with acids (neutralisation); reaction with ammonium salts to produce ammonia gas; reaction with some metals
  • Distinguish between strong and weak acids and bases and calculate pH
    • Define strong acids as acids that fully dissociate in water (e.g., HCl, H2SO4, HNO3) and weak acids as acids that partially dissociate (e.g., CH3COOH, H2CO3)
    • Define strong bases as bases that fully dissociate in water (e.g., NaOH, KOH) and weak bases as bases that partially dissociate (e.g., NH3)
    • Explain the differences in properties between strong and weak acids/bases: conductivity, reactivity with metals, enthalpy of neutralisation
    • Define pH as pH = -log[H+] and use it to measure acidity
    • Calculate pH of strong acid and strong base solutions from concentration
    • Explain acid dissociation constant (Ka) as a measure of the strength of a weak acid
  • Perform and interpret acid-base titrations including back titrations
    • Name and describe the function of apparatus used in titration: burette, pipette, pipette filler, conical flask, retort stand and clamp, indicators
    • Describe the procedure for a simple acid-base titration: fill burette with titrant (known concentration), add indicator to analyte (unknown concentration), add titrant until endpoint (colour change)
    • Identify appropriate indicators for different acid-base combinations: phenolphthalein for strong base/weak acid, methyl orange for strong acid/weak base
    • Calculate the concentration of an unknown solution using C1V1 = C2V2 and stoichiometric ratios
    • Describe back titration as a method used when a direct titration is not possible (e.g., insoluble solids, volatile substances)
    • Calculate unknown quantities using back titration data: moles of excess reagent = moles added minus moles reacted back
    • Solve titration problems involving molar mass determination of unknown salts and acids
Trends of Chemical and Physical Properties of Elements and Their Compounds in the Periodic Table
  • Describe trends in physical and chemical properties of Period 3 elements
    • State the Modern Periodic Law: properties of elements repeat in a regular pattern based on atomic number and electron arrangement
    • List Period 3 elements: Na, Mg, Al, Si, P, S, Cl, Ar, each with outer electrons in the third energy level
    • Describe the trend in metallic character across Period 3: decreases from Na (metal) to Cl (non-metal); Si is a metalloid; Ar is a noble gas
    • Describe the trend in ionisation energy across Period 3: generally increases from Na to Ar (with exceptions at Al and S due to subshell effects)
    • Describe the trend in reactivity with water across Period 3: Na reacts vigorously, Mg slowly, Al forms protective oxide layer, Si and P do not react, S and Cl react slightly
    • Describe the trend in reactivity with oxygen across Period 3: metals form basic oxides, non-metals form acidic oxides
    • Describe the trend in reactivity with chlorine across Period 3: all elements except Ar form chlorides
  • Describe and explain the physical and chemical properties of hydrides of Period 3 elements
    • List Period 3 hydrides: NaH, MgH2 (ionic), AlH3 (partly ionic/covalent), SiH4, PH3, H2S, HCl (covalent)
    • Compare melting and boiling points: ionic hydrides (NaH, MgH2) have high melting points due to strong ionic bonds; covalent hydrides have lower melting points due to weak Van der Waals forces
    • Compare solubility and reactivity with water: ionic hydrides react to form hydrogen gas and hydroxide; covalent hydrides are generally soluble forming acidic or neutral solutions
    • Compare thermal stability: ionic hydrides are more thermally stable; covalent hydrides decompose at lower temperatures
    • Describe the acid-base behaviour of Period 3 hydrides: hydrides become more acidic across the period
  • Describe and explain the physical and chemical properties of oxides of Period 3 elements
    • List Period 3 oxides: Na2O, MgO (ionic), Al2O3 (amphoteric), SiO2 (giant covalent), P4O10, SO3, Cl2O7 (molecular)
    • Compare the structures: ionic oxides (strong ionic bonds), giant covalent oxides (network of covalent bonds), molecular oxides (small molecules with weak Van der Waals forces)
    • Explain trends in melting/boiling points: high for ionic and giant covalent; low for molecular oxides
    • Explain electrical conductivity: ionic oxides conduct when molten; covalent and molecular oxides do not conduct
    • Classify oxides as basic (Na2O, MgO), amphoteric (Al2O3), or acidic (SiO2, P4O10, SO3) based on reaction with acids and bases
    • Write equations for reactions of Period 3 oxides with water: basic oxides form hydroxides (e.g., Na2O + H2O gives 2NaOH); acidic oxides form acids (e.g., SO3 + H2O gives H2SO4)
  • Describe and explain the physical and chemical properties of hydroxides and chlorides of Period 3 elements
    • List Period 3 hydroxides: NaOH, Mg(OH)2, Al(OH)3, Si(OH)4, P(OH)3
    • Describe the trend in acid-base character of Period 3 hydroxides: from strongly basic (NaOH) to amphoteric (Al(OH)3) to strongly acidic (HClO4)
    • List Period 3 chlorides: NaCl, MgCl2 (ionic), AlCl3 (intermediate), SiCl4, PCl3, PCl5, SCl2 (covalent)
    • Compare the structure and bonding of Period 3 chlorides: ionic chlorides (Na, Mg) vs. covalent chlorides (Si-Cl)
    • Describe the hydrolysis of Period 3 chlorides with water: ionic chlorides dissolve to form neutral/basic solutions; covalent chlorides hydrolyse to form acidic solutions and HCl gas
Physical and Chemical Properties of the Halogens
  • Describe the physical properties and trends of the halogens (Group 17)
    • Identify halogens as Group 17 elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At)
    • Describe physical states at room temperature: F2 and Cl2 are gases, Br2 is a liquid, I2 is a solid, At is a radioactive solid
    • Explain why halogens are reactive: they have ns2np5 configuration and need one electron to achieve a full outer shell
    • Describe trends down Group 17: colour darkens (pale yellow to purple-black), density increases, melting/boiling points increase, intermolecular forces strengthen, oxidising strength decreases
    • Describe and explain the trend in bond energy of X-X bonds: decreases down the group due to increasing atomic size and lone pair repulsion
    • Describe the solubility of halogens in water and organic solvents
  • Describe the chemical properties of halogens and their reactions
    • Describe reactions of halogens with metals to form ionic metal halides (e.g., Na + 0.5Cl2 gives NaCl)
    • Describe reactions of halogens with hydrogen to form hydrogen halides (e.g., H2 + Cl2 gives 2HCl)
    • Describe reactions of halogens with water and alkalis
    • Explain halogen displacement reactions: a more reactive halogen displaces a less reactive halide ion from its salt solution (e.g., Cl2 + 2KBr gives 2KCl + Br2)
    • Use displacement reactions to determine the order of oxidising power: F2 > Cl2 > Br2 > I2
    • Describe the use of silver nitrate solution to identify halide ions: Cl- forms white AgCl precipitate; Br- forms cream AgBr; I- forms yellow AgI
  • Describe the reactions of halide salts, compare acid strengths of hydrogen halides, and identify uses of halogens
    • Describe reactions of solid halide salts with concentrated H2SO4: NaCl/NaBr/NaI produce HCl/HBr/HI respectively; HBr reduces H2SO4 to SO2; HI reduces H2SO4 further to S and H2S
    • Identify the relative reducing power of halide ions: F- < Cl- < Br- < I- (iodide is the strongest reducing agent)
    • Explain the trend in acid strength of hydrogen halides: HF << HCl < HBr < HI; HF is a weak acid (strong H-F bond), HCl, HBr, HI are strong acids
    • Relate H-X bond strength to acid strength: weaker bond leads to easier dissociation and stronger acid
    • Describe uses of halogens: chlorine in water purification and bleaching, fluorine in toothpaste and non-stick coatings (PTFE), iodine as antiseptic, bromine in flame retardants
Structure, Chemical Bonding and Properties of Molecular Compounds
  • Predict the shape and bond angles of molecular compounds using VSEPR theory and explain sigma and pi bonds
    • Define electronegativity using the Pauling scale and explain trends: increases across a period, decreases down a group
    • Classify bonds by electronegativity difference: non-polar covalent (less than 0.5), polar covalent (0.5 to 1.7), ionic (greater than 1.7)
    • Define dipole moment and explain how bond polarity determines molecular polarity
    • Apply VSEPR theory to predict molecular geometry based on the number of bonding pairs and lone pairs around the central atom
    • Predict and state bond angles for common molecular shapes: linear (180 degrees), trigonal planar (120 degrees), tetrahedral (109.5 degrees), trigonal pyramidal (~107 degrees), bent/V-shaped (~104.5 degrees)
    • Describe shapes of specific molecules: CH4 (tetrahedral), CO2 (linear), NH3 (trigonal pyramidal), H2O (bent), BF3 (trigonal planar), PCl5 (trigonal bipyramidal), SF6 (octahedral)
    • Define sigma bond as formed by head-on overlap of orbitals, allowing free rotation
    • Define pi bond as formed by side-by-side overlap of p orbitals, restricting rotation
  • Explain hybridisation and describe bonding in key organic and inorganic molecules
    • Define hybridisation as the mixing of atomic orbitals to form new hybrid orbitals of equivalent energy
    • Describe sp3 hybridisation: mixing of one s and three p orbitals to form four equivalent sp3 orbitals in tetrahedral arrangement (109.5 degrees); examples include CH4 and C2H6
    • Describe sp2 hybridisation: mixing of one s and two p orbitals to form three sp2 orbitals in trigonal planar arrangement (120 degrees), with one unhybridised p orbital forming a pi bond; examples include C2H4 (ethene) and benzene
    • Describe sp hybridisation: mixing of one s and one p orbital to form two sp orbitals in linear arrangement (180 degrees), with two unhybridised p orbitals forming two pi bonds; examples include C2H2 (ethyne) and CO2
    • Relate hybridisation to molecular geometry and bond angles in organic compounds: sp3 (single bonds), sp2 (double bonds), sp (triple bonds)
    • Explain the bonding in benzene: each carbon is sp2 hybridised, forming a flat hexagonal ring with delocalised pi electrons above and below the ring, giving resonance stability
Organic Compounds
  • Describe the structure, naming, isomerism, and properties of alkanes
    • Define alkanes as saturated hydrocarbons with only C-C single bonds; general formula CnH(2n+2)
    • Name straight-chain alkanes using IUPAC nomenclature: methane, ethane, propane, butane, pentane, hexane, heptane, octane, nonane, decane
    • Name branched-chain alkanes by identifying the longest chain, numbering from the end nearest a substituent, and naming substituents (methyl-, ethyl-) with locants
    • Describe structural isomerism in alkanes: chain isomerism where same molecular formula has different carbon chain arrangements
    • Describe physical properties of alkanes: low polarity, low boiling points increasing with chain length, less dense than water, insoluble in water
    • Describe chemical properties of alkanes: combustion (complete combustion gives CO2 + H2O; incomplete gives CO + soot), free radical substitution (halogenation) with UV light
    • Explain the mechanism of free radical substitution: initiation (homolytic bond cleavage by UV light), propagation (chain reaction), termination (radical combination)
  • Describe the structure, naming, isomerism, and properties of alkenes
    • Define alkenes as unsaturated hydrocarbons with at least one C=C double bond; general formula CnH(2n)
    • Name alkenes using IUPAC nomenclature with suffix '-ene' and a locant for the double bond position
    • Describe structural isomerism in alkenes: chain isomerism and position isomerism (double bond in different positions)
    • Describe geometric (cis-trans) isomerism in alkenes: arises from restricted rotation around C=C double bond; cis (same side) and trans (opposite side) arrangements
    • Describe physical properties of alkenes: colourless gases (up to 4 carbons) or volatile liquids, insoluble in water, less dense than water
    • Describe chemical properties of alkenes: addition reactions with H2 (hydrogenation), HX (hydrogen halides), H2O (hydration with H2SO4 catalyst), X2 (halogenation); oxidation with KMnO4
    • Explain Markovnikov's rule for addition of HX to unsymmetrical alkenes: H adds to the carbon with more hydrogen atoms
    • Describe the test for alkenes: bromine water decolourises in the presence of an alkene (electrophilic addition)
  • Describe the structure, naming, and properties of alkynes
    • Define alkynes as unsaturated hydrocarbons with at least one C triple bond C; general formula CnH(2n-2)
    • Name alkynes using IUPAC nomenclature with suffix '-yne' (e.g., ethyne, propyne)
    • Describe the structure of alkynes: sp hybridised carbon atoms, linear geometry, bond angle of 180 degrees
    • Explain the triple bond in alkynes: one sigma bond (head-on overlap of sp orbitals) and two pi bonds (side-by-side overlap of unhybridised p orbitals)
    • Describe physical properties of alkynes: non-polar, insoluble in water, soluble in organic solvents
    • Describe chemical reactions of alkynes: addition reactions with H2, HX, and X2; combustion
  • Describe the structure, bonding, and properties of benzene
    • Describe the historical development of benzene's structure: Kekule's alternating double bond model and its limitations (equal bond lengths, unexpected stability)
    • Explain the delocalised structure of benzene: each carbon is sp2 hybridised, forming a flat hexagonal ring with pi electrons delocalised above and below the ring
    • Explain resonance in benzene: the two Kekule structures are resonance hybrids; electrons are not fixed but spread over all six carbons
    • Explain the unique stability of benzene due to delocalisation (aromatic stability)
    • Describe electrophilic substitution reactions of benzene: nitration (with HNO3/H2SO4 forming nitrobenzene) and halogenation
    • Write the equation for the complete combustion of benzene: 2C6H6 + 15O2 gives 12CO2 + 6H2O
  • Describe the structure, naming, classification, and properties of alkanols
    • Define alkanols as organic compounds containing one or more hydroxyl (-OH) groups; general suffix '-anol'
    • Classify alkanols as primary (-OH on carbon bonded to one other carbon), secondary (-OH on carbon bonded to two carbons), or tertiary (-OH on carbon bonded to three carbons)
    • Name alkanols using IUPAC rules: identify the longest chain containing the -OH group, replace '-ane' with '-anol', number from end nearest to -OH
    • Describe physical properties of alkanols: lower members are liquids with relatively high boiling points due to hydrogen bonding, miscible with water, density less than water
    • Describe chemical properties of alkanols: combustion, oxidation (primary to aldehyde to carboxylic acid; secondary to ketone; tertiary is resistant to oxidation), dehydration to form alkenes, esterification with carboxylic acids
  • Describe the structure, naming, and properties of alkanoic acids
    • Define alkanoic acids as organic compounds containing a carboxyl group (-COOH); general formula CnH(2n+1)COOH
    • Name alkanoic acids using IUPAC rules: replace the final '-e' of the parent alkane with '-oic acid'; the carboxyl carbon is always carbon number 1
    • List the first four alkanoic acids: methanoic acid (HCOOH), ethanoic acid (CH3COOH), propanoic acid (C2H5COOH), butanoic acid (C3H7COOH)
    • Describe physical properties of alkanoic acids: weak acids (partially dissociate in water), hydrogen bonding causes higher boiling points, lower members are miscible with water
    • Describe chemical properties of alkanoic acids: reaction with metals to form salts and hydrogen gas; reaction with bases (neutralisation) to form salts and water; reaction with carbonates to produce CO2; esterification with alcohols to form esters and water
    • Explain the effect of electron-withdrawing substituents (e.g., halogens) on acid strength: stabilise the conjugate base, increasing acidity

Year 3

10 topics
Chemical Industry
  • Explain the concept of chemical industry and identify chemical plants operating in Ghana
    • Define 'industry' as the large-scale production of goods or services; define 'chemical industry' as industries that convert raw materials into chemical products through chemical processes
    • Explain what a chemical plant is and distinguish it from a chemical industry
    • Identify chemical plants in Ghana and describe what they produce: crude oil refinery, cement production, soap making, salt making, gold refinery, steel production, aluminium production, brewing
    • Explore raw materials used by each plant and state whether they are local or foreign; explain why some Ghanaian plants depend on foreign raw materials
    • Discuss the major products and by-products of the chemical plants listed
    • Describe basic chemical processes used to transform raw materials into products: saponification, electrolysis, solvent extraction, fermentation, precipitation, decomposition
Extraction of Metals
  • Identify mineral deposits in Ghana and outline the extraction and economic importance of key metals
    • Describe properties and reactivity series of metals (review from Year 2)
    • Define mineral/ore as a naturally occurring solid material from which a useful metal or mineral can be profitably extracted
    • Identify mineral deposits and their locations in Ghana: metallic minerals (gold, bauxite, manganese, iron ores); precious stones (diamond); industrial minerals (limestone, clay, kaolin, solar salt, crude oil)
    • Identify the ores of gold (quartz/pyrite), aluminium (bauxite), iron (haematite, magnetite), and manganese (pyrolusite)
    • Outline the extraction of gold from quartz and aluminium from bauxite; write chemical equations for key steps in extraction processes
    • Discuss the domestic and industrial uses of aluminium and gold and their alloys in Ghana
    • Explain why duralumin (an aluminium alloy) is preferred to pure aluminium in many industrial applications
Extraction of Crude OIL and Petroleum Processing
  • Describe the formation, composition, extraction, and processing of crude oil
    • Explain the theory of crude oil formation from zooplankton, algae, diatoms, foraminifera, and radiolaria over millions of years under heat and pressure
    • Describe the chemical composition of crude oil: mainly hydrocarbons (aliphatic and aromatic), with impurities containing N, O, S
    • Classify crude oil by API gravity (light vs heavy), geographic location, and sulphur content (sweet vs sour); explain why light and sweet crude oils are preferred and command higher prices
    • Describe crude oil extraction from an oil well: primary recovery (natural gas pressure drives oil to surface) and secondary recovery (gas lift and gas injection methods)
    • Describe the fractional distillation of crude oil and the fractions obtained (refinery gas, petrol/gasoline, naphtha, kerosene, diesel, fuel oil, bitumen) and their uses
    • Explain cracking (breaking large hydrocarbon molecules into smaller, more useful ones) and reforming (rearranging molecules to improve quality); state their importance to the petroleum refining industry
    • Define petrochemicals as chemicals derived from crude oil or natural gas used as raw materials in industry; give examples and their uses
    • Define octane number as a measure of a fuel's resistance to knocking; explain the use of anti-knocking agents; explain why lead tetraethyl (TEL) has been phased out
Environmental Pollution
  • Describe the sources, types, and effects of air, water, and land pollution
    • Define pollution as the introduction of harmful substances into the environment; identify systems that can be polluted: air, water, land
    • Describe natural causes of air pollution: wildfires (increase CO2 and particulate matter), windblown dust and wind storms in the West African sub-region
    • Describe human activities that cause air pollution: combustion of fossil fuels (producing CO2, CO, soot, SO2 from sulphur compounds, NO2 from nitrogen impurities); roasting metal sulphides (producing SO2)
    • Describe gold mining pollution: release of arsenic oxide (As2O3) into the atmosphere
    • Explain acid rain: SO2 and NO2 react with water in the atmosphere to form H2SO3 and HNO3, which fall as acid rain and damage vegetation, aquatic life, and structures
    • Explain the greenhouse effect: CO2 and other greenhouse gases (CH4, H2O vapour) trap infrared radiation near Earth's surface, causing global warming
    • Explain ozone depletion: CFCs and other chemicals break down ozone (O3) in the stratosphere; reduced ozone allows more UV radiation to reach Earth, causing skin cancer and crop damage
    • Describe the consequences of acid rain, greenhouse effect, and ozone depletion on vegetation, water bodies, and human wellbeing
Biotechnology
  • Describe the concept, products, and services of biotechnology and relate it to traditional indigenous technology
    • Define biotechnology as the use of living organisms or biological processes to develop products and services; describe its applications in food and drink, waste treatment, genetic engineering, pharmaceuticals, mining, and fuel production
    • Describe biotechnology products: chemicals (food additives, ethanol, polymers); fuel (biogas/methane from anaerobic digestion; gasohol, a blend of ethanol and gasoline)
    • Describe biotechnology services: bioleaching (using microorganisms to extract metals from low-grade ores); treatment of oil spills using hydrocarbon-degrading bacteria; domestic waste treatment and sewage treatment
    • Describe the production process for a traditional indigenous product (e.g. gari, kenkey, or local gin/akpeteshie): raw materials, stages in production using a flow chart, source of energy, chemical transformations, products and by-products, waste management, and health/safety considerations
Cement and its Uses
  • Describe cement production from raw materials, its uses, and its environmental impact
    • Identify raw materials for clinker production: limestone (CaCO3) and clay (aluminosilicates)
    • Describe the physical treatment of raw materials: crushing and grinding of limestone; drying and milling of clay; mixing of powdered limestone and clay
    • Describe the chemical treatment: roasting the limestone-clay mixture in a rotary kiln at approximately 1600°C to produce clinker (mainly calcium silicates)
    • Explain cement production from clinker: mix clinker with gypsum (to control setting time), then mill and pack for sale
    • State uses of cement in the construction industry: bridges, roads, buildings, dams; note that asbestos (an older cement-based material) is highly carcinogenic
    • Describe the environmental pollution associated with cement production: cement dust released into the atmosphere, skin and respiratory irritation on contact with cement dust, heat dissipation from kilns, noise pollution from machinery
Fats and Oils
  • Describe the sources, structure, properties, and uses of fats and oils including soap production
    • Identify sources of fats and oils: plant sources (palm oil, coconut oil, groundnut oil, soya bean oil) and animal sources (lard, butter, fish oil)
    • Describe the physical properties of fats and oils: insoluble in water, soluble in organic solvents; oils are liquid at room temperature (unsaturated fats) while fats are solid (more saturated)
    • Describe the general structure of fats and oils as mono-, di-, and tri-esters of propan-1,2,3-triol (glycerol) with long-chain fatty acid groups
    • Describe acidic and alkaline hydrolysis of fats and oils; describe hydrogenation of unsaturated oils to produce solid fats (margarine production)
    • Perform the test for fats and oils (e.g. translucent grease spot on paper, or emulsification test)
    • Outline saponification: reaction of fats or oils with NaOH or KOH to produce soap (sodium or potassium salt of fatty acid) and glycerol; apply using local materials (cocoa husks or plantain peel as alkali source)
    • Distinguish soapy detergents from soapless (synthetic) detergents in terms of raw materials, structure, and effectiveness in hard water
    • State uses of fats and oils: food preparation, lubrication of machine parts, medicinal use; body functions (energy source, thermal insulation, cell membrane components)
Proteins
  • Describe the sources, structure, and biological functions of proteins as natural polymers
    • Identify sources of dietary protein: meat, fish, eggs, beans, milk, groundnuts, soya beans, cashew nuts
    • Describe the solubility of proteins in water; explain that hydrolysis of proteins yields amino acids
    • Describe the general structure of alpha-amino acids: contain an amino group (-NH2) and a carboxyl group (-COOH) on the same alpha-carbon; use glycine and tyrosine as specific examples
    • Explain that approximately 20 common alpha-amino acids (obtained from organisms' proteins) are the building blocks of all proteins
    • Describe protein formation: condensation reaction between amino acids forms peptide (amide) bonds, producing polypeptide chains that fold into proteins
    • Classify proteins as natural condensation polymers of amino acid monomers
    • State uses and biological functions of proteins: dietary food, structural material (keratin, collagen), biological catalysts (enzymes), hormones (insulin), genetic material (histones), energy source
Carbohydrates
  • Describe the sources, classification, and functions of carbohydrates as natural polymers
    • Identify sources of carbohydrates: sugars (sucrose, glucose), cereals (rice, maize, wheat), tubers (cassava, yam), fruits, honey
    • Describe the solubility of sugars in water; explain that hydrolysis of disaccharides and polysaccharides produces monosaccharides
    • Perform tests for reducing sugars: use Fehling's solution or Benedict's solution (brick-red precipitate indicates reducing sugar); use Tollens' reagent (silver mirror test); use glucose test strips
    • Classify carbohydrates: monosaccharides (glucose, galactose, fructose — simplest sugars, cannot be hydrolysed further); disaccharides (lactose from milk, sucrose from cane/beet sugar — hydrolysed to two monosaccharides); polysaccharides (starch, cellulose, glycogen — hydrolysed to many monosaccharides)
    • Describe starch as a natural addition polymer of glucose units; explain the hydrolysis of polysaccharides to monosaccharides
    • State uses of carbohydrates: industrial (paper, textiles, fermentation), commercial (sweeteners, food additives), medicinal (glucose drips, pharmaceutical excipients), domestic (food), pharmaceutical
    • Describe the functions of polysaccharides in the body: starch/glucose as primary energy source; glycogen as short-term energy reserve in liver and muscles; cellulose as dietary fibre in plant-based foods
Synthetic Polymers
  • Describe the formation, classification, properties, and uses of synthetic polymers
    • Define polymer as a large molecule built from many repeating monomer units joined by covalent bonds; define polymerization as the chemical process of forming polymers
    • Classify synthetic polymers by monomer type: polystyrene (from styrene), synthetic rubber (from isoprene), polythene (from ethene), PTFE/Teflon (from tetrafluoroethene), orlon (from acrylonitrile), nylon (from diamine and diacid), bakelite (from phenol and methanal)
    • Describe addition polymerization: unsaturated monomers (alkenes) open their double bonds and link together without any by-product; examples: polythene from ethene, PVC (polyvinyl chloride) from vinyl chloride/chloroethene
    • Describe condensation polymerization: monomers react and join with loss of a small molecule (e.g. water); example: polyurethane formed from diisocyanate and diol monomers; nylon from diamine and diacid
    • Differentiate thermoplastics (soften and can be reshaped on heating, e.g. polythene, PVC) from thermosets (harden permanently on heating and cannot be remoulded, e.g. bakelite, phenol-methanal plastics)
    • Describe chemical tests on plastics: behaviour on heating (melts or chars), reaction with acids, reaction with alkalis
    • Discuss advantages of synthetic polymers: high strength, low density, electrical insulation, chemical resistance, low cost, versatility; discuss disadvantages: poor biodegradability, disposal problems, use of non-renewable resources
    • State uses of synthetic polymers as replacements for metals, glass, ceramics, wood, cardboard, and paper; examples: polythene for packaging, PVC for pipes and electrical cable insulation, polyurethane foams for thermal insulation, Teflon for non-stick cookware
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