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2024/2025 WAEC Chemistry Syllabus – Latest waec syllabus for chemistry

The West African Examination Council (WAEC) 2024 WAEC Chemistry Syllabus, the Latest WAEC syllabus for Chemistry Exam, is out.

We are pleased to inform all WAEC students participating in the WASSCE exam that the WAEC official has released the latest WAEC 2024 chemistry syllabus.

The 2024 WAEC syllabus for the chemistry exam contains all the recommended topics you need to read and study against your WAEC exam.

2024 WAEC Chemistry Syllabus – The latest WAEC syllabus for chemistry.


(a)(i) Measurement of physical quantities.

(ii) Scientific measurements and their importance in chemistry.

(b) Scientific Methods


(a) Gross features of the atom.

(b) (i) Atomic number/proton number, number of neutrons, isotopes, atomic mass, mass number.

(1) Measurement of mass, length, time, temperature, and volume.

(2) Appropriate SI units and significant figures.

(3) Precision and accuracy in measurement.

Outline the scientific method to include:

Observation, hypothesis, experimentation, formulation of laws and theories.

(1) Short account of Dalton’s atomic theory and limitations, J.J. Thompson’s experiment, and Bohr’s model of the atom.

(2) Outline the description of Rutherford’s alpha scattering experiment to establish the structure of the atom.

Meaning and representation in symbols of atoms and sub-atomic particles.

CONTENT NOTES  (ii) Relative atomic mass (Ar) and relative molecular mass (Mr) based on the Carbon-12 scale.(iii) Characteristics and

nature of matter.

(c) Particulate nature of matter: physical and chemical changes.

(d) (i) Electron Configuration

(ii) Orbitals

(iii) Rules and principles

for filling in electrons.

(1) Atomic mass as the weighted average mass of isotopes. Calculation of the relative mass of chlorine should be used as an example. (2) Carbon-12 scale as a unit of measurement.

Definition of atomic mass unit.

Atoms, molecules, and ions.

Definition of particles and treatment of particles as building blocks of matter.

Explain physical and chemical changes with examples.

Physical change- melting of solids, magnetizing iron, dissolution of salt, etc.

Chemical change- burning of wood, rusting of iron, decay of leaves, etc.

Detailed electron configurations (s,p,d) for atoms of the first thirty elements.

Origin of s,p, and d orbitals as sub-energy levels; shapes of s and p orbitals only.

(1) Aufbau Principle, Hund’s Rule of Maximum Multiplicity, and Pauli Exclusion Principle.

(2) Abbreviated and detailed electron configuration regarding s, p, and d.


(b) Separation techniques

(c) Criteria for purity.


(a) Periodicity of the elements.

(b) Different categories of elements in the periodic table.

(c) Periodic law:

(i) Trends on the periodic table;

(ii) Periodic gradation of the elements in the third period (Na – Ar).

Solid-solid, solid-liquid, liquid-liquid, gas-gas with examples.

Crystallization, distillation, precipitation, magnetization, chromatography, sublimation, etc.

Boiling point for liquids and melting point for solids.

Electron configurations lead to group and periodic classifications.

Metals, semi-metals, and non-metals are in the periodic table, as well as halogens. Alkali metals, alkaline earth metals, and transition metals are metals.

Explanation of the periodic law.

Periodic properties: atomic size, ionic size, ionization energy, electron affinity, and electronegativity.

Simple discrepancies regarding beryllium, boron, oxygen, and nitrogen should be accounted for.

(1) Progression from:

(i) metallic to the non-metallic character of the element;

(ii) ionic to covalent bonding in compounds.

CONTENTS NOTES(d) Reactions between acids and metals, their oxides and trioxocarbonates (IV).

(e) Periodic gradation of elements in group seven, the halogens: F, Cl, Br, and I.

(f) Elements of the first transition series.

21Sc – 30Zn

(2) Differences and similarities in the properties between the second and third-period elements should be stated. (1) Period three metals (Na, Mg, Al).

(2) Period four metals (K, Ca).

(3) Chemical equations.

(4) pH of solutions of the metallic oxides and trioxocarbonates.

Recognition of group variations, noting any anomalies.

Treatment should include the following:

(a) physical states, melting and boiling points;

(b) variable oxidation states;

(c) redox properties of the elements;

(d) displacement reaction of one halogen by another;

(e) the response of the elements with water and alkali (balanced equations required).

(1) Their electron configurations, physical properties, and chemical reactivity of the elements and their compounds.

(2) Physical properties should include physical states and metallic and magnetic properties.

(3) Reactivity of the metals with air, water, and acids, and comparison with s-block elements (Li, Na, Be, Mg).


(a) Interatomic bonding

(b) (i) Formation of ionic bonds and compounds.

(ii) Properties of ionic compounds.

(c) Naming of ionic compounds.

(d) Formation of covalent bonds and compounds.

(e) (i) Properties of covalent compounds.

(ii) Coordinate (dative) covalent bonding.

(4) Other properties of transition metals should include (a) variable oxidation states;

(b) formation of colored compounds;

(c) complex formation;

(d) catalytic abilities;

(e) paramagnetism;

(f) hardness.

I was meaning of chemical bonding.

Lewis dot structure for simple ionic and covalent compounds.

Formation of stable compounds from ions. Factors influencing formation: ionization energy, electron affinity, and electronegativity difference.

Solubility in polar and non-polar solvents, electrical conductivity, hardness, and melting point.

IUPAC system for simple ionic compounds.

Factors influencing covalent bond formation. Electron affinity, ionization energy, atomic size, and electronegativity.

Solubility in polar and non-polar solvents, melting point, boiling point, and electrical conductivity.

Formation and difference between pure covalent and coordinate (dative) covalent bonds.

CONTENT NOTES (f) Shapes of molecular compounds. (g) (i) Metallic Bonding

(ii) Factors influencing its formation.

(iii) Properties of metals.

(h) (i) Intermolecular bonding

(ii) Intermolecular forces in covalent compounds.

(iii) Hydrogen bonding

(iii) van der Waals forces

(iv) Comparison of all bond types.


Linear, planar, tetrahedral, and shapes for some compounds, e.g., BeCl2, BF3, CH4, NH3, CO2. Factors should include atomic radius, ionization energy, and number of valence electrons. Types of specific packing are not required.

Typical properties include heat and electrical conductivity, malleability, luster, ductility, sonority, and hardness.

Relative physical properties of polar and non-polar compounds.

Description of formation and nature should be treated.

Dipole-dipole, induced dipole-dipole, and induced dipole-induced dipole forces should be treated under van der Waal’s forces.

Variation of the melting points and boiling points of noble gases, halogens, and alkanes in the homologous series explained in terms of van der Waal’s forces, and variation in the boiling points of H2O and H2S explained using Hydrogen bonding.

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6.0 STOICHIOMETRY AND CHEMICAL REACTIONS(a) (i) Symbols, formulae, and equations.

(ii) chemical symbols

(iii) Empirical and molecular formulae.

(iv) Chemical equations and IUPAC names of chemical compounds.

(v) Laws of chemical combination.

(b) Amount of substance.


Symbols of the first thirty elements and other common elements that are not among the first thirty elements.

Calculations involving formulae and equations will be required. Mass and volume relationships in chemical reactions and the stoichiometry of such responses include the calculation of the percentage composition of an element.

(1) Combustion reactions (including combustion of simple hydrocarbons)

(2) Synthesis

(3) Displacement or replacement

(4) Decomposition

(5) Ionic reactions

(1) Laws of conservation of mass.

(2) Law of constant composition.

(3) Law of multiple proportions. Explanation of the laws to balance given equations.

(4) Experimental illustration of the law of conservation of mass.

(1) Mass and volume measurements.

(2) The mole as a unit of measurement; Avogadro’s constant, L= 6.02 x 1023 entities mol-1.

(3) Molar quantities and their uses.

(4) Moles of electrons, atoms, molecules, formula units, etc.


(c) Mole ratios(d) (i) Solutions

(ii) Concentration terms

(iii) Standard solutions.

(e) Preparation of solutions from liquid solutes using the dilution method.


Use of mole ratios in determining stoichiometry of chemical reactions. Simple calculations determine the number of entities, amount of substance, mass, concentration, volume, and percentage yield of product. (1) Concept of a solution as made up of solvent and solute.

(2) Distinguishing between dilute solution and concentrated solution.

(3) Basic, acidic, and neutral solutions.

Mass (g) or moles (mol) per unit volume. Emphasis on current IUPAC chemical terminology, symbols, and conventions. Concentration can be expressed as mass concentration, g dm-3, molar concentration, mol dm-3.

(1) Preparation of some primary standards e.g anhydrous Na2CO3, (COOH)2, 2H2O/H2C2O4.2H2O.

(2) Meaning of the terms primary standard, secondary standard, and standard solution.

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Dilution factor


7.0 STATES OF MATTER(a) (i) Kinetic theory of matter.

(ii) Changes of state of matter.

(iii) Diffusion


(1) Postulates of the kinetic theory of matter.

(2) Use the kinetic theory to explain the processes: melting of solids, boiling of liquids, evaporation of liquids, dissolution of solutes, Brownian motion, and diffusion.

(1) Changes in the state of matter should be explained in terms of the movement of particles. It should be emphasized that randomness decreases (and orderliness increases) from a gaseous state to a liquid state and a solid state and vice versa.

(2) Illustrations of changes of state using the different forms of water, iodine, sulfur, naphthalene, etc.

(3) Brownian motion to be illustrated using any of the following experiments:

(a) pollen grains/powdered sulfur in water (viewed under a microscope);

(b) smoke in a glass container illuminated by an intense light from the side;

(c) a dusty room being swept and viewed from outside under sunlight.

(1) Experimental demonstration of diffusion of two gases.

(2) Relationship between the speed at which different gas particles move and the masses of particles.

(3) Experimental demonstration of diffusion of solute particles in liquids.


(b) Gases:(i) Characteristics and nature of gases;

(ii) The gas laws;

(iii) Laboratory preparation and properties of some gases.

(c) (i) Liquids

(ii) Vapour and gases.


Arrangement of particles, density, shape, and compressibility.

The Gas laws: Charles’, Boyle’s, Dalton’s law of partial pressure, Graham’s law of diffusion, and Avogadro’s law. The ideal gas equation of state. A qualitative explanation of each of the gas laws was provided using the kinetic model.

The use of Kinetic molecular theory to explain changes in gas volumes, pressure, and temperature.

Mathematical relations of the gas law


Ideal and Real gases

Factors responsible for the deviation of natural gases from a perfect situation.

(1) Preparation of the following gases: H2, NH3 and CO2. Principles of purification and collection of gases.

(2) Physical and chemical properties of the gases.

Characteristics and nature of liquids based on the arrangement of particles, shape, volume, compressibility, density, and viscosity.

(1) Concept of vapor, vapor pressure, saturated vapor pressure, boiling, and evaporation.

(2) Distinction between vapor and gas.

(3) Effect of vapor pressure on boiling points of liquids.

(4) Boiling at reduced pressure.


(d) Solids:(i) Characteristics and nature;

(ii) Types and structures;

(iii) Properties of solids.

(e) Structures, properties, and uses of diamond and graphite.

(f) Determination of melting points of covalent solids.


(a) Energy and enthalpy

(b) Description, definition, and illustrations of energy changes and their effects.


(1) Ionic, metallic, covalent network, and molecular solids. Examples in each case.

(2) Arrangements of particles, ions, molecules, and atoms in the solid state.

Relate the properties of solids to the interatomic and intermolecular bonding in the solids—identification of the types of chemical bonds in graphite and differences in the physical properties.

The uses of diamond and graphite are related to the structure.

The use of iodine in everyday life.

Melting points indicate the purchase city of solids e.,g, ph, enyl methanedioic acid (benzoic acid), ethanedioic acid (oxalic), and ethanamide.

Explanation of the terms energy and enthalpy. Energy changes associated with chemical processes.

(1) Exothermic and endothermic processes.

(2) Total energy of a system as the sum of various forms of energy, e.g., kinetic, potential, electrical, heat, sound, etc.

(3) Enthalpy changes are involved in the following processes: combustion, dissolution, and neutralization.


9.0 ACIDS, BASES AND SALTS(a) Definitions of acids and bases.

(b) Physical and chemical properties of acids and bases.

(c) Acids, bases, and salts as electrolytes.

(d) Classification of acids and bases.

(e) Concept of pH

(1) Arrhenius concepts of acids and bases regarding H3O+ and OH– ions in water.

(2) Effects of acids and bases on indicators, metal Zn, Fe, and trioxocarbonate (IV) salts and hydrogentrioxocarbonate (IV) salts.

Characteristic properties of acids and bases in aqueous solution include:

(a) conductivities, taste, litmus/indicators, feel, etc.;

(b) balanced chemical equations of all reactions.

Electrolytes and non-electrolytes; strong and weak electrolytes. Evidence from conductivity and enthalpy of neutralization.

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(1) Strength of acids and bases.

(2) Classify acids and bases into strong and weak.

(3) Extent of dissociation reaction with water and conductivity.

(4) Behaviour of weak acids and weak bases in water as examples of equilibrium systems.

(1) Definition of pH and knowledge of pH scale.

(2) Measure the pH of solutions using a pH meter, calometric methods, or universal indicator.

(3) Significance of pH values in everyday life, e.g., acid rain, soil pH, blood, urine.

CONTENT NOTES (f) Salts:(i) Laboratory and industrial preparation of salts;

(ii) Uses;

(iii) Hydrolysis of salt.

(g) Deliquescent, efflorescent and hygroscopic compound.

(h) Acid-Base indicators

(i) Acid-base titration


I was meaning of salts. Types of salts: standard, acidic, basic, double, and complex salts.

(1) Description of laboratory and industrial production of salts.

(2) Mining of impure sodium chloride and conversion into coarse salt.

(3) Preparation of NaOH, Cl2 and H2.

(1) Explain how salts form aciform alkaline and neu, trial aqueous solutions.

(2) Behaviour of some salts (e.g NH4Cl, AlCl3, Na2CO3, CH3COONa) in water as examples of equilibrium systems.

(3) Effects of charge density of some cations and anions on the hydrolysis of their aqueous solution. Examples are to be taken from group 1, group 2, group 3, and the d-block element.

The use of hygroscopic compounds as a drying agent should be emphasized.

(1) Qualitative description of how acid-base indicator works.

(2) Indicators as weak organic acids or bases (organic dyes).

(3) The indicator’s color at any pH depends on the relative amounts of acid and its forms.

(4) Working pH ranges of methyl orange and phenolphthalein.

(1) Knowledge and correct use of relevant apparatus.

(2) Knowledge of how acid-base indicators work in titrations.



(a) General principles

(b) Practical application of solubility.

(3) Acid-base titration experiments involving HCl, HNO3, H2SO4 and NaOH, KOH, Ca(OH)2, CO32-, HCO3–.(4) Titration involving weak acids versus strong bases, strong acids versus weak bases, and strong acids versus strong bases using the appropriate indicators and their applications in quantitative determination, e.g., concentrations, mole ratio, purity, water of crystallization, and composition.

(1) Meaning of Solubility.

(2) Saturated and unsaturated solutions.

(3) Saturated solution as an equilibrium system.

(4) Solubility expressed in terms of mol dm-3 and g dm-3 of solution/solvent.

(5) Solubility curves and their uses.

(6) Effect of temperature on solubility of a substance.

(7) Relationship between solubility and crystallization.

(8) Crystallization/recrystallization as a method of purification.

(9) Knowledge of soluble and insoluble salts of stated cations and anions.

(10) Calculations on solubility.

Generalization about the solubility of salts and their applications to qualitative analysis. e.g. Pb2+, Ca2+, Al3+, Cu2+, Fe2+, Fe3+, Cl–, Br–, I–, SO42-, S2-, and CO32-, Zn2+, NH4+, SO32-

Explanation of solubility rules.


(i) Factors affecting rates;

(ii) Theories of reaction rates;

(iii) Analysis and interpretation of graphs.

(b) Equilibrium:

(i) General Principle;


  (1) Definition of reaction rate.

(2) Observable physical changes: color, mass, temperature, pH, formation of precipitate, etc.

(1) Physical states, concentration/ pressure of reactants, temperature, catalysts, light, particle size, and nature of reactants.

(2) Appropriate experimental demonstration for each factor is required.

(1) Collision and transition state theories to be treated qualitatively only.

(2) Factors influencing collisions: temperature and concentration.

(3) Effective collision.

(4) Activation energy.

(5) Energy profile showing activation energy and enthalpy change.

Drawing of graphs and charts.

Explanation of reversible and irreversible reactions. Reversible reaction, i.e., dynamic equilibrium. Equilibrium constant K must be treated qualitatively. It must be stressed that K for a system is stable at a constant temperature.

Simple experiment to demonstrate reversible reactions.

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(ii) Le Chatelier’s principle.12.0 REDOX REACTIONS

(a) Oxidation and reduction process.

(b) Oxidizing and reducing agents.

(c) Redox equations

(d) Electrochemical cells:

(i) Standard electrode potential;

(ii) Draw a cell diagram and write cell notation.


  Prediction of the effects of external influence of concentration, temperature pressure, and volume changes on equilibrium systems.

(1) Oxidation and reduction in terms of:

(a) addition and removal of oxygen and hydrogen;

(b) loss and gain of electrons;

(c) change in oxidation numbers/states.

(2) Determination of oxidation numbers/states.

(1) Description of oxidizing and reducing agents in terms of:

(a) addition and removal of oxygen and hydrogen;

(b) loss and gain of electrons;

(c) change in oxidation numbers/state.

Balancing redox equations by:

(a) ion, electron, or change in oxidation number/states;

(b) half-reactions and overall reaction.


(1) Standard hydrogen electrode means standard electrode potential (Eo) and measurement.

(2) Only metal/metal ion systems should be used.


(iii) e.m.f of cells;(iv) Application of Electrochemical cells.

(e) Electrolysis:

(i) Electrolytic cells;

(ii) Principles of electrolysis;

(iii) Factors influencing discharge of species;

(iv) Faraday’s laws;

(v) Practical application;


(1) Electrochemical cells are a combination of two half-cells. (2) The meaning of magnitude and sign of the e.m.f.

(1) Distinction between primary and secondary cells

(2) Daniell cells, lead acid battery cells, dry cells, fuel cells, and their use as electrical energy generators from chemical reactions.


Comparison of electrolytic and electrochemical cells; weak and strong electrolyte.

Mechanism of electrolysis.

Limit electrolytes to molten PbBr2

and NaCl, dilute NaCl solution, concentrated NaCl solution, CuSO4(aq), dilute H2SO4, NaOH(aq), and CaCl2(aq) (using platinum or graphite and copper electrodes).

Simple calculations based on the relation 1F= 96,500 C and mole ratios to determine the mass, the volume of gases, the number of entities, charges, etc., using half and overall reactions.

Electroplating, extraction, and purification of metals.


(vi) Corrosion of metals.13.0 CHEMISTRY OF CARBON COMPOUNDS

(a) Classification

(b) Functional group

(b) Separation and purification of organic compounds.

(c) Petroleum/crude oil


(1) Corrosion treated as a redox process.

(2) Rusting of iron and its economic costs.

(3) Prevention based on the relative magnitude of electrode potentials and preventive methods like galvanizing, sacrificial/cathodic protection, and non-redox methods (painting, greasing/oiling, etc.).

Broad classification into straight chain, branched chain, aromatic and alicyclic compounds.

Systematic nomenclature of compounds with the following functional groups: alkanes, alkenes, alkynes, hydroxyl compounds (aliphatic and aromatic), alkanoic acids, alkyl alkanoates (esters and salts), and amines.

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Methods to be discussed include distillation, crystallization, drying, and chromatography.

(1) Composition and classification.

(2) Fractional distillation and significant products.

(3) Cracking and reforming.

(4) Petrochemicals: sources; uses, e.g., as starting materials of organic synthesis.

(5) Quality of petrol, meaning of octane number and its importance to the petroleum industry.


(d) Determination of organic compounds’ empirical and molecular formulae and molecular structures. (e) General properties of organic compounds:

(i) Homologous series;

(ii) Isomerism.

(f) Alkanes:

(i) Sources, properties;

(ii) Uses.

(g) Alkenes:

(i) Sources and properties;


(1) Gradation in physical properties.

(2) Effects on the physical properties by introducing active groups into the inert alkane.

(1) Examples should be limited to compounds having a maximum of five carbon atoms.

(2) Differences between structural and geometric/stereoisomerism.

(1) Laboratory and industrial preparations and other sources.

(2) Nomenclature and structure.

(3) Reactivity:

(a) combustion;

(b) substitution reactions;

(c) cracking of large alkane molecules.

As fuels and starting materials for synthesis. Uses of haloalkanes and pollution effects.

(1) Laboratory preparation.

(2) Nomenclature and structure.


(ii) Uses;

(iii) Laboratory detection.

(h) Alkynes:

(i) Sources, characteristic properties, and uses;

(ii) Chemical reactions.

(i) Benzene:

(i) Structure and physical properties;

(ii) Chemical properties.


(3) Addition reactions with halogens hydrogen, bromine water, hydrogen halides, and acidified water.(4) Oxidation: hydroxylation with aqueous KMnO4.

(5) Polymerization.

The reaction with Br2/water, Br2/CCl4, and KMnO4(aq) is used to characterize alkenes.

(1) Nomenclature and structure.

(2) Industrial production of ethyne.

(3) Uses of ethyne.

(4) Distinguishing test between terminal and non-terminal alkynes.

(5) Test to distinguish between alkane, alkene, and alkyne.

Chemical reactions: halogenation, combustion, hydration, and hydrogenation.

Resonance in benzene. Stability leads to substitution reactions.

(1) Addition reactions: hydrogenation and halogenation (mechanism not required).

(2) Compare reactions with those of alkenes.


(J) Alkanols:(i) Sources, nomenclature and structure;

(ii) Classification;

(iii) Physical properties;

(iv) Chemical properties;

(v) Laboratory test;

(vi) Uses.

(k) Alkanoic acids:

(i) Sources, vocabulary, and structure;

(ii) Physical properties;


(1) Laboratory preparation, including hydration of alkenes.

(2) Industrial and local production of ethanol, including alcoholic beverages,

(3) Harmful impurities and methods of purification should be mentioned.

(4) Recognition of mono-, di- and triol structure.

Primary, secondary, and tertiary alkanols.

Boiling point, solubility in water. Including the hydrogen bonding effect.

(1) Reaction with:

(a) Na;

(b) alkanoic acids (esterification);

(c) conc. H2SO4.

(2) Oxidation by:

(a) KMnO4(aq);

(b) K2Cr2O7(aq);

(c) I2 in NaOH­(aq).

Laboratory test for ethanol.

Methanoic acid –insect bite.

Ethanoic acid – vinegar.

Recognition of mono and dioic acid.

Boiling point, solubility in water.

Including the hydrogen bonding effect.


(iii) Chemical properties;(iv) Laboratory test;

(iv) Uses.

(l) Alkanoates as derivatives of alkanoic acids:

(i) Sources, nomenclature, preparation and structure;

(ii) Physical properties;

(iii) Chemical properties;

(iv) Uses.


(a) Chemical industry


Acid properties only, i.e., reactions with H2O, NaOH, NH3, NaHCO3, Zn, and Mg.

Reaction with NaHCO3, Na2CO3.

Ethanoic and phenyl methanoic (benzoic) acids are used as examples of aliphatic and aromatic acids, respectively.

Preparation of alkyl alkanoates (esters) from alkanoic acids.

Solubility, boiling, and melting point.

Hydrolysis of alkyl alkanoates (mechanism not required).

Alkanoates are used in soap production, flavoring agents, plasticizers, solvents, and perfumes.

(1) Natural resources in the candidate’s won country.

(2) Chemical industries in the candidates’ country and their corresponding raw materials.

(3) Distinction between delicate and heavy chemicals.


(b) Pollution: air, water and soil pollution;

(c) Biotechnology.


     (a) Proteins:

(i) Sources and properties;

(ii) Uses of protein.

(b) Amino acids


(4) Factors that determine location of chemical industries.(5) Effect of industries on the community.

(1) Sources, effects, and control.

(2) Greenhouse effect and depletion of the ozone layer.

(3) Biodegradable and non-biodegradable pollutants.

Food processing and fermentation, including production of gari, bread, and alcoholic beverages, e.g., Local gin.

Proteins are polymers of amino acid molecules linked by peptide or amide linkage.

Physical properties, e.g., solubility

Chemical properties include:

(a) hydrolysis of proteins;

(b) laboratory test using Ninhydrin/Biuret reagent/Millons reagent.

(1) Nomenclature and general structure of amino acids.

(2) Difunctional nature of amino acids.


(c) Fats/oils:(i) Sources and properties;

(ii) General structure of fats/oils;

(iii) Preparation of soap;

(iv) Uses of fats/oils.

(d) Carbohydrates:

(i) Sources and nomenclature;

(ii) Properties;


As alkyl alkanoates (esters).

From animals and plants.

Physical properties such as solubility.

Chemical properties:

(a) acidic and alkaline hydrolysis;

(b) hydrogenation;

(c) test for fats and oil.

As mono-, di-, and tri-esters of propane-1,2,3-triol (glycerol).

(1) Preparation of soap (saponification) from fats and oils.

(2) Comparison of soapless detergents and their action on soft and hard water.

(1) Classes of carbohydrates as:

(a) monosaccharides;

(b) disaccharides;

(c) polysaccharides.

(2) Name and components of various classes of carbohydrates.

(1) Physical properties such as solubility of sugars.

(2) Chemical properties- Hydrolysis of disaccharides into monosaccharides.

(3) Test for reducing sugars using sugar strips, Fehling’s or Benedict’s solution, or Tollen’s reagent.


(iii) Carbohydrates as examples of polymer;(iv) Uses.

(e) Synthetic polymers:

(i) Properties;

(ii) Uses of polymers.

(1) Starch is a polymer made up of glucose units.

(2) Condensation of monosaccharides to form disaccharides and polysaccharides.

(1) Definition of terms: monomers, polymers, and polymerization.

(2) Addition and condensation polymerization.

(3) Classification and preparation based on the monomers and comonomers.

(1) Thermoplastics and thermosets.

(2) Modification of properties of polymers.

(3) Plastics and resins.

(4) Chemical test on plastics using:

(a) heat;

(b) acids;

(c) alkalis.

Frequently Asked Questions (FAQs)

Q1: How should I approach organic chemistry?

A1: Break it down into smaller sections. Focus on understanding reaction mechanisms and recognizing patterns.

Q2: Is memorization necessary?

A2: While understanding concepts is crucial, memorization aids quick recall during the exam. Balance both approaches.

Q3: Do you have any tips for mastering physical chemistry?

A3: Practice numerical problems regularly. Understanding the underlying concepts will make problem-solving easier.


The first step toward exam success is equipping yourself with the 2024 WAEC Chemistry Syllabus. Follow our guide, stay consistent in your preparation, and conquer every aspect of the syllabus. Best of luck!

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