Are you an SSS3 student about to sit for the forthcoming West African Examination Council (WAEC) examination and you are desperately searching for the latest WAEC Syllabus for Chemistry? In this article, we shall painstakingly give an updated version of the latest WAEC Chemistry Syllabus rundown uploaded on the website.
This curriculum was developed by the West African Examination Council for students preparing to take the senior secondary school exams, whether it be a school examination or a private examination (first and second series).
A list of all the topics that will be covered in the test or class is the list of documents contained in the WAEC Chemistry Syllabus.
Look for a different textbook to read and practice from after going over the concepts as a student. Utilize the online WAEC Chemistry previous questions and answers to assess your knowledge or level of preparedness.
How to Use WAEC Chemistry Syllabus 2023/2024
1. The syllabus can be downloaded in PDF format.
2. Obtain a different Chemistry text.
3. In the textbook, look up the subject. Study what you’ve read, reflect on it, and internalize it.
4. To gauge how well you’ve prepared, look for the WAEC Chemistry Series (Past Question).
WAEC Syllabus for Chemistry 2023/2024
This syllabus is designed solely for WAEC examination purposes, for this reason, topics are not necessarily arranged in the order in which they should be taught in class.
When creating the curriculum, the following presumptions were made:
1. The applicant must have completed the integrated science/basic science or general science and mathematics courses in junior high school (J.H.S.) or junior secondary school (JSS);
2. To build the necessary competencies and skills as outlined in the pertinent Chemistry teaching syllabuses, that applicant would engage in as many of the prescribed activities and project work as they could;
3. that laboratories in schools that teach the subject are well-equipped.
Note: The significant figures, SI units, and the traditional/IUPAC naming scheme must all be familiar to candidates.
The syllabus’s goals and objectives are to gauge candidates’ understanding of
1. Having an understanding of fundamental chemistry principles;
2. The degree to which laboratory skills have been acquired, including knowledge of risks and precautions;
3. Comprehension of how chemistry and other sciences are related to one another;
A person’s degree of knowledge on the advantages and risks associated with the relationship between chemistry and industry, the environment, and daily life;
4. Rational and analytical thinking abilities.
WAEC Chemistry 2023 Examination Scheme
There shall be three papers – Papers 1, 2, and 3 all of which must be taken. Papers 1 and 2 shall be composite papers to be taken in one sitting.
PAPER 1: 50 multiple-choice questions based on the syllabus Section A will make up the exam. (ie the portion of the syllabus which is common to all candidates). For a total of 50 points, candidates must respond to every question within a single hour.
PAPER 2: This will be a two-hour essay covering the curriculum with a 100-point grade. The essay will be divided into Sections A and B.
Section A: comprises ten briefs, and structured questions chosen from the common section of the course. (i.e. Section A of the syllabus). Each question must be answered correctly for 25 marks, according to the candidates.
Section B: Two questions will be from the syllabus segment that is universal to all candidates (segment A of the syllabus), and the other two questions will be from the section that is unique to the candidate’s nation (Section B or C of the syllabus).
Three of the questions may be chosen by the candidates to be answered. There are 25 marks assigned to each question.
PAPER 3: The practical exam will last two hours for school-bound applicants or one hour and thirty minutes for private candidates as an alternative to the practical work test.
The following curriculum elements will be covered in the questions:
1. One question on quantitative analysis;
2. One question on qualitative analysis;
3. The third question shall test candidates’ familiarity with the practical activities suggested in their teaching syllabuses.
Current WAEC syllabus for Chemistry 2023/2024 PDF Download
This manual communicates information about a specific topic of study and defines expectations and responsibilities.
Introduction To Chemistry
(a) (i) Measurement of physical quantities.
(ii) Scientific measurements and their importance in chemistry.
(b) Scientific Methods
Structure of the Atom
(a) Gross features of the atom.
(b) (i) Atomic number/proton number, number of neutrons, isotopes, atomic mass, mass number.
(I) Measurement of mass, length, time, temperature and volume.
(II) Appropriate SI units and significant figures.
(III) Precision and accuracy in measurement.
Outline the scientific method to include:
Observation, hypothesis, experimentation, formulation of laws and theories.
(b) a Short account of Dalton’s atomic theory and limitations, J.J. Thompson’s experiment and Bohr’s model of the atom.
(c) 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.
(i) 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
(iii) Rules and principles for filling in electrons.
(e) 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 the 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, the magnetization of 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
Aufbau Principle, Hund’s Rule of Maximum Multiplicity and Pauli Exclusion Principle.
Abbreviated and detailed electron configuration in terms of s, p, and d.
Standard Separation Techniques For Mixtures
(a) Classification of mixtures.
(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.
The boiling point for liquids and melting point for solids.
Electron configurations lead to group and periodic classifications.
Metals, semi-metals, non-metals in the periodic table and halogens. Alkali metals, alkaline earth metals and transition metals as metals.
Explanation of the periodic law.
Periodic properties; atomic size, ionic size, ionization energy, electron affinity and electronegativity.
Simple discrepancies should be accounted for with respect to beryllium, boron, oxygen and nitrogen.
(i) metallic to the non-metallic character of the element;
(ii) ionic to covalent bonding in compounds.
(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
(1) 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 reaction 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, metallic properties 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) Formation of ionic bonds and compounds.
(ii) Properties of ionic compounds.
(c) Naming of ionic compounds.
(d) Formation of covalent bonds and compounds.
(e) Properties of covalent compounds.
(ii) Coordinate (dative) covalent bonding.
(4) Other properties of transition metals should include:
(a) variable oxidation states;
(b) formation of coloured compounds;
(c) complex formation;
(d) catalytic abilities;
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.
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 shapes for some compounds e.g. BeCl2, BF3, CH4, NH3, CO2. Factors should include atomic radius, ionization energy and the number of valence electrons. Types of specific packing are not required.
Typical properties include heat and electrical conductivity, malleability, lustre, 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.
Biochemistry and Chemical Reactions
(a) 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 reactions such as the calculation of the percentage composition of an element.
(1) Combustion reactions (including combustion of simple hydrocarbons)
(3) Displacement or replacement
(5) Ionic reactions
(6) Laws of conservation of mass.
(7) Law of constant composition.
(8) Law of multiple proportions. Explanation of the laws to balance given equations.
(9) Experimental illustration of the law of conservation of mass.
(10) Mass and volume measurements.
(11) The mole as a unit of measurement; Avogadro’s constant, L= 6.02 x 1023 entities mol-1.
(12) Molar quantities and their uses.
(13) 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 by the method of dilution.
Use of mole ratios in determining the stoichiometry of chemical reactions. Simple calculations to 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 be expressed as a 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. (Dilution factor)
States of Matter
(a) (i) Kinetic theory of matter.
(ii) Changes of state of matter.
(1) Postulates of the kinetic theory of matter.
(2) Use of the kinetic theory to explain the following processes: melting of solids, boiling of liquids, evaporation of liquids, dissolution of solutes, Brownian motion and diffusion.
(3) 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 to a solid state and vice versa.
(4) Illustrations of changes of state using the different forms of water, iodine, sulphur, naphthalene etc.
(5) Brownian motion to be illustrated using any of the following experiments:
(a) pollen grains/powdered sulphur in water (viewed under a microscope);
(b) smoke in a glass container illuminated by a strong light from the side;
(c) a dusty room being swept and viewed from outside under sunlight.
(6) Experimental demonstration of diffusion of two gases.
(7) Relationship between the speed at which different gas particles move and the masses of particles.
(8) Experimental demonstration of diffusion of solute particles in liquids.
(i) Characteristics and nature of gases;
(ii) The gas laws;
(iii) Laboratory preparation and properties of some gases.
(b) (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. Qualitative explanation of each of the gas laws 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 real gases from ideal 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 vapour, vapour pressure, saturated vapour pressure, boiling and evaporation.
(2) Distinction between vapour and gas.
(3) Effect of vapour pressure on boiling points of liquids.
(4) Boiling at reduced pressure.
(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.
Energy and Energy Changes
(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 type of 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 as indicators of the purity of solids e.g. Phenyl methanoic acid (benzoic acid), ethanedioic acid (oxalic) and ethanamide.
Explanation of the terms energy and enthalpy. Energy changes are 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 involved in the following processes: combustion, dissolution and neutralization.
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’s concepts of acids and bases in terms of 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 the 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.
(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.
(5) Definition of pH and knowledge of pH scale.
(6) Measurement of pH of solutions using pH meter, calorimetric methods or universal indicator.
(7) Significance of pH values in everyday life e.g. acid rain, pH of the soil, blood, and urine.
(i) Laboratory and industrial preparation of salts;
(iii) Hydrolysis of salt.
(g) Deliquescent, efflorescent and hygroscopic compound.
(h) Acid-Base indicators
(i) Acid-Base titration
Meaning of salts.
Types of salts: normal, acidic, basic, double and complex salts.
(1) Description of laboratory and industrial production of salts.
(2) Mining of impure sodium chloride and conversion into granulated salt.
(3) Preparation of NaOH, Cl2 and H2.
(4) Explanation of how salts form acidic, alkaline and neutral aqueous solutions.
(5) Behaviour of some salts (e.g NH4Cl, AlCl3, Na2CO3, CH3COONa) in water as examples of equilibrium systems.
(6) Effects of charge density of some cations and anions on the hydrolysis of their aqueous solution. Examples 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) Colour of the indicator at any pH is dependent on relative amounts of acid and forms.
(4) Working pH ranges of methyl orange and phenolphthalein.
(5) Knowledge and correct use of relevant apparatus.
(6) Knowledge of how acid-bases indicators work in titrations.
Solubility of Substances
(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, the 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.
Chemical Kinetics and Equilibrium System
(a) Rate of reactions:
(i) Factors affecting rates;
(ii) Theories of reaction rates;
(iii) Analysis and interpretation of graphs.
(i) General Principle;
(1) Definition of reaction rate.
(2) Observable physical changes: colour, mass, temperature, pH, formation of precipitate etc.
(3) Physical states, concentration/ pressure of reactants, temperature, catalysts, light, particle size and nature of reactants.
(4) Appropriate experimental demonstration for each factor is required.
(5) Collision and transition state theories are to be treated qualitatively only.
(6) Factors influencing collisions: temperature and concentration.
(7) Effective collision.
(8) Activation energy.
(9) 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 constant at a constant temperature.
A simple experiment to demonstrate reversible reactions.
Le Chatelier’s principle.
(a) Oxidation and reduction process.
(b) Oxidizing and reducing agents.
(c) Redox equations
(d) Electrochemical cells:
(i) Standard electrode potential;
(ii) Drawing of cell diagram and writing 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: meaning of standard electrode potential (Eo) and its measurement.
(2) Only metal/metal ion systems should be used.
(iii) e.m.f of cells;
(iv) Application of Electrochemical cells.
(i) Electrolytic cells;
(ii) Principles of electrolysis;
(iii) Factors influencing discharge of species;
(iv) Faraday’s laws;
(v) Practical application;
(1) Electrochemical cells as a combination of two half-cells.(2) The meaning of magnitude and sign of the e.m.f.
(2) Distinction between primary and secondary cells
(3) Daniell cells, lead acid battery cells, dry cells, fuel cells and their use as generators of electrical energy from a chemical reaction
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
Chemistry of Carbon Compounds
(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 the 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.
Methods to be discussed should include distillation; crystallization; drying and chromatography.
(1) Composition and classification.
(2) Fractional distillation and major 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 empirical and molecular formulae and molecular structures of organic compounds.(e) General properties of organic compounds:
(i) Homologous series;
(i) Sources, properties;
(i) Sources and properties;
(1) Gradation in physical properties.
(2) Effects on the physical properties by the introduction of 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.
(b) substitution reactions;
(c) cracking of large alkane molecules.
As fuels, as starting materials for synthesis. Uses of haloalkanes and pollution effects.
(1) Laboratory preparation.
(2) Nomenclature and structure.
(iii) Laboratory detection.
(i) Sources, characteristic properties and uses;
(ii) Chemical reactions.
(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.
Use of reaction with Br2/water, Br2/CCl4 and KMnO4(aq) as means of characterizing alkenes.
(1) Nomenclature and structure.
(2) Industrial production of ethyne.
(3) Uses of ethyne.
(4) Distinguishing test between the 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.
(i) Sources, nomenclature and structure;
(iii) Physical properties;
(iv) Chemical properties;
(v) Laboratory test;
(k) Alkanoic acids:
(i) Sources, nomenclature 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 the structure of mono-, di- and triols.
Primary, secondary and tertiary alkanols.
Boiling point, solubility in water. Including the hydrogen bonding effect.
(1) Reaction with:
(b) alkanoic acids (esterification);
(c) conc. H2SO4.
(2) Oxidation by:
(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;
(l) Alkanoates as derivatives of alkanoic acids:
(i) Sources, nomenclature, preparation and structure;
(ii) Physical properties;
(iii) Chemical properties;
Chemistry, Industry and the Environment
(a) Chemical industry
Acid properties only i.e. reactions with H2O, NaOH, NH3, NaHCO3, Zn and Mg.
Reaction with NaHCO3, Na2CO3.
Uses of ethanoic and phenyl methanoic (benzoic) acids 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).
Uses of alkanoates include the production of soap, flavouring agents, plasticizers, solvents and in perfumes.
(1) Natural resources in the candidate’s won country.
(2) Chemical industries in candidates’ own country and their corresponding raw materials.
(3) Distinction between fine and heavy chemicals.
(b) Pollution: air, water and soil pollution;
Basic Biochemistry and Synthetic Polymers
(i) Sources and properties;
(ii) Uses of protein.
(b) Amino acids
(4) Factors that determine the 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 as polymers of amino acid molecules linked by peptide or amide linkage.
Physical properties e.g. solubility
Chemical properties include:
(a) hydrolysis of proteins;
(b) A laboratory test using Ninhydrin/Biuret reagent/Millons reagent.
(1) Nomenclature and general structure of amino acids.
(2) Difunctional nature of amino acids.
(i) Sources and properties;
(ii) General structure of fats/oils;
(iii) Preparation of soap;
(iv) Uses of fats/oils.
(i) Sources and nomenclature;
1. As alkyl alkanoates (esters).
2. From animals and plants.
3. Physical properties such as solubility.
4. Chemical properties:
(a) acidic and alkaline hydrolysis;
(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:
(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) chemistry as examples of polymer;(iv) Uses.
(e) Synthetic polymers:
(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:
If you find this article helpful someone might also need it so don’t hesitate to share.