WAEC Chemistry

WAEC Syllabus
Syllabus 2019/2020
- Uploaded | See
online by Current Chemistry Syllabus Here
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CHEMISTRY

1. INTRODUCTION

This syllabus is drawn purposely for examination, hence the topics are not necessarily
arranged in the order in which they should be taught.

The following assumptions were made in drawing of the syllabus:

(1) That candidates must have covered the Integrated Science/Basic Science or General
Science and Mathematics syllabuses at the Junior Secondary School (JSS)/Junior High
School (J.H.S) level;

(2) That candidates would carry out as many of the suggested activities and project work as
possible, and consequently develop the intended competencies and skills as spelt out in
the relevant Chemistry teaching syllabuses;

(3) That schools which offer the subject have well-equipped laboratories.

Note: Candidates are required to have the knowledge of the significant figures, S.I.
units and the conventional/IUPAC system of nomenclature.

2. AIMS

The aims and objectives of the syllabus are to assess candidates’

(1) understanding of basic chemistry concepts;

(2) level of acquisition of laboratory skills including awareness of hazards and safety
measures;
(3) level of awareness of the inter-relationship between chemistry and other discipline;
(4) level of awareness of the linkage between chemistry and industry/environment/everyday
life in terms of benefits and hazards;
(5) skills of critical and logical thinking.

3. EXAMINATION SCHEME

There shall be three papers - Papers 1, 2 and 3 all of which must be taken. Paper 1 and 2
shall be a composite paper to be taken at one sitting.

PAPER 1: Will consist of fifty multiple choice objective questions drawn from Section A of
the syllabus (ie the portion of the syllabus which is common to all candidates) .
Candidates will be required to answer all the questions within 1 hour for 50
marks.

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PAPER 2: Will be a 2-hour essay paper covering the entire syllabus and carrying
100 marks. The paper will be in two sections; Sections A and B.

Section A: Will consist of ten short structured questions drawn from the
common portion of the syllabus. (i.e. Section A of the syllabus).
Candidates will be required to answer all the questions for 25
marks.

Section B: Will consist of two questions from the common portion of the
syllabus (i.e. Section A of the syllabus) and two other questions
from the section of the syllabus which is perculiar to the country of
the candidate (i.e. either Section B or C of the syllabus).
Candidates will be required to answer any three of the questions.
Each question shall carry 25 marks.

PAPER 3: This shall be a 2-hour practical test for school candidates or 1 hour
30 minutes alternative to practical work test for private candidates. Each
version of the paper shall contain three compulsory questions and carry 50
marks.

The questions shall be on the following aspects of the syllabus:

- One question on quantitative analysis;

- One question on qualitative analysis;
- The third question shall test candidates’ familiarity with the
practical activities suggested in their teaching syllabuses.

Details of the input into the continuous assessment shall be given by the Council.

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SECTION A

(For all candidates)

CONTENT NOTES

1.0 INTRODUCTION TO CHEMISTRY

(a) (i) Measurement of physical quantities.

(ii) Scientific measurements and their (1) Measurement of mass, length, time,
importance in chemistry. temperature and volume.
(2) Appropriate SI units and significant
figures.
(3) Precision and accuracy in
measurement.

(b) Scientific Methods Outline the scientific method to include:

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

2.0 STRUCTURE OF THE ATOM

(a) Gross features of the atom. (1) Short account of Dalton’s atomic
theory and limitations, J.J.
Thompson’s experiment and Bohr’s
model of the atom.
(2) Outline description of the
Rutherford’s alpha scattering
experiment to establish the structure
of the atom.

(b) (i) Atomic number/proton number, Meaning and representation in symbols of

number of neutrons, isotopes, atomic atoms and sub-atomic particles.
mass, mass number.

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CONTENT NOTES

(ii) Relative atomic mass (Ar) and (1) Atomic mass as the weighted average
relative molecular mass (Mr) based mass of isotopes. Calculation of
on Carbon-12 scale. relative mass of chlorine should be
used as an example.
(2) Carbon-12 scale as a unit of
measurement.
Definition of atomic mass unit.

(iii) Characteristics and Atoms, molecules and ions.

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

(c) Particulate nature of mater: physical and Explain physical and chemical changes
chemical changes. with examples.
Physical change- melting of solids,
magnetization of iron, dissolution of salt
etc.
Chemical change- burning of wood,
rusting of iron, decay of leaves etc.

(d) (i) Electron Configuration Detailed electron configurations (s,p,d)

for atoms of the first thirty elements.

(ii) Orbitals Origin of s,p and d orbitals as sub-energy

levels; shapes of s and p orbitals only.

(iii) Rules and principles (1) Aufbau Principle, Hund’s Rule of

for filling in electrons. Maximum Multiplicity and Pauli
Exclusion Principle.
(2) Abbreviated and detailed electron
configuration in terms of s, p, and d.

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CONTENT NOTES

3.0 STANDARD SEPARATION

TECHNIQUES
FOR MIXTURES
(a) Classification of mixtures. Solid-solid, solid-liquid, liquid-liquid,
gas-gas with examples.

(b) Separation techniques Crystallization, distillation, precipitation,

magnetization, chromatography,
sublimation etc.

(c) Criteria for purity. Boiling point for liquids and melting
point for solids.

4.0 PERIODIC CHEMISTRY

(a) Periodicity of the elements. Electron configurations leading to group
and periodic classifications.

(b) Different categories of elements in the Metals, semi-metals, non-metals in the

periodic table. periodic table and halogens. Alkali
metals, alkaline earth metals and
transition metals as metals.

(c) Periodic law: Explanation of the periodic law.

(i) Trends on periodic table; Periodic properties; atomic size, ionic

size, ionization energy, electron affinity
and electronegativity.
Simple discrepancies should be
accounted for in respect to beryllium,
boron, oxygen and nitrogen.

(ii) Periodic gradation of the elements in (1) Progression from:

the third period (Na - Ar). (i) metallic to non-metallic character
of element;
(ii) ionic to covalent bonding in
compounds.

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CONTENTS NOTES

(2) Differences and similarities in the

properties between the second and the
third period elements should be
stated.

(d) Reactions between acids and metals, (1) Period three metals (Na, Mg, Al).
their oxides and trioxocarbonates (IV). (2) Period four metals (K, Ca).
(3) Chemical equations.
(4) pH of solutions of the metallic oxides
and trioxocarbonates.

(e) Periodic gradation of elements in group Recognition of group variations noting

seven, the halogens: F, Cl, Br and I. 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) reaction of the elements with water
and alkali (balanced equations
required).

(f) Elements of the first transition series. (1) Their electron configurations,
21Sc – 30Zn 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, acids and comparison with s-
block elements (Li, Na, Be, Mg).

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CONTENT NOTES

(4) Other properties of transition metals

should include:
(a) variable oxidation states;
(b) formation of coloured
compounds;
(c) complex formation;
(d) catalytic abilities;
(e) paramagnetism;
(f) hardness.

5.0 CHEMICAL BONDS

(a) Interatomic bonding Meaning of chemical bonding.
Lewis dot structure for simple ionic and
covalent compounds.

(b) (i) Formation of ionic bonds and Formation of stable compounds from
compounds. ions. Factors influencing formation:
ionzation energy; electron affinity and
electronegativity difference.

(ii) Properties of ionic compounds. Solubility in polar and non-polar

solvents, electrical conductivity, hardness
and melting point.

(c) Naming of ionic compounds. IUPAC system for simple ionic

compounds.

(d) Formation of covalent bonds and Factors influencing covalent bond

compounds. formation. Electron affinity, ionization
energy, atomic size and electronegativity.

(e) (i) Properties of covalent compounds. Solubility in polar and non-polar

solvents, melting point, boiling point and
electrical conductivity.

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

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CONTENT NOTES

(f) Shapes of molecular compounds. Linear, planar, tetrahedral and shapes for
some compounds e.g. BeCl2, BF3, CH4,
NH3, CO2.

(g) (i) Metallic Bonding

(ii) Factors influencing its formation. Factors should include: atomic radius,
ionization energy and number of valence
electrons. Types of specific packing not
required.

(iii) Properties of metals. Typical properties including heat and

electrical conductivity, malleability,
lustre, ductility, sonority and hardness.

(h) (i) Inter molecular bonding Relative physical properties of polar and
non-polar compounds.
(ii) Intermolecular forces in covalent Description of formation and nature
compounds. should be treated.
Dipole-dipole, induced dipole-dipole,
induced dipole-induced dipole forces
should be treated under van der Waal’s
forces.

(iii) Hydrogen bonding 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
(iii) van der Waals forces forces; and variation in the boiling points
of H2O, and H2S explained using
Hydrogen bonding.

(iv) Comparison of all bond types.

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CONTENT NOTES

6.0 STOICHIOMETRY AND CHEMICAL

REACTIONS
(a) (i) Symbols, formulae and equations. Symbols of the first thirty elements and
other common elements that are not
among the first thirty elements.

(ii) chemical symbols

(iii) Empirical and molecular formulae. Calculations involving formulae and

equations will be required. Mass and
volume relationships in chemical
reactions and the stoichiometry of such
reactions such as: calculation of
percentage composition of element.

(iv) Chemical equations and IUPAC (1) Combustion reactions (including

names of chemical compounds. combustion of simple hydrocarbons)
(2) Synthesis
(3) Displacement or replacement
(4) Decomposition
(5) Ionic reactions

(v) Laws of chemical combination. (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.

(b) Amount of substance. (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.

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CONTENT NOTES

(c) Mole ratios Use of mole ratios in determining

stoichiometry of chemical reactions.
Simple calculations to determine the
number of entities, amount of substance,
mass, concentration, volume and
percentage yield of product.

(d) (i) Solutions (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.

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

(iii) Standard solutions. (1) Preparation of some primary

standards e.g anhydrous Na2CO3,
(COOH)2, 2H2O/H2C2O4.2H2O.
(2) Meanning of the terms primary
standard, secondary standard and
standard solution.

(e) Preparation of solutions from liquid Dilution factor

solutes by the method of dilution.

.
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CONTENT NOTES

7.0 STATES OF MATTER

(a) (i) Kinetic theory 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.

(ii) Changes of state of matter. (1) Changes of state of matter should be

explained in terms of movement of
particles. It should be emphasized that
randomness decreases (and orderliness
increases) from gaseous state to liquid
state and to solid state and vice versa.
(2) Illustrations of changes of state using the
different forms of water, iodine, sulphur,
naphthalene etc.
(3) 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.

(iii) Diffusion
(1) Experimental demonstration of
diffusion of two gases.
(2) Relationship between speed at which
different gas particles move and the
masses of particles.
(3) Experimental demonstration of
diffusion of solute particles in
liquids.

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CONTENT NOTES

(b) Gases:
(i) Characteristics and nature of gases; Arrangement of particles, density, shape
and compressibility.

(ii) The gas laws; The Gas laws: Charles’; Boyle’s;

Dalton’s law of partial pressure;
Graham’s law of diffusion, 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, temperature.
Mathematical relations of the gas law
PV= nRT
Ideal and Real gases
Factors responsible for the deviation of
real gases from ideal situation.

(iii) Laboratory preparation and properties of (1) Preparation of the following gases:
some gases. H2, NH3 and CO2. Principles of
purification and collection of gases.
(2) Physical and chemical properties of
the gases.

(c) (i) Liquids Characteristics and nature of liquids

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

(ii) Vapour and gases. (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.

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CONTENT NOTES

(d) Solids:
(i) Characteristics and nature; (1) Ionic, metallic, covalent network and
molecular solids. Examples in each
case.
(2) Arrangements of particles ions,
molecules and atoms in the solid
state.

(ii) Types and structures; 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.

(iii) Properties of solids.

(e) Structures, properties and uses of The uses of diamond and graphite related
diamond and graphite. to the structure.
The use of iodine in everyday life.

(f) Determination of melting points of Melting points as indicator of purity of

covalent solids. solids e.g. Phenyl methanedioic acid
(benzoic acid), ethanedioic acid (oxalic)
and ethanamide.

8.0 ENERGY AND ENERGY CHANGES

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

(b) Description, definition and illustrations (1) Exothermic and endothermic

of energy changes and their effects. 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.

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CONTENT NOTES

9.0 ACIDS, BASES AND SALTS

(a) Definitions of acids and bases. (1) Arrhenius 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.

(b) Physical and chemical properties of acids Characteristic properties of acids and
and bases. bases in aqueous solution to include:
(a) conductivities, taste,
litmus/indicators, feel etc.;
(b) balanced chemical equations of all
reactions.

(c) Acids, bases and salts as electrolytes. Electrolytes and non-electrolytes; strong
and weak electrolytes. Evidence from
conductivity and enthalpy of
neutralization.

(d) Classification of acids and bases. (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 example of
equilibrium systems.

(e) Concept of pH (1) Definition of pH and knowledge of

pH scale.
(2) Measurement of pH of solutions
using pH meter, calometric methods
or universal indicator.
(3) Significance of pH values in
everyday life e.g. acid rain, pH of
soil, blood, urine.

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CONTENT NOTES

(f) Salts: Meaning of salts.

Types of salts: normal, acidic, basic,
double and complex salts.

(i) Laboratory and industrial preparation (1) Description of laboratory and

of salts; industrial production of salts.
(ii) Uses; (2) Mining of impure sodium chloride
and conversion into granulated salt.
(3) Preparation of NaOH, Cl2 and H2.

(iii) Hydrolysis of salt. (1) Explanation of how salts forms

acidic, alkaline and neutral 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
to be taken from group 1, group 2,
group 3 and the d-block element.
(g) Deliquescent, efflorescent and
hygroscopic compound. Use of hygroscopic compounds as
drying agent should be emphasized.
(h) Acid-Base indicators
(1) Qualitative description of how acid-
base indicator works.
(2) Indicators as weak organic acids or
bases (organic dyes).
(3) Colour of indicator at any pH
dependent on relative amounts of
acid and forms.
(4) Working pH ranges of methyl orange
and phenolphthalein.
(i) Acid-Base titration
(1) Knowledge and correct use of
relevant apparatus.
(2) Knowledge of how acid-bases
indicators work in titrations.
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CONTENT
NOTES

(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.

10.0 SOLUBILITY OF SUBTANCES

(a) General principles (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.

(b) Practical application of solubility. Generalization about 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.

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CONTENT NOTES

11.0 CHEMICAL KINETICS AND

EQUILIBRIUM SYSTEM
(a) Rate of reactions: (1) Definition of reaction rate.
(2) Observable physical and changes:
colour, mass, temperature, pH,
formation of precipitate etc.

(i) Factors affecting rates; (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.

(ii) Theories of reaction rates; (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.

(iii) Analysis and interpretation of graphs. Drawing of graphs and charts.

(b) Equilibrium:
(i) General Principle; 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 constant
temperature.
Simple experiment to demonstrate
reversible reactions.

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CONTENT NOTES

(ii) Le Chatelier’s principle. Prediction of the effects of external

influence of concentration, temperature
pressure and volume changes on
equilibrium systems.

12.0 REDOX REACTIONS

(a) Oxidation and reduction process. (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.

(b) Oxidizing and reducing agents. (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.
(c) Redox equations
Balancing redox equations by:
(a) ion, electron or change in oxidation
number/states;
(b) half reactions and overall reaction.
(d) Electrochemical cells:
Definition/Explanation
(i) Standard electrode potential; (1) Standard hydrogen electrode:
meaning of standard electrode
potential (Eo) and its measurement.
(2) Only metal/metal ion systems should
be used.

(ii) Drawing of cell diagram and writing

cell notation.

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CONTENT NOTES

(1) Electrochemical cells as a

(iii) e.m.f of cells; combination of two half-cells.
(2) The meaning of magnitude and sign
of the e.m.f.

(1) Distinction between primary and

(iv) Application of Electrochemical cells.
secondary cells
(2) Daniell cell, lead acid battery cell, dry
cells, fuel cells and their use as
generators of electrical energy from
chemical reactions.

(e) Electrolysis: Definition.

(i) Electrolytic cells; Comparison of electrolytic and

electrochemical cells; weak and strong
electrolyte.
(ii) Principles of electrolysis;
Mechanism of electrolysis.

(iii) Factors influencing discharge of Limit electrolytes to molten PbBr2

species; and NaCl, dilute NaCl solution,
concentrated NaCl solution, CuSO4(aq),
dilute H2SO4, NaOH(aq) and CaCl2(aq)
(using platinum or graphite and copper
electrodes).
(iv) Faraday’s laws;
Simple calculations based on the relation
1F= 96,500 C and mole ratios to
determine mass, volume of gases,
number of entities, charges etc. using
half and overall reactions.
(v) Practical application;
Electroplating, extraction and purification
of metals.

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CONTENT NOTES

(vi) Corrosion of metals. (1) Corrosion treated as a redox process.

(2) Rusting of iron and its economic
costs.
(3) Prevention based on relative
magnitude of electrode potentials
and preventive methods like
galvanizing, sacrificial/cathodic
protection and non-redox methods
(painting, greasing/oiling etc.).
13.0 CHEMISTRY OF CARBON
COMPOUNDS

(a) Classification Broad classification into straight chain,

branched chain, aromatic and alicyclic
compounds.
(b) Functional group
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:

(b) Separation and purification of organic
distillation; crystallization; drying and
compounds.
chromatography.

(1) Composition and classification.

(c) Petroleum/crude oil (2) Fractional distillation and major
products.
(3) Cracking and reforming.
(4) Petro-chemicals: 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.

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CONTENT NOTES

(d) Determination of empirical and

molecular formulae and molecular
structures of organic compounds.

(e) General properties of organic

compounds:

(i) Homologous series; (1) Gradation in physical properties.

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

(ii) Isomerism.
(1) Examples should be limited to
compounds having maximum of five
carbon atoms.
(2) Differences between structural and
geometric/stereo isomerism.

(f) Alkanes:

(i) Sources, properties; (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, as starting materials for

(ii) Uses. synthesis. Uses of haloakanes and
pollution effects.

(g) Alkenes:

(i) Sources and properties;

(1) Laboratory preparation.
(2) Nomenclature and structure.

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CONTENT NOTES

(3) Addition reactions with halogens

hydrogen, bromine water, hydrogen
halides and acidified water.
(4) Oxidation: hydroxylation with
aqueous KMnO4.
(5) Polymerization.

(ii) Uses;

(iii) Laboratory detection. Use of reaction with Br2/water, Br2/CCl4

and KMnO4(aq) as means of
characterizing alkenes.
(h) Alkynes:

(i) Sources, characteristic properties and (1) Nomenclature and structure.

uses; (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.

(ii) Chemical reactions.

Chemical reactions: halogenation,
combustion, hydration and
hydrogenation.
(i) Benzene:

(i) Structure and physical properties; Resonance in benzene. Stability leading

to substitution reactions.

(ii) Chemical properties.

(1) Addition reactions: hydrogenation
and halogenation (mechanism not
required).
(2) Compare reactions with those of
alkenes.

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CONTENT NOTES

(J) Alkanols:

(i) Sources, nomenclature and structure; (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.

(ii) Classification; Primary, secondary and tertiary alkanols.

(iii) Physical properties; Boiling point, solubility in water.

Including hydrogen bonding effect.

(iv) Chemical properties; (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).
(v) Laboratory test;
Laboratory test for ethanol.
(vi) Uses.

(k) Alkanoic acids:

Methanoic acid –insect bite.
Ethanoic acid – vinegar.

(i) Sources, nomenclature and structure; Recognition of mono and dioic acid.
(ii) Physical properties;
Boiling point, solubility in water.
Including hydrogen bonding effect.

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CONTENT NOTES

(iii) Chemical properties; Acid properties only i.e. reactions with

H2O, NaOH, NH3, NaHCO3, Zn and Mg.

(iv) Laboratory test; Reaction with NaHCO3, Na2CO3.

(iv) Uses.
Uses of ethanoic and phenyl methanoic
(benzoic) acids as examples of aliphatic
and aromatic acids respectively.

(l) Alkanoates as drivatives of alkanoic acids:

(i) Sources, nomenclature, preparation

and structure; Preparation of alkyl alkanoates (esters)
from alkanoic acids.

(ii) Physical properties; Solubility, boiling and melting point.

(iii) Chemical properties; Hydrolysis of alkyl alkanoates

(mechanism not required).

(iv) Uses. Uses of alkanoates to include

production of soap, flavouring agent,
plasticizers, as solvents and in
perfumes.

14.0 CHEMISTRY, INDUSTRY AND THE

ENVIRONMENT

(a) Chemical industry (1) Natural resources in candidate’s won

country.
(2) Chemical industries in candidates own
country and their corresponding raw
materials.
(3) Distinction between fine and heavy
chemicals.

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CONTENT NOTES

(4) Factors that determine location of

chemical industries.
(5) Effect of industries on the
community.

(b) Pollution: air, water and soil

(1) Sources, effects and control.
pollution;
(2) Greenhouse effect and depletion of
the ozone layer.
(3) Biodegradable and non-biodegradable
pollutants.
(c) Biotechnology.
Food processing, fermentation including
production of gari, bread and alcoholic
beverages e.g. Local gin.
15.0 BASIC BIOCHEMISTRY AND
SYNTHETIC POLYMERS

(a) Proteins:
Proteins as polymers of amino acids
molecules linked by peptide or amide
linkage.

(i) Sources and properties; Physical properties e.g. solubility

Chemical properties to include:
(a) hydrolysis of proteins;
(b) laboratory test using
Ninhydrin/Biuret reagent/Millons
reagent.

(ii) Uses of protein.

(b) Amino acids

(1) Nomenclature and general structure
of amino acids.
(2) Difunctional nature of amino acids.

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CONTACT NOTES

(c) Fats/oils: As alkyl alkanoates (esters).

(i) Sources and properties;
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-

(ii) General structure of fats/oils;
1,2,3-triol (glycerol).

(iii) Preparation of soap; (1) Preparation of soap (saponification)

from fats and oils.
(2) Comparison of soap less detergents
and their action on soft and hard
water.
(iv) Uses of fats/oils.

(d) Carbohydrates:

(i) Sources and nomenclature;

(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

(ii) Properties;
of sugars.
(2) Chemical properties- Hydrolysis of
disaccharides into monosaccharides.
(3) Test for reducing sugars using sugar
strips, Fehling’s or Benedicts solution
or Tollen’s reagent.

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CONTENT NOTES

(iii) Carbohydrate as examples of (1) Starch as a polymer made up of

polymer; glucose units.
(2) Condensation of monosaccharides to
form disaccharides and
polysaccharides.

(iv) Uses.

(e) Synthetic polymers: (1) Definition of terms: monomers,

polymers and polymerization.
(2) Addition and condensation
polymerization.
(3) Classification and preparation based
on the monomers and comonomers.
(i) Properties;
(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.

(ii) Uses of polymers.

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SECTION B
(For candidates in Ghana only)

CONTENTS NOTES

1.0 STRUCTURE OF THE ATOM

(a) Elementary treatment of mass (1) Qualitative knowledge of the mass
spectrometer. spectrometer: principles and operations of the
mass spectrometer; and its use to detect
isotopes, determination of Relative atomic and
molecular masses only.
(2) Wave nature of electrons.
(3) Quantum numbers and their importance.

Meaning of terms: Nucleons, nuclide.

(b) (i) Nuclear chemistry
Charges, relative mass and penetrating power of
(ii) Types and nature of radiations:
radiations. Meaning of radioactivity. Difference
alpha, beta particles and gamma
between spontaneous nuclear reactions
radiation.
(radioactivity) and induced nuclear reactions.

Natural and artificial radioactivity.

(iii) Radioactivity:
Detection of radiation by Geiger-Muller counter.
induced/stimulated.

Distinction between ordinary chemical reactions

(iv) Nuclear reactions: fission and
and nuclear reactions. Generations of electricity;
fusion in nuclear reactions.
atomic bombs. Balanced equations of nuclear
reactions

(1) Carbon dating (qualitative treatment only).

(v) Effects and application of
(2) Use of radioactivity in agriculture, medicine
radioactivity and industries.
(3) Hazards associated with nuclear radiations.

Factors affecting stability of nuclides:

Binding energy, neutron-proton ratio, and half life.
Calculations involving half-life

2.0 PERIODIC CHEMISTRY

(a) Reactions between acids (1) Period three metals (Na, Mg, Al)
and metals their oxides and (2) Period four metals (K, Ca)
trioxocarbonates (IV). (3) Chemical equations
(4) pH of solutions of the metallic oxides and
trioxocarbonates.
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CONTENTS NOTES

(b) Acidic properties of oxides of non- (1) Oxides of carbon, nitrogen, sulphur,
metals. phosphorus and chlorine.
(2) pH of aqueous solutions of the oxides.
(3) Chemical equations.

(c) Physical and chemical properties of (1) Comparison of the physical and chemical
period 3 elements and their properties of period three elements.
compounds. (2) Comparison of the physical and chemical
properties of (hydrides, oxides, hydroxides
and chlorides) compounds.
(3) Thermal stability of CO32- and NO3- of Li,
Na, K, Mg and Ca.
(4) Experiment to compare thermal stability of
Na2CO3/LiCO3/CuSO4.

(d) Silicon (1) Structures for SiO2 and CO2 account for the
differences between physical and chemical
properties of the two oxides.
(2) Uses of silicon and its compounds e.g.
ceramics, glass, silica gel and microchips.

(e) Periodic gradation of elements in (1) Inter- atomic bond energies.

group seven i.e. the halogens. (2) Hydrides and their acid strength comparison
of the Ka values of the hydrogen halides.
(3) Variable oxidation states of the halogens.
(f) Bonding in complex compounds.
Definition of ligands and central ions
Examples of ligands
(1) Formation of coordination compounds.
(2) Nomenclature of complex ions and
compounds (Cl-, F-, I-, NO3-, NH3, H2O,
SO42-).

(g) Shapes of complex compounds. Tetrahedral, square planar, octahedral e.g.

(Fe(CN)6]3-, [Cu(NH3)4]2+, [Ag)NH3)2]+ [Cu)CN)4]2

(h) Elements of the first transition Reactivity of the metals with air, water, acids and
series. comparison with s-block elements (Li, Na, Be,
Mg).
3.0. CHEMICAL BONDS
(a) Formation of Ionic bonds:
(i) Factors that influence ionic bond
Factors should include lattice energy.
formation;

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CONTENTS NOTES

(ii) Covalent character in ionic bond;

(iii) Polar covalent bonds. (1) Ionic character (polarity) in covalent bonds
based on electronegativity difference
between the species involved.
(2) Effects of covalent and ionic character in
ionic and covalent bonds on the solubility,
thermal stability and boiling points of ionic
and covalent compounds.
(b)(i) Hybridization of atomic orbitals.
Definition of Hybridization.

(ii) Formation of hybrid orbitals.

(1) Description of sp, sp2, sp3 hybrid orbitals.
(2) Shapes of sp, sp2, sp3 and sp3d2 hybrid
orbitals. Treatment should be limited to the
following molecules only. CH4, H2O, NH3,
BCl3, C2H2, BeCl2, C2H4 and SF6.
(iii) Formation of sigma (σ) and pi (π)
Description of sigma and pi bonds. Using C2H2 and
bonds.
C6H6.

4.0 SOLUTIONS
(a) Preparation of solutions from liquid
(1) Outline of steps involved in the preparation
solutes by the method of dilution. of solutions from liquid solutes.
(2) Determination of concentration of liquid
solutes (stock solution) given the density,
w/v, w/w), specific gravity, relative
molecular mass, molar mass, and % purity.
(3) Primary standard, secondary standard and
standardized solution.
5.0 ENERGY AND ENERGY CHANGES
(a) Energy changes in physical and
(1) Definition and understanding of the
isolated systems.
meaning of the energy terms: systems,
surroundings, open and closed.
(2) Enthalpy change involved in the following
processes: combustion, atomization,
sublimation, hydration/salvation and
dissolution.
(b) Hess’s Law of heat summation and
Born-Haber cycle. Explanation of Hess’s law and its application in
the development of the Born-Haber cycle.
(1) Use of difference cycles to illustrate Hess’s
law.

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CONTENTS NOTES

(2) Simple calculations using chemical

equations, energy cycles or diagrams with
given energy changes.

(c) Bond Energy Explanation of bond energy and bond

dissociation energy.
(1) Bond energy as an average value.
Differences in bond energy and bond
dissociation energy.
(2) Bond energy in molecules and its use in
assessment of bond strength, energy content
and enthalpy of reaction.
(3) Calculations using summation of bond
energies in reactants and products as a
measure of enthalpy of reaction.

6.0 ACIDS, BASES AND SALTS

(a) Definitions of acids and bases. (1) Bronsted – Lowry and Lewis concept of
acids and bases.
(2) Conjugate acid-base pair concept in terms
of equilibrium.
(b) pH, pOH and pKw
(1) Ionic product constant of water Kw =
[H+(aq)][OH(aq)] = 1.0 x 10-14 mol2dm-6.
(2) pH and pOH as a measure of acidity and
alkalinity respectively pH = -log[H3O+].
(3) Knowledge of pH scale.
(4) Calculation of [H+], [OH-] and the
corresponding pH and pOH of given
solutions.
(c) Partial ionization of weak acids and
weak bases. Explanation of pKa and pKb of weak acids and
bases.
(1) Behaviour of weak acids and weak bases in
water as example of equilibrium systems.
(2) Calculations involving Ka, pKa and Kb,
pKb.
(3) Ka, pKa and Kb, pKb as measurements of
acid and basic strengths respectively.

(d) Buffer Solutions (1) Qualitative definition of buffers.

(2) Examples of buffers from the laboratory.
(3) Preparation of buffer solutions.

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CONTENTS NOTES

(e) Acid base titrations Double indicator titrations (continuous and

Discontinuous) and back titration.
Calculations involving concentration, composition
and % purity.
Graphs for acid-based titrations. Nature of graphs
of strong acid and strong base, strong acid and
weak base and strong base and weak acid.

7.0 SOLUBILITY OF SUBSTANCES

(a) Solubility and solubility product. (1) Explanation of solubility products (Ksp) of
sparingly soluble ionic compounds.
(2) Calculations involving solubility and
solubility products.
(3) Factors affecting solubility.
(c) Crystallization and recrystallization.
Explanation of the effect of lattice energy and
hydration energy on crystallization and
recrystallization.
8.0 CHEMICAL KINETICS AND
EQUILIBRIUM SYSTEMS
(a) Rate law and Order of reaction
(1) Deduction of order and rate law from
experimental data.
(2) Simple relationship between rates and
concentration of zero, first and second order
reactions. Graphical representation of zero,
first and second order reactions.
(3) Half-life for first order reactions and its
significance.
(4) General rate law equation.
(5) Derivation of the rate expression from
experimentally determined rate data:
R = k[A]x [B]y where k = rate constant.

(b) Rate determining step of a multi-step

reaction.

(c) Equilibrium

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CONTENTS NOTES

(d) Equilibrium Law of Mass Action. (1) Mathematical expression for the
determination of equilibrium constant K
(2) K is constant for a system at constant
temperature.
(3) Relationship between Kp and Kc.
(4) Calculation of Kp and Kc from given set
of data.
(5) Difference between homogeneous and
heterogeneous equilibrium systems.

9.0 CHEMISTRY OF CARBON

COMPOUNDS
(a) Separation and Purification. Other methods should include solvent extraction
and melting point determinations.

(b) Determination of empirical and Outline of steps in:

molecular formulae. (a) Detection of N, S and the halogens.
(b) Estimation of C, H and O.

(c) Reactivity of Organic Compounds. (1) Inductive effect and Mesomeric effect.
(2) Resonance illustrated with benzene
molecule.
(3) Explanation of the terms:
nucleophiles, electrophiles, free radicals and
ions. homolytic fission, heterolytic fission.
(d) Alkanes
Halogenation – free radical mechanism.

(e) (i) Reactions of benzene. Mono substituted reactions of benzene: toluene,

phenol, aniline, benzoic acid and nitrobenzene.
(IUPAC and trivial names)

(ii) Comparison or reactions of Differences between the reactivity of benzene and

benzene and alkenes. alkenes towards certain reagents.
Uses of hexachlorocyclobezane and benzene
hexachloride (BHC).

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CONTENT NOTES

10.0 CHEMICAL INDUSTRY AND

ENVIRONMENT

(a) (i) Sources of raw materials Location of mineral deposits and their nature.
(ii) Mining of mineral as ore.

(iii) Extraction of metals Mineral (1) Metals – gold, bauxite, manganese and iron.
deposits in Ghana. (2) Precious stone – diamond.
(3) Industrial mining of limestone CaCO3, clay
Kaolin, solar salt
(4) Processing of Au, Al, Fe as main products
(5) Uses of the metals

(b) Cement and its uses (1) Sources of raw materials for cement
sproduction.
(2) Processes involved in the production of
cement.
(3) Uses of cement.
(4) Environmental impact.

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SECTION C

(For candidates in Nigeria, Sierra-Leone, Liberia and The Gambia)

CONTENT NOTES

1.0 NON METALS AND THEIR

COMPOUNDS

(a) Carbon:
(i) Allotropes of carbon; (1) Graphite, diamond and amorphous
Carbon;
(2) Structures, properties and uses.
(3) The uses of the allotropes should be
correlated with their properties and
structures.
(4) Combustion of allotropes.

(ii) Coal:
I. Types; Different types should include anthracite,
peat and lignite.
II. Destructive distillation
of coal and uses of the
products.

(iii) Coke:
I. Classification and uses; Water gas and producer gas.
II. Manufacture of synthetic
gas and uses.

(iv) Oxides of carbon

I. Carbon (IV) oxides; (1) Laboratory preparation.
(2) Properties and uses.
(3) Test for carbon (IV) oxides.

II. Carbon (II) oxides; Properties and uses only.

III. Trioxocarbonate (IV) salt. (1) Properties: solubility, action of heat,

reaction with dilute acid.

(2) Uses.

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CONTENT NOTES

(b) Oxygen:
(i) Laboratory and industrial
preparation;
(ii) Properties and uses; Test for oxygen will be required.

(iii) Binary compounds of oxygen:

acidic, basic, amphoteric and
neutral oxides.

(c) Hydrogen:
(i) Laboratory preparations;
(ii) Properties and uses. Test for hydrogen will be required.

(d) Water and solution: Test for water will be required.

(i) Composition of water; Reference should be made to the
electrolysis of acidified water.

(ii) Water as a solvent;

(iii)Hardness of water, causes and (1) Advantages and disadvantages of hard
methods of removing it; water and soft water.
(2) Experiments to compare the degrees of
hardness in different samples of water.
(iv) Treatment of water for town
supply.

(e) Halogens: Redox properties of the elements;

displacement reaction of one halogen by
another.
(i) Chlorine:
I. Laboratory preparation;
II. Properties and reactions. Properties should include:
(a) variable oxidation states;
(b) reaction with water and alkali
(balanced equation required).
(f) Hydrogen chloride gas:
(i) Laboratory preparation;
(ii) Properties and uses; (1) Test for HCl gas.
(2) Fountain experiment.

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CONTENT NOTES

(iii) Uses of halogen compounds. Uses should include silver halide in

photography and sodium oxochlorate (I) as
a bleaching agent.
(g) Nitrogen:
(i) Preparation and properties; Both laboratory and industrial preparations
from liquefied air are required.
(ii) Uses of nitrogen;
(iii) Compounds of nitrogen:
I. Ammonia; (1) Laboratory and industrial preparations.
(2) Properties and uses.
(3) Test for ammonia.
(4) Fountain experiment.

II. Trioxonitrate (V) acid; (1) Laboratory preparation.

(2) Properties and uses.

III. Trioxonitrate (V) salts. (1) Action of heat will be required.

(2) Test for trioxonitrate (V) ions.

(h) Sulphur:
(i) Allotropes and uses;
(ii) Compound of sulphur;
(iii) Trioxosulphate (IV) acids and
its salts;
(iv) Tetraoxosulphate (VI) acid: Contact process should be discussed.
industrial preparation, reactions
and uses.

(i) The noble gases: properties and uses.

2.0 METALS AND THEIR COMPOUNDS (1) Raw materials, processing, main
(a) Extraction of metals: products and by-products.

(i) Aluminium; (2) Uses of metals.

(ii) Iron;
(iii) Tin.

(b) Alloys. Common alloys of Cu, Al, Pb, Fe, Sn

and their uses.

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CONTENT NOTES

(c) Properties and uses of sodium and its

compounds. Compounds must be limited to NaCl,
NaOH, Na2CO3, NaNO3, Na2SO4 and
NaClO.
(d) Properties and uses of calcium and its The compounds must be limited to
compounds. CaCO3, CaO, CaSO4, CaCl2, and
Ca(OH)2
(e) Reactivity of iron and aluminium with
air, water and acids.
(f) Properties and uses of copper and its
The compounds must be limited to
compounds.
CuSO4, CuO and CuCl2.

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.16.0 PRACTICALS
(a) GENERAL SKILLS AND PRINCIPLES
Candidates will be expected to be familiar with the following skills and principles:
(i) Measurement of mass and volume;
(ii) Preparation and dilution of standard solutions;
(iii) Filtration, recrystallisation and melting point determination;
(iv) Measurement of heats of neutralization and solutions;
(v) Determination of pH value of various solutions by colorimetry;
(vi) Determination of rates of reaction from concentration versus time curves;

(vii) Determination of equilibrium constants for simple system.

(b) QUANTITATIVE ANALYSIS

Acid-base titrations

The use of standard solutions of acids and alkalis and the indicators; methyl orange,
methyl red and phenolphthalein to determine the following:

(i) The concentrations of acid and alkaline solutions;

(ii) The molar masses of acids and bases and water of crystallization.
(iii)The solubility of acids and bases;
(iv) The percentage purity of acids and bases;
(v) Analysis of Na2CO3/NaHCO3 mixture by double
indicator methods (Ghanaians only).
(vi) Stoichiometry of reactions.

Redox titrations
Titrations of the following systems to solve analytical problems:
(i) Acidic MnO4- with Fe2+;
(ii) Acidic MnO4- with C2O42-;
(iii) I2 in KI versus S2O32-.

(d) QUALITATIVE ANALYSIS

No formal scheme of analysis is required.

(i) Characteristic tests of the following cations with dilute NaOH(aq) and NH3(aq);

NH4; Ca2+; Pb2+; Cu2+; Fe2+; Fe3+; Al3+; and Zn2+.

(ii) Confirmatory tests for the above cations.

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(iii) Characteristic reaction of dilute HCl on solids or aqueous solutions and conc.
H2SO4 on solid samples of the following:
Cl- ; SO32- ; CO32- ; NO3- and SO42-.

(iv) Confirmatory tests for the above anions

(v) Comparative study of the halogens; displacement reactions.

(vi) Characteristic tests for the following gases: H2; NH3; CO2; HCl and SO2.

(vii) Characteristic test tube reactions of the functional groups in the following simple
organic compounds: Alkenes; alkanols; alkanoic acids, sugars (using Fehiling’s
and Benedict’s solutions only); starch (iodine test only) and proteins (using the
Ninhydrin test, Xanthoporteic test, Biuret test and Millon’s test only).

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