NABTEB Chemistry Syllabus For 2022

• Physical and Chemical

changes

• Definition of Elements,

Compounds and

Mixtures.

• Methods of separation

of Mixtures.

Experimental illustration of the methods of separation of mixtures is required.

1.1 Explain the concepts of atoms, molecules and ions (atomic or molecular ions)

1. 1.2  Explain the features of the atom.
2. 1.3  Draw and label the main electronic shells of the atom and explain atomic- number, isotopes, relative atomic mass (Ar), relative molecular mass (Mr) and calculate the relative atomic masses of elements that exhibit isotopy.
3. 1.4  List the main and sub- energy levels in an atom and explain the arrangement of electrons in these energy levels.

1. Concepts of Atoms, Molecules and Ions: Definition and treatment of particles as building block of matter.
2. Gross features of the Atom:
   • Account of Dalton’s Atomic theory.
   • Outline of the J.J. Thomson’s Experiment and Bohr-Rutherford’s Alpha Particles Scattering Experiment to establish the structure of the atoms.
3. Electronic Shells.
4. Atomic number, Mass

number, Isotopes, Relative atomic number (Ar), Relative molecular mass (Mr).

1. Electronic Energy Levels:

Ӣ Arrangement of

electrons in the main

and sub-energy levels. 6. Orbitals.

Ӣ Originofs,p,d,andf orbitals as sub-energy levels.

”¢ Shapes of s and p ”“

Arrangement of electrons in the main shells (K,L,M) and calculation of relative atomic mass of chlorine as example is required.

1. 1.5  Discuss the orbitals of atom and explain the rules for filling electrons in the sub-energy levels.
2. 1.6  Distinguish between chemical reaction and nuclear reaction.
3. 1.7  Explain the nature of the three main types of radiation.
4. 1.8  Define radioactivity. Distinguish between natural and artificial. Explain nuclear fission and nuclear fusion.
5. 1.9  Define ”˜Half-life’ as a measure of the stability of the atomic nucleus.

1.10 Write and balance simple nuclear equation. 1.11 Explain the effects of

radioactive radiation on human beings and state the uses of radio isotopes.

• Aufbau Principle,

Hund’s rule of Maximum Multiplicity and Pauli’s Exclusion Principle.

• Abreviated and detailed electronic configuration in terms of s, p and d orbitals from hydrogen to zinc.

7. Nuclear Chemistry:

• Chemical reaction and

nuclear reaction.

• Balancing of simple

nuclear equation.

• Types and nature of

radiations (alpha, Beta

and Gamma rays).

• Natural and artificial

radioactivity

• Nuclear fission and

nuclear fusion.

• Devices used for

detecting

radioactiveities.

• Qualitative treatment of

hald life.

• Effects and

applications of radioactivity, carbon dating, uses in agriculture, medicine and industry.

Symbols too, should be used to identify the type of radiation. Geiger Muller counter of detecting radiation should be described.

1.1 State the periodic law and its application in the formulation of the periodic table.

1. 1.2  Outline the periodic properties and their trends across the period and down a group.
2. 1.3  Outline the periodic graduation of properties

configurations leading to group and periodic classifications.

Ӣ The uniqueness of hydrogen atom in the periodic table.

2. Periodic properties of the first 30 elements, atomic size, ionization energy,

i. Metallic to

1. Periodic graduation of the

properties of halogens.

1. • Physical states, melting

points.

1. • Redox properties,

displacement reaction of one halogen by another.

1. Transition metals and their characteristic properties: Electronic configuration, Metallic properties, Chemical reactivities, Magnetic properties, variable oxidation states, formation of complex ions and catalytic properties.

4.1 Explain chemical bonding and list the types of chemical combinations.

1. 4.2  Illustrate with appropriate examples electrovalent or ionic bonding. List factors that influence the formation of electrovalent compound and outline the properties of an electrovalent or ionic compound.
2. 4.3  Illustrate with appropriate examples, ordinary covalent and

1. Chemical bonding and types of chemical combination.
2. Inter-atomic bonding:
   a. ionic bonding factors.
   * ionization energy, electron affinity and electro negativity differences. Properties: Melting points, boiling points and solubility in various solvents.
   b. Covalent bonding factors: ionization energy, electron affinity and electron negativity differences. Properties such as melting points,

4.4 Draw the shape of simple molecules.

1. 4.5  Explain metallic bonding. List factors that influence the formation of metallic bonds and outline the typical properties of metals..
2. 4.6  Explain vander Waal’s forces. Give examples of the compound in which they exist and explain the variation of melting points and boiling points of noble gases; halogen and alkenes in terms of vander Waal’s forces.
3. 4.7  Explain the unusually higher boiling point of HF, H20 and NH3 over HCL, H2S and PH3 respectively due to hydrogen bonding.

solvents.

1. Simple molecules and their

shapes i.e. (i) Linear (ii) Non-linear (iii) Tetrachedral.)

1. Metallic bonding: Factors such as atomic radius, ionization potential, and number of valence electrons. Properties to include conductivity, malleability, ductility.
2. Intermolecular bonding.
   a. vander Waal’s forces: relative physical properties of polar and non-polar compounds.

b. Hydrogen bonding.
* Variation in the boiling points of compounds such as H20, H2S

These should be demonstrated using metals such as Mg, Zn, Sn and Fe.

Description of formation and nature should be treated. Dipole-dipole and induced dipolar forces is required.

formulae and equations. List the rules for writing of balanced equations and write balanced chemical equations by applying the rules.

1. 5.2  State and explain the laws of chemical combination.
2. 5.3  Explain the concepts of moles, Avogadro number and constant molar

1. Symbols, formulae and equations.

Ӣ Chemical symbols. Ӣ Empirical and

molecular formulae. Ӣ Chemical equations. Combining power of

elements and oxidation

numbers.

1. Laws of chemical

combination.

1. • Law of conservation of

mass.

1. • Law of constant

composition.

5.4 Explain the concept of solution.

3. Amount of substance:

• mass and volume

measurements.

• Avogadro constant. L

as the number of carbon atoms in 1 mole (12.00g of 12C).

• Molar quantities and their uses.
• Mole of electrons, atoms, molecules, formula units.
• Mole ratio.
• Calculation of mass

concentration and molar concentration volume and other quantities in chemical reactions.

4. Solutions:

• concept of solution as

made up of solvent and solute (in a single phase).

• Concentration terms.
• Standard solutions.

Preparation of some primary standard solutions using anhydrous Na2CO3, (COOH), 2H2O

• dilution factor determination.

1.1 Define And explain the concept if law, theory and hypothesis as usual in science.

1.2 State the postulate of Kinetic theory of matter. Explain the nature of solids, liquids and gasses, change of state of matter

• Postulates of the kinetic

theory.

• Nature of solids,

liquids and gases.

• Change of states of

matter.

• Diffusion,

demonstration using

kinetic model.

1. 1.3  State and explain the gas

laws. Explain each law using the kinetic model. Represent the laws mathematically and graphically (where applicable). Derive the general gas lwaw.

1. 1.4  State the relationship of vapour pressure with the boiling points of liquids. Describe simple methods for the determination of boiling points.
2. 1.5  Classify solids. Compare the properties of the types of solids. Describe the arrangement of ions, molecules and atoms in three dimensions in the solid state.

Explain melting points and describe the structure, properties, and uses of diamond and graphite.

1. Gases:
   The gas laws. Charles’ Boyle’s Dalton’s, Graham’s Avogadro’s law and ideal gas equation. Mathematical relation of the gas laws and calculations. Molar volume of a gas at S.T.P. = 22.4dm3.
   PV = K

T
Derivation of general gas law and calculations.

1. Liquids:
   • Liquids as an

intermediate state between gases and solids in the kinetic molecular sense.

1. • Concept of vapour pressure: simple determination of boiling points.
   • Standard boiling points.
2. Solids:

Types and structure.

1. • Ionic, metallic,

covalent and molecular

solids

1. • Comparism of the

properties of the types

of solids.

1. • Regular arrangement of

ions, molecules and atoms in three dimensions in the solid state.

1. • Melting points.

Pollen grain/sulphur in water (viewed under microscope).

Smokes in glass container illuminated by a strong light from the side.

A dusty room being swept and viewed from outside under sunlight is required.

Teachers should endeavour to point out the differences between Dalton’s law and dalton’s atomic theory.

Specific packing arrangement knowledge is NOT required.

7.1 Define energy. List different forms of energy. State the laws of conservation of energy and explain its units.

1. 7.2  Explain the term exothermic reactions, endothermic reactions, heat of reaction, heat of formation, heat of combustion, heat of neutralization and heat of solution. Measure and calculate heat of neutralization, heat of solution and heat of combustion.
2. 7.3  Concepts of free energy and entropy. Discussion

of ˆ†G=ˆ†H-Tˆ†S

• Exothermic and

endothermic processes.

• Total energy of a

system as the sum of various forms of energy e.g. kinetic, potential, electrical, heat, sound, etc.

• Enthalpy changes of the following:

Formation, combustion,

solution, neutralization.

• Practical knowledge of the measurement of the heats of neutralization,

solution and

combustion.

• Uses of energy changes

Ӣ Conditions for spontaneous changes as consequences of balance between tendency towards lower enthalpy and tendency toward higher entropy.

1.1 Define acid, base and salt. Explain Arrhenius theory of acids and basicity of an acid.

1. 1.2  Outline the various physical and chemical properties and chemical properties of acids, bases and salts. Balance chemical equations of ionic reactions.
2. 1.3  Explain the preparation of acids and salts by various methods.
3. 1.4  Define electrolyte and non-electrolyte. Distinguish between strong and weak electrolytes.
   Determine the conductances and enthalpy of neuralization of acids, bases and salts.
4. 1.5  Explain the pH scale. Use it to determine the acidity and alkalinity of aqueous solutions.
5. 1.6  Explain the behaviour of weak acids and bases in water. Compare the conductances of molar solutions of strong and weak acids and bases.
6. 1.7  Explain hydrolysis and

1. Acids, bases and salts: Ӣ Definition.
   • Arrhenius concept of acids and bases.
   • Basicity of and acid
2. Physical and chemical properties of acids and

bases:

1. • Conductivities, taste

etc.

1. • Concept of

amphoterism

1. • Balanced chemical

equations of all

reactions.

1. Preparation of acids and

salts:
Ӣ Deliquescent

efflorescent and hygroscopic substances.

1. Acids, bases and salts as electrolytes.
   • Electrolytes and non- electrolytes, strong and weak electrolytes.
   • Evidence from conductivity and enthalpy of neutralization.
2. The pH
   • knowledge of the pH

scale.

1. • as a measure of

acididity and alkalinity

salts.

1. 1.8  Explain the elementary

theory of indicators: State the working oH ranges of methyl orange and phenolphthalein.

1. 1.9  Identify correctly the relevant apparatus for acid-base titrations. Carry out titrations using acids, bases and appropriate indicators.

pH and poH from given

data

1. Weak acids and weak

bases:

1. • Behaviour of acids and

bases in water as an example of equilibrium system.

1. • Comparison of the conductances of molar solutions of strong and weak acids and bases.
2. Hydrolysis: Qualitaive
   • Explanation of the

hydrolysis

1. • Behaviour of NH4CI

AICI3, CuSO4, Na2CO3H, Na2S, CuNO3 CH3COONa in water. Compare with NaCI, CaCI2, Ba(NO3)2, K2SO4.

1. Acid-base indicators:
   • Indicators as weak

organic acids and

bases.

1. • Colour of indicator at

any pH dependent on relative amounts of acid and base forms.

1. • Working ranges of methyl orange and phenolphthalein
2. Acid-base titrations:
   • Relevant apparatus for

acid-base titrations.

1. • Workings of indicators

in acid-base titrations.

1. • Determine

concentrations, % purity, water of crystallization from the experimental result.

Teacher is advised to approach this practically, testing the various solutions with litmust paper or litmus solution.

9.1 Explain the general principles of solubility.

1. Define solution, saturated solution, super saturated solution and solubility.
2. Factors affecting solubility
3. Determination of

solubilities of substance.

1. Solubility curves and their

applications.

1. 1.1  Define rate of reaction. Explain factors

affecting rates of reactions. Discuss the theory of reaction rates.

1. 1.2  Explain the general principles of

equilibrium, Le Chatelier’s principle and factors affecting positions\ of equilibrium in chemical reactions.

1. Rates of reaction:
   • definition of the rates

of reaction for gaseous systems: Pressure may be used as concentration terms.

1. • Collision theory and activation energy theory.
   • Factors influencing collision such as temperature and concentration, surface area/nature of reactions.
   • Concentration time graph
   • Effective collision.
   • Activation energy
   • Energy profile

diagrams showing activation energy and enthalpy change.

1. Equilibrium systems:

Ӣ Reversible and

irreversible reactions.

• Meaning of equilibrium constant and its mathematical expression.
• Statement of Le Chatelier’s principles.
• Factors affecting the position of equilibrium of chemical reactions.

1.1 Explain the concepts of oxidation and

reduction, reducing and oxidizing agents, redox reactions. Outline the rules for the determination of oxidation numbers of elements in substances.

1.2 Describe electrochemical cells and outline their applications.

1. Definitions of oxidation and reduction, reducing and oxidizing agents in terms of:
   • Addition and removal of oxygen and

hydrogen.

1. • Loss and gain of

electrons.

1. • Change in oxidation

number.

1. Oxidation numbers/states
2. Tests for oxidants and

reductants.

1. Balancing of redox

equations by:

1. • ion, electron, or change

in oxidation

number/state.

1. • Half reactions and

overall reactions.

1. • Categorising of processes at the

electrodes.

1. Standard electrode

potentials of

electrochemical cells.

1. Drawing and writing of the

cell diagrams.

1. Electromotive force of

cells.

1. Distinction between

primary and secondary

(b) the same reaction e.g. disproportion IUPAC system of nomenclature is required.

1. 1.3  Explain the principles/mechanisms of electrolysis. Discuss the factors that determine the preferential discharge of ions at the electrode and their practical applications.
2. 1.4  Explain the concept of corrosion on metals.

1. Principles of electrolysis
2. Comparison of the

mechanism of electrolysis

with electrochemical cells

1. Faraday’s Law
2. Practical applications

1. Corrosion treated as redox process.
2. Rusting of iron and its economic implication.
3. Prevention based on relative magnitude of electrode potentials, preventive methods like galvanizing, sacrificial cathodic protection, and non-redox methods.

1. Oxygen: Laboratory and industrial preparation, properties, uses and tests, binary compounds of oxygen: acidic oxides, basic oxides, amphoteric oxides, neutral oxides and higher oxides..
2. Hydrogen: Laboratory and industrial preparation, properties, uses and tests, isotopes of hydrogen.
3. Water and solution: Composition of H2O2 hardness of water, properties and test.

The contributions of 2H the isotopes and to the making of 3H heavy water and hydrogen bomb respectively. Reference should be made to acidulated water.

1. Nitrogen: Laboratory and industrial preparation, properties and uses, compounds of nitrogen:
   • NH3 laboratories/

industrial preparation and uses.

1. • HNO3: Laboratory preparation, reactions and uses.
   • NO2: Laboratory/ industrial

preparation,

properties and uses.

1. Carbon: Allotropes,

general properties, CO & CO2 reparation, properties and uses.

1. Sulphur: Allotropes and uses. Compound of sulphur; sulphides, trioxosulphate (IV) acid and salt, tetraoxosulphate (VI) acid ”“ industrial preparation, reactions and uses.

8. Noble gases: Properties and uses.

Fountain experiment to illustrate solubility of HCL and NH3 Demonstration of the oxidation of Halides to other halogen by chlorine.

Coal ”“ different types and destructive distillation is required.

Candidates are expected to appreciate the fact that some noble gases do form compound when specially treated with fluorine.

1. Distinction between metals and non-metals. Relative abundance of metals/metallic ors in nature.
2. Principle of extraction of metals. Preliminary preparations.
   Ӣ Electrolysis
   • Reduction of chlorides.
   • Reduction of

oxides/sulphides.

1. • Thermal and chemical

reduction

1. Extraction and uses of:

1. 1. Alkali metals ”“ lithium, sodium and potassium. Alkaline earthmetals.
   2. Ca, Mg, and AL
   3. transition of metals ”“

Fe, Cu, Ag and Au

1. 1. Important compounds of Na, K, Mg, Ca, A1,

Ba, Cu, Ag, Zn, Hg, Fe

and their uses.

1. Alloys of Fe, Cu, AL Pb

and Zn

1. Activity series

The uses of alkali metals as precipitating agents in cation analysis should be stressed.

1. Classification and nomenclature.

Root names.

Functional groups

1. Separation and purification

methods ”“ distillation, crystallization, drying, chromatography.

1. Determination of the empirical and molecular formulae and the molecular structure.
2. Homologous series.
3. Isomerism:
4. Differences between structural geometric/ stereoisomerism.

Example should be limited to compounds having maximum if five carbon atoms.

1.1 Explain the major classification of

hydrocarbons. Write the general and structural formulae and identify their functional groups.

1.2 Explain the sources, properties and uses of the alkanes.

1.3 Describe and explain the process of refining petroleum.

1. Definition of

hydrocarbons.

1. Sources of

hydrocarbons

• coal.
• Natural gas and petroleum.

c. Classifications of hydrocarbons.

• Aliphatic.
• Cyclic”“aliphatic and

aromatic.

1. Classifications of

aliphatic hydrocarbons.

1. Distinction between

saturated and unsaturated hydrocarbons.

1. General and structural formulae of alkanes, alkenes and Alkynes.
2. Identification of the functional groups.

2. Alkanes:

• Laboratory and

industrial preparation.

• Nomenclature and

structure.

• Reactivity:

Combustion, substitution reactions, cracking of large alkane molecules.

• Uses of alkanes: As fuels, starting materials for synthesis.
• Isomerism in alkanes. 3. Petroleum:

Ӣ Composition.

• Fractional distillation

and major products.

• Cracking process and

reforming.

• Petrochemicals.
• Octane numbr and

The uses of haloalkanes and pollution effects should be treted.

1.5 Explain the sources, properties and uses of alkynes.

1.6 Describe the structure and properties of the benzene (both physical and chemical).

4. Alkenes:

• Laboratory preparation.
• Nomenclature and

structure.

• Main reactions such as

addition reactions with halogens, bromine water, hydrogen halides: Oxidation: Hydroxylation with aqeous KMnO4

• Laboratory detection: Use of reaction with Br2.CCI4 as a means of characterizing alkenes.

5. Alkynes:

• Laboratory preparation

or production.

• Nomenclature and

structure.

• Industrial uses of

ethyne e.g. oxyacetylene (oxyethyne) in lamps.

6. Benzene:

• Resonance in benzene.
• Halogenations.
• Addition reactions.
• Comparison of its

reaction with those of Alkenes.

16.1 Explain the sources, nomenclature, structure, classification, properties and laboratory test of alkanas.

1. General formula of alkanols as CnH2+1OH.
2. Functional group of alkanols as OH.
3. Molecular and structural formula and the IUPAC names of the first vew,
4. General methods of preparation.
5. Fermentation process as a method of preparation and from ethene a, a by product of the cracking process.
6. Classification of alkanols. Ӣ Primary

Ӣ Secondary

Ӣ Tertiary

1. Physical and chemical

properties.

1. Uses of ethanol
2. Test

• Functional groups as

COOH.

• Molecular and

structural formula and IUPAC names of the first few members.

• Physical and chemical properties.
• Uses and properties of ethanoic and phenyl methanoic (benzoic) acids as examples of aliphatic and aromaic acids respectively.
• Laboratory est.
• Reaction with NaHCO3 2. Alkanoates:

• Preparation of alkyl alkanoates (esters) from alkanoic acids.
• Physical and chemical

Acidic properties should be emphasized.

Ӣ Saponification, hardening of oils.

• Detergents.
• Comparison of soapy

and soapless detergents with respect to their action with soft and hard water. respectively.

Laboratory preparation of soapy detergents and soapless detergents.

1.1 Explain the source, properties and

important uses of some macro-molecules.

1. Amino acids:
   * Difunctional nature of amino acids.
2. Natural and synthetic polymers:

Ӣ Definition

Ӣ Types of polymers.

Natural polymers:

1. Carbohydrates:

formulae, properties,

classification and uses.

1. Proteins: Polymers of

amino molecules linked by peptides or amide linkage.

1. Hydrolysis uses in living system.

Synthetic polymers: Classification and monomers and co- polymers.

Experimental differentiation of reducing and non reducing of sugars should be carried out.

• History of the

development of the

chemical industry.

• Important chemical

industries in Nigeria, and their corresponding raw materials.

• Distinction between fine and heavy

chemicals.

• Factors influencing the

siting of chemical

industries.

• Effects of industries on

the community.

2. Extraction of metals:

• Raw materials, processing, main products, by-products, recycling of AL, and Fe, Au or Sn.
• Uses of metals.
• Define and list the

common alloys: Cu, AL, Pb, and Fe and their uses.

• Sources, effects and

control.

• Green house effect and

depletion of the ozone

layer.

• Biodegradable

pollutants.

4. Biotechnology:

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