CHLORINE AND ITS COMPOUNDS

Introduction and Preparation of Chlorine
Physical Properties of Chlorine
Chemical Properties of Chlorine - Reaction with Water

By the end of the lesson, the learner should be able to:
Define chlorine and state its position in the periodic table. Describe the occurrence of chlorine in nature. Describe laboratory preparation of chlorine gas. Write balanced equations for chlorine preparation.
Investigate the physical properties of chlorine gas. Explain the method of collection used for chlorine. Test the solubility of chlorine in water. State the density and color of chlorine gas.

Q/A: Review Group VII elements and electron configuration of chlorine ( 8.7). Discussion: Occurrence as sodium chloride in sea water and rock salt. Practical work: Experiment 6.1 - Preparation using MnO2 + concentrated HCl. Setup apparatus as in Figure 6. Safety precautions for handling chlorine gas.
Practical work: Experiment 6.2 - Testing chlorine gas preserved from previous experiment. Recording observations in Table 6. Testing: Color, smell (caution - no direct smelling), density, solubility in water. Demonstration: Inverting gas jar in water trough. Discussion: Why collected by downward delivery.

Manganese(IV) oxide, Concentrated HCl, Gas collection apparatus, Water, Concentrated H2SO4, Blue litmus paper, Gas jars
Preserved chlorine gas, Water trough, Gas jars, Observation tables, Safety equipment
Chlorine gas, Distilled water, Blue and red litmus papers, Colored flower petals, Gas jars, Boiling tubes

KLB Secondary Chemistry Form 4, Pages 195-196
KLB Secondary Chemistry Form 4, Pages 196-197

CHLORINE AND ITS COMPOUNDS

Chemical Properties of Chlorine - Reaction with Metals
Chemical Properties of Chlorine - Reaction with Non-metals
Oxidising Properties of Chlorine

By the end of the lesson, the learner should be able to:
Investigate reactions of chlorine with metals. Write balanced equations for metal-chlorine reactions. Explain the formation of metal chlorides. Demonstrate exothermic nature of these reactions.

Practical work: Experiment 6.4 - Reactions with burning magnesium, hot iron wire, dry chlorine over hot iron coil (Figure 6.2). Recording observations in Table 6. Observations: White fumes (MgCl2), glowing iron wire, black crystals (FeCl3). Discussion: Formation of higher oxidation state chlorides. Safety: Proper ventilation and eye protection.

Magnesium ribbon, Iron wire, Chlorine gas, Deflagrating spoon, Combustion tube, Anhydrous CaCl2, Gas jars
Red phosphorus, Hydrogen gas, Chlorine gas, Deflagrating spoon, Gas jars, Bunsen burner, Safety equipment
Sodium sulphite solution, Barium nitrate, Lead nitrate, Hydrogen sulphide gas, Aqueous ammonia, Chlorine gas, Test tubes

KLB Secondary Chemistry Form 4, Pages 199-201

CHLORINE AND ITS COMPOUNDS

Reaction of Chlorine with Alkali Solutions
Oxidising Properties - Displacement Reactions

By the end of the lesson, the learner should be able to:
Investigate reactions of chlorine with alkalis. Compare reactions with cold dilute and hot concentrated alkalis. Write equations for formation of chlorates and hypochlorites. Explain formation of bleaching powder.

Practical work: Experiment 6.7 - Bubbling chlorine through cold dilute NaOH and hot concentrated NaOH. Recording observations in Table 6. Formation of pale-yellow solution (cold) vs colorless solution (hot). Equations: 3Cl2 + 6NaOH → 5NaCl + NaClO3 + 3H2O (hot), Cl2 + 2NaOH → NaCl + NaClO + H2O (cold). Discussion: Industrial production of bleaching powder.

Sodium hydroxide solutions (dilute cold, concentrated hot), Chlorine gas, Beakers, Bunsen burner, Thermometer
Potassium bromide solution, Potassium iodide solution, Chlorine gas, Test tubes, Observation charts

KLB Secondary Chemistry Form 4, Pages 202-203

CHLORINE AND ITS COMPOUNDS

Test for Chloride Ions
Uses of Chlorine and its Compounds

By the end of the lesson, the learner should be able to:
Carry out confirmatory tests for chloride ions. Distinguish between different chloride tests. Practice qualitative analysis techniques. Write equations for chloride ion tests.

Practical work: Experiment 6.9 - Testing sodium chloride with concentrated H2SO4, testing with lead(II) nitrate solution. Recording observations in Table 6. Tests: White fumes with H2SO4 + ammonia test, white precipitate with Pb(NO3)2 that dissolves on warming. Writing equations: NaCl + H2SO4 → NaHSO4 + HCl, Pb²⁺ + 2Cl⁻ → PbCl

Sodium chloride, Concentrated H2SO4, Lead(II) nitrate solution, Aqueous ammonia, Glass rod, Test tubes, Bunsen burner
Charts showing industrial uses, Samples of bleaching agents, PVC materials, Photographs of water treatment plants, Industrial application diagrams

KLB Secondary Chemistry Form 4, Pages 204-205

CHLORINE AND ITS COMPOUNDS

Hydrogen Chloride - Laboratory Preparation
Chemical Properties of Hydrogen Chloride

By the end of the lesson, the learner should be able to:
Describe laboratory preparation of hydrogen chloride gas. Set up apparatus for HCl preparation. Investigate physical properties of HCl gas. Explain the method of collection used.
Prepare aqueous hydrogen chloride (hydrochloric acid). Investigate acid properties of HCl solution. Test reactions with metals, bases, and carbonates. Compare HCl in water vs organic solvents.

Practical work: Experiment 6.10 - Preparation using rock salt (NaCl) + concentrated H2SO Setup apparatus as in Figure 6.3(b). Testing physical properties and recording in Table 6.6. Tests: Solubility (fountain experiment), reaction with ammonia, effect on litmus. Collection by downward delivery due to density. Writing equation: NaCl + H2SO4 → NaHSO4 + HCl.
Practical work: Experiment 6.11 - Preparation of aqueous HCl using apparatus in Figure 6. Testing with metals (Zn, Fe, Mg, Cu), NaOH, carbonates, lead nitrate. Recording observations in Table 6.7. Testing HCl in methylbenzene - no acid properties. Discussion: Ionization in water vs molecular existence in organic solvents. Writing equations for acid reactions.

Rock salt (NaCl), Concentrated H2SO4, Gas collection apparatus, Ammonia solution, Litmus papers, Water trough, Gas jars
Distilled water, Filter funnel, Metals (Zn, Fe, Mg, Cu), NaOH solution, Carbonates, Lead nitrate, Methylbenzene, Indicators

KLB Secondary Chemistry Form 4, Pages 207-208
KLB Secondary Chemistry Form 4, Pages 208-211

CHLORINE AND ITS COMPOUNDS

Large-scale Manufacture of Hydrochloric Acid

By the end of the lesson, the learner should be able to:
Describe industrial production of hydrochloric acid. Identify raw materials and conditions used. Explain the controlled combustion process. Draw flow diagrams of the industrial process.

Study of Figure 6.4 - Large-scale manufacture setup. Discussion: Raw materials (H2 from electrolysis/cracking, Cl2 from electrolysis). Controlled combustion: H2 + Cl2 → 2HCl in jet burner. Dissolving HCl gas in water over glass beads. Safety: Explosive nature of H2/Cl2 mixture, use of excess chlorine. Industrial considerations: 35% concentration, transport in rubber-lined steel tanks.

Flow diagrams, Industrial photographs, Glass beads samples, Charts showing electrolysis processes, Safety equipment models

KLB Secondary Chemistry Form 4, Pages 211-212

CHLORINE AND ITS COMPOUNDS

Uses of Hydrochloric Acid
Environmental Pollution by Chlorine Compounds and Summary

By the end of the lesson, the learner should be able to:
List the industrial uses of hydrochloric acid. Explain applications in metal treatment. Describe use in water treatment and manufacturing. Relate acid properties to industrial applications.

Discussion: Applications - rust removal and descaling, galvanizing preparation, electroplating preparation, water treatment (chlorination), sewage treatment. Manufacturing uses: dyes, drugs, photographic materials (AgCl), pH control in industries. Q/A: How acid properties make HCl suitable for these uses. Case studies: Metal cleaning processes, water purification systems.

Samples of rusted and cleaned metals, Photographic materials, pH control charts, Industrial application videos, Water treatment diagrams
Environmental pollution charts, Ozone layer diagrams, DDT restriction documents, PVC waste samples, NEMA guidelines, Summary charts of reactions

KLB Secondary Chemistry Form 4, Pages 212-213

THE MOLE

Relative Mass - Introduction and Experimental Investigation

By the end of the lesson, the learner should be able to:
Define relative mass using practical examples
Compare masses of different objects using a reference standard
Explain the concept of relative atomic mass
Identify carbon-12 as the reference standard

Experiment: Weighing different sized nails using beam balance. Use smallest nail as reference standard. Q/A: Discuss everyday examples of relative measurements. Teacher exposition: Introduction of carbon-12 scale and IUPAC recommendations. Calculate relative masses from experimental data.

Different sized nails ( 5-15cm), Beam balance, Fruits of different masses, Reference charts

KLB Secondary Chemistry Form 3, Pages 25-27

THE MOLE

Avogadro's Constant and the Mole Concept
Interconversion of Mass and Moles for Elements

By the end of the lesson, the learner should be able to:
Define Avogadro's constant and its value
Explain the concept of a mole as a counting unit
Relate molar mass to relative atomic mass
Calculate number of atoms in given masses of elements
Apply the formula: moles = mass/molar mass
Calculate mass from given moles of elements
Convert between moles and number of atoms
Solve numerical problems involving moles and mass

Experiment: Determine number of nails with mass equal to relative mass in grams. Teacher exposition: Introduce Avogadro's constant (6.023 × 10²³). Discussion: Mole as counting unit like dozen. Worked examples: Calculate moles from mass and vice versa.
Worked examples: Mass-mole conversions using triangle method. Supervised practice: Calculate moles in given masses of common elements. Problem solving: Convert moles to atoms using Avogadro's number. Assignment: Practice problems on interconversion.

Beam balance, Various sized nails, Scientific calculators, Avogadro's constant charts
Scientific calculators, Periodic table, Worked example charts, Formula triangles

KLB Secondary Chemistry Form 3, Pages 27-30
KLB Secondary Chemistry Form 3, Pages 30-32

THE MOLE

Molecules and Moles - Diatomic Elements
Empirical Formula - Experimental Determination

By the end of the lesson, the learner should be able to:
Distinguish between atoms and molecules
Define relative molecular mass
Calculate moles of molecules from given mass
Determine number of atoms in molecular compounds

Discussion: Elements existing as molecules (O₂, H₂, N₂, Cl₂). Teacher exposition: Difference between atomic and molecular mass. Worked examples: Calculate moles of molecular elements. Problem solving: Number of atoms in molecular compounds.

Molecular models, Charts showing diatomic elements, Scientific calculators
Crucible and lid, Magnesium ribbon, Bunsen burner, Beam balance, Tongs, Safety equipment

KLB Secondary Chemistry Form 3, Pages 29-30

THE MOLE

Empirical Formula - Reduction Method

By the end of the lesson, the learner should be able to:
Determine empirical formula using reduction reactions
Calculate empirical formula from reduction data
Apply reduction method to copper oxides
Analyze experimental errors and sources

Experiment: Reduction of copper(II) oxide using laboratory gas. Measure masses before and after reduction. Calculate moles of copper and oxygen. Determine empirical formula from mole ratios. Discuss experimental precautions.

Combustion tube, Porcelain boat, Copper(II) oxide, Laboratory gas, Beam balance, Bunsen burner

KLB Secondary Chemistry Form 3, Pages 35-37

THE MOLE

Empirical Formula - Percentage Composition Method

By the end of the lesson, the learner should be able to:
Calculate empirical formula from percentage composition
Convert percentages to moles
Determine simplest whole number ratios
Apply method to various compounds

Worked examples: Calculate empirical formula from percentage data. Method: percentage → mass → moles → ratio. Practice problems: Various compounds with different compositions. Discussion: When to multiply ratios to get whole numbers.

Scientific calculators, Percentage composition charts, Worked example displays

KLB Secondary Chemistry Form 3, Pages 37-38

THE MOLE

Molecular Formula - Determination from Empirical Formula
Molecular Formula - Combustion Analysis
Concentration and Molarity of Solutions

By the end of the lesson, the learner should be able to:
Define molecular formula
Relate molecular formula to empirical formula
Calculate molecular formula using molecular mass
Apply the relationship (empirical formula)ₙ = molecular formula
Determine molecular formula from combustion data
Calculate moles of products in combustion
Relate product moles to reactant composition
Apply combustion analysis to hydrocarbons

Teacher exposition: Difference between empirical and molecular formulas. Worked examples: Calculate molecular formula from empirical formula and molecular mass. Formula: n = molecular mass/empirical formula mass. Practice problems with various organic compounds.
Worked examples: Hydrocarbon combustion producing CO₂ and H₂O. Calculate moles of C and H from product masses. Determine empirical formula, then molecular formula. Practice: Various combustion analysis problems.

Scientific calculators, Molecular mass charts, Worked example displays
Scientific calculators, Combustion analysis charts, Molecular models of hydrocarbons
Scientific calculators, Molarity charts, Various salt samples for demonstration

KLB Secondary Chemistry Form 3, Pages 38-40
KLB Secondary Chemistry Form 3, Pages 40-41

THE MOLE

Preparation of Molar Solutions

By the end of the lesson, the learner should be able to:
Describe procedure for preparing molar solutions
Use volumetric flasks correctly
Calculate masses needed for specific molarities
Prepare standard solutions accurately

Experiment: Prepare 1M, 0.5M, and 0.25M NaOH solutions in different volumes. Use volumetric flasks of 1000cm³, 500cm³, and 250cm³. Calculate required masses. Demonstrate proper dissolution and dilution techniques.

Volumetric flasks (250, 500, 1000cm³), Sodium hydroxide pellets, Beam balance, Wash bottles, Beakers

KLB Secondary Chemistry Form 3, Pages 43-46

THE MOLE

Dilution of Solutions

By the end of the lesson, the learner should be able to:
Define dilution process
Apply dilution formula M₁V₁ = M₂V₂
Calculate concentrations after dilution
Prepare dilute solutions from concentrated ones

Experiment: Dilute 25cm³ of 2M HCl to different final volumes (250cm³ and 500cm³). Calculate resulting concentrations. Worked examples using dilution formula. Safety precautions when diluting acids.

Volumetric flasks, Hydrochloric acid (2M), Measuring cylinders, Pipettes, Safety equipment

KLB Secondary Chemistry Form 3, Pages 46-50

THE MOLE

Stoichiometry - Experimental Determination of Equations

By the end of the lesson, the learner should be able to:
Determine chemical equations from experimental data
Calculate mole ratios from mass measurements
Write balanced chemical equations
Apply stoichiometry to displacement reactions

Experiment: Iron displacement of copper from CuSO₄ solution. Measure masses of iron used and copper displaced. Calculate mole ratios. Derive balanced chemical equation. Discuss spectator ions.

Iron filings, Copper(II) sulphate solution, Beam balance, Beakers, Filter equipment

KLB Secondary Chemistry Form 3, Pages 50-53

THE MOLE

Stoichiometry - Precipitation Reactions
Stoichiometry - Gas Evolution Reactions
Volumetric Analysis - Introduction and Apparatus

By the end of the lesson, the learner should be able to:
Investigate stoichiometry of precipitation reactions
Determine mole ratios from volume measurements
Write ionic equations for precipitation
Analyze limiting and excess reagents
Determine stoichiometry of gas-producing reactions
Collect and measure gas volumes
Calculate mole ratios involving gases
Write equations for acid-carbonate reactions

Experiment: Pb(NO₃)₂ + KI precipitation reaction. Use different volumes to determine stoichiometry. Measure precipitate heights. Plot graphs to find reaction ratios. Identify limiting reagents.
Experiment: HCl + Na₂CO₃ reaction. Collect CO₂ gas in plastic bag. Measure gas mass and calculate moles. Determine mole ratios of reactants and products. Write balanced equation.

Test tubes, Lead(II) nitrate solution, Potassium iodide solution, Burettes, Ethanol, Rulers
Conical flask, Thistle funnel, Plastic bags, Rubber bands, Sodium carbonate, HCl solution
Pipettes (10, 20, 25cm³), Burettes (50cm³), Pipette fillers, Conical flasks, Various solutions

KLB Secondary Chemistry Form 3, Pages 53-56
KLB Secondary Chemistry Form 3, Pages 56-58

THE MOLE

Titration - Acid-Base Neutralization

By the end of the lesson, the learner should be able to:
Perform acid-base titrations accurately
Use indicators to determine end points
Record titration data properly
Calculate average titres from multiple readings

Experiment: Titrate 25cm³ of 0.1M NaOH with 0.1M HCl using phenolphthalein. Repeat three times for consistency. Record data in tabular form. Calculate average titre. Discuss accuracy and precision.

Burettes, Pipettes, 0.1M NaOH, 0.1M HCl, Phenolphthalein indicator, Conical flasks

KLB Secondary Chemistry Form 3, Pages 59-62

THE MOLE

Titration - Diprotic Acids

By the end of the lesson, the learner should be able to:
Investigate titrations involving diprotic acids
Determine basicity of acids from titration data
Compare volumes needed for mono- and diprotic acids
Write equations for diprotic acid reactions

Experiment: Titrate 25cm³ of 0.1M NaOH with 0.1M H₂SO₄. Compare volume used with previous HCl titration. Calculate mole ratios. Explain concept of basicity. Introduce dibasic and tribasic acids.

Burettes, Pipettes, 0.1M H₂SO₄, 0.1M NaOH, Phenolphthalein, Basicity reference chart

KLB Secondary Chemistry Form 3, Pages 62-65

THE MOLE

Standardization of Solutions

By the end of the lesson, the learner should be able to:
Define standardization process
Standardize HCl using Na₂CO₃ as primary standard
Calculate accurate concentrations from titration data
Understand importance of primary standards

Experiment: Prepare approximately 0.1M HCl and standardize using accurately weighed Na₂CO₃. Use methyl orange indicator. Calculate exact molarity from titration results. Discuss primary standard requirements.

Anhydrous Na₂CO₃, Approximately 0.1M HCl, Methyl orange, Volumetric flasks, Analytical balance

KLB Secondary Chemistry Form 3, Pages 65-67

THE MOLE

Back Titration Method
Redox Titrations - Principles
Redox Titrations - KMnO₄ Standardization

By the end of the lesson, the learner should be able to:
Understand principle of back titration
Apply back titration to determine composition
Calculate concentrations using back titration data
Determine atomic masses from back titration
Standardize KMnO₄ solution using iron(II) salt
Calculate molarity from redox titration data
Apply 1:5 mole ratio in calculations
Prepare solutions for redox titrations

Experiment: Determine atomic mass of divalent metal in MCO₃. Add excess HCl to carbonate, then titrate excess with NaOH. Calculate moles of acid that reacted with carbonate. Determine metal's atomic mass.
Experiment: Standardize KMnO₄ using FeSO₄(NH₄)₂SO₄·6H₂O. Dissolve iron salt in boiled, cooled water. Titrate with KMnO₄ until persistent pink color. Calculate molarity using 5:1 mole ratio.

Metal carbonate sample, 0.5M HCl, 0M NaOH, Phenolphthalein, Conical flasks
Potassium manganate(VII), Potassium dichromate(VI), Iron(II) solutions, Color change charts
Iron(II) ammonium sulfate, KMnO₄ solution, Dilute H₂SO₄, Pipettes, Burettes

KLB Secondary Chemistry Form 3, Pages 67-70
KLB Secondary Chemistry Form 3, Pages 70-72

THE MOLE

Water of Crystallization Determination

By the end of the lesson, the learner should be able to:
Determine water of crystallization in hydrated salts
Use redox titration to find formula of hydrated salt
Calculate value of 'n' in crystallization formulas
Apply analytical data to determine complete formulas

Experiment: Determine 'n' in FeSO₄(NH₄)₂SO₄·nH₂O. Dissolve known mass in acid, titrate with standardized KMnO₄. Calculate moles of iron(II), hence complete formula. Compare theoretical and experimental values.

Hydrated iron(II) salt, Standardized KMnO₄, Dilute H₂SO₄, Analytical balance

KLB Secondary Chemistry Form 3, Pages 72-73

THE MOLE

Atomicity and Molar Gas Volume

By the end of the lesson, the learner should be able to:
Define atomicity of gaseous elements
Classify gases as monoatomic, diatomic, or triatomic
Determine molar gas volume experimentally
Calculate gas densities and molar masses

Experiment: Measure volumes and masses of different gases (O₂, CO₂, Cl₂). Calculate densities and molar masses. Determine volume occupied by one mole. Compare values at different conditions.

Gas syringes (50cm³), Various gases, Analytical balance, Gas supply apparatus

KLB Secondary Chemistry Form 3, Pages 73-75

THE MOLE

Combining Volumes of Gases - Experimental Investigation
Gas Laws and Chemical Equations

By the end of the lesson, the learner should be able to:
Investigate Gay-Lussac's law experimentally
Measure combining volumes of reacting gases
Determine simple whole number ratios
Write equations from volume relationships

Experiment: React NH₃ and HCl gases in measured volumes. Observe formation of NH₄Cl solid. Measure residual gas volumes. Determine combining ratios. Apply to other gas reactions.

Gas syringes, Dry NH₃ generator, Dry HCl generator, Glass connecting tubes, Clips
Scientific calculators, Gas law charts, Volume ratio examples

KLB Secondary Chemistry Form 3, Pages 75-77

ORGANIC CHEMISTRY I

Introduction to Organic Chemistry and Hydrocarbons
Sources of Alkanes - Natural Gas, Biogas, and Crude Oil

By the end of the lesson, the learner should be able to:
Define organic chemistry and hydrocarbons
Explain why carbon forms many compounds
Classify hydrocarbons into alkanes, alkenes, and alkynes
Identify the bonding in carbon compounds
Identify natural sources of alkanes
Describe composition of natural gas and biogas
Explain crude oil as major source of alkanes
Describe biogas digester and its operation

Teacher exposition: Definition of organic chemistry. Discussion: Unique properties of carbon - tetravalency, catenation, multiple bonding. Q/A: Examples of hydrocarbons in daily life. Introduction to three main groups of hydrocarbons.
Discussion: Natural gas composition (80% methane). Explanation: Biogas formation from organic waste decomposition. Teacher demonstration: Biogas digester model/diagram. Q/A: Environmental benefits of biogas production.

Carbon models, Hydrocarbon structure charts, Molecular model kits
Biogas digester model/diagram, Natural gas composition charts, Organic waste samples

KLB Secondary Chemistry Form 3, Pages 86-87

ORGANIC CHEMISTRY I

Fractional Distillation of Crude Oil

By the end of the lesson, the learner should be able to:
Explain fractional distillation process
Perform fractional distillation of crude oil
Identify different fractions and their uses
Relate boiling points to molecular size

Experiment: Fractional distillation of crude oil using improvised column. Collect fractions at different temperatures (120°C intervals up to 350°C). Test fractions for appearance, flammability, and viscosity. Record observations and relate to molecular size.

Crude oil sample, Boiling tubes, High-temperature thermometer, Sand/porcelain chips, Bunsen burner, Test tubes

KLB Secondary Chemistry Form 3, Pages 87-89

ORGANIC CHEMISTRY I

Cracking of Alkanes - Thermal and Catalytic Methods
Alkane Series and Homologous Series Concept

By the end of the lesson, the learner should be able to:
Define cracking of alkanes
Distinguish between thermal and catalytic cracking
Write equations for cracking reactions
Explain industrial importance of cracking

Teacher exposition: Definition and purpose of cracking. Discussion: Thermal vs catalytic cracking conditions. Worked examples: Cracking equations producing smaller alkanes, alkenes, and hydrogen. Q/A: Industrial applications and hydrogen production.

Cracking process diagrams, Chemical equation charts, Catalyst samples for demonstration
Alkane series chart, Molecular formula worksheets, Periodic table

KLB Secondary Chemistry Form 3, Pages 89-90

ORGANIC CHEMISTRY I

Nomenclature of Alkanes - Straight Chain and Branched

By the end of the lesson, the learner should be able to:
Name straight-chain alkanes using IUPAC rules
Identify parent chains in branched alkanes
Name branched alkanes with substituent groups
Apply systematic naming rules correctly

Teacher demonstration: Step-by-step naming of branched alkanes. Rules application: Longest chain identification, numbering from nearest branch, substituent naming. Practice exercises: Various branched alkane structures. Group work: Name complex branched alkanes.

Structural formula charts, IUPAC naming rules poster, Molecular model kits

KLB Secondary Chemistry Form 3, Pages 90-92

ORGANIC CHEMISTRY I

Isomerism in Alkanes - Structural Isomers
Laboratory Preparation of Methane

By the end of the lesson, the learner should be able to:
Define isomerism in alkanes
Draw structural isomers of butane and pentane
Distinguish between chain and positional isomerism
Predict number of isomers for given alkanes
Describe laboratory preparation of methane
Perform methane preparation experiment safely
Test physical and chemical properties of methane
Write equation for methane preparation

Teacher exposition: Isomerism definition and types. Practical exercise: Draw all isomers of butane and pentane. Discussion: Physical property differences between isomers. Model building: Use molecular models to show isomeric structures.
Experiment: Heat mixture of sodium ethanoate and soda lime. Collect methane gas over water. Tests: Color, smell, combustion, reaction with bromine in dark. Record observations in table format. Safety precautions during gas collection.

Molecular model kits, Isomerism charts, Structural formula worksheets
Sodium ethanoate, Soda lime, Round-bottomed flask, Gas collection apparatus, Bromine water, Wooden splints

KLB Secondary Chemistry Form 3, Pages 92-94
KLB Secondary Chemistry Form 3, Pages 94-96

ORGANIC CHEMISTRY I

Laboratory Preparation of Ethane
Physical Properties of Alkanes

By the end of the lesson, the learner should be able to:
Prepare ethane using sodium propanoate and soda lime
Compare preparation methods of methane and ethane
Test properties of ethane gas
Write general equation for alkane preparation

Experiment: Prepare ethane from sodium propanoate and soda lime. Compare with methane preparation method. Carry out similar tests as for methane. Discussion: General pattern for alkane preparation from sodium alkanoates.

Sodium propanoate, Soda lime, Gas collection apparatus, Testing materials
Physical properties data tables, Graph paper, Calculators, Solubility demonstration materials

KLB Secondary Chemistry Form 3, Pages 94-96

ORGANIC CHEMISTRY I

Chemical Properties of Alkanes - Combustion and Substitution

By the end of the lesson, the learner should be able to:
Write equations for complete and incomplete combustion
Explain substitution reactions with halogens
Describe conditions for halogenation reactions
Name halogenated alkane products

Worked examples: Combustion equations for various alkanes. Teacher demonstration: Methane + bromine in sunlight (or simulation). Discussion: Free radical mechanism in substitution. Practice: Write equations for chlorination of methane.

Molecular models, Halogenation reaction charts, Chemical equation worksheets

KLB Secondary Chemistry Form 3, Pages 97-98

ORGANIC CHEMISTRY I

Uses of Alkanes in Industry and Daily Life

By the end of the lesson, the learner should be able to:
List major uses of different alkanes
Explain industrial applications of alkanes
Describe environmental considerations
Evaluate economic importance of alkanes

Discussion: Uses of gaseous alkanes as fuels. Teacher exposition: Industrial applications - carbon black, methanol production, hydrogen source. Q/A: Environmental impact and cleaner fuel initiatives. Assignment: Research local uses of alkane products.

Industrial application charts, Product samples, Environmental impact materials

KLB Secondary Chemistry Form 3, Pages 98-100