GAS LAWS

Boyle's Law - Introduction and Experimental Investigation

By the end of the lesson, the learner should be able to:
State Boyle's law
Explain Boyle's law using kinetic theory of matter
Investigate the relationship between pressure and volume of a fixed mass of gas
Plot graphs to illustrate Boyle's law

Teacher demonstration: Use bicycle pump to show volume-pressure relationship. Students observe force needed to compress gas. Q/A: Review kinetic theory. Class experiment: Investigate pressure-volume relationship using syringes. Record observations in table format. Discuss observations using kinetic theory.

Bicycle pump, Syringes, Gas jars, Chart showing volume-pressure relationship

KLB Secondary Chemistry Form 3, Pages 1-3

GAS LAWS

Boyle's Law - Mathematical Expression and Graphical Representation
Boyle's Law - Numerical Problems and Applications
Charles's Law - Introduction and Temperature Scales
Charles's Law - Experimental Investigation and Mathematical Expression

By the end of the lesson, the learner should be able to:
Express Boyle's law mathematically
Apply the equation PV = constant
Plot and interpret pressure vs volume graphs
Plot pressure vs 1/volume graphs
State Charles's law
Convert temperatures between Celsius and Kelvin scales
Define absolute zero temperature
Explain the concept of absolute temperature

Q/A: Recall previous lesson observations. Teacher exposition: Derive P₁V₁ = P₂V₂ equation from experimental data. Students plot graphs of pressure vs volume and pressure vs 1/volume. Analyze graph shapes and interpret mathematical relationship.
Teacher demonstration: Flask with colored water column experiment. Q/A: Observe volume changes with temperature. Exposition: Introduce Kelvin scale and absolute zero concept. Practice: Temperature conversions between °C and K. Discuss absolute zero and ideal gas concept.

Graph papers, Scientific calculators, Chart showing mathematical expressions
Scientific calculators, Worked example charts, Unit conversion tables
Round-bottomed flask, Narrow glass tube, Colored water, Rubber bung, Hot and cold water baths
Glass apparatus, Thermometers, Graph papers, Water baths at different temperatures

KLB Secondary Chemistry Form 3, Pages 3-4
KLB Secondary Chemistry Form 3, Pages 6-8

GAS LAWS

Charles's Law - Numerical Problems and Applications
Combined Gas Law and Standard Conditions

By the end of the lesson, the learner should be able to:
Solve numerical problems using Charles's law
Apply V₁/T₁ = V₂/T₂ in calculations
Predict gas behavior with temperature changes
Relate Charles's law to everyday phenomena

Worked examples: Step-by-step problem solving with temperature conversions. Supervised practice: Calculate volumes at different temperatures. Discuss applications: hot air balloons, tire pressure changes, weather balloons. Assignment: Practice problems with real-life contexts.

Scientific calculators, Temperature conversion charts, Application examples
Scientific calculators, Combined law derivation charts, Standard conditions reference table

KLB Secondary Chemistry Form 3, Pages 10-12

GAS LAWS

Introduction to Diffusion - Experimental Investigation
Rates of Diffusion - Comparative Study

By the end of the lesson, the learner should be able to:
Define diffusion process
Investigate diffusion in liquids and gases
Compare rates of diffusion in different media
Explain diffusion using kinetic theory

Class experiments: (a) KMnO₄ crystal in water - observe spreading over time. (b) Bromine vapor in gas jars - observe color distribution. (c) Ammonia gas in combustion tube with litmus paper. Record observations over time. Discuss particle movement and kinetic energy.

KMnO₄ crystals, Bromine liquid, Gas jars, Combustion tube, Litmus papers, Stopwatch
Glass tube (25cm), Cotton wool, Concentrated NH₃ and HCl, Stopwatch, Ruler, Safety equipment

KLB Secondary Chemistry Form 3, Pages 14-16

GAS LAWS

Graham's Law of Diffusion - Theory and Mathematical Expression

By the end of the lesson, the learner should be able to:
State Graham's law of diffusion
Express Graham's law mathematically
Relate diffusion rate to molecular mass and density
Explain the inverse relationship between rate and √molecular mass

Teacher exposition: Graham's law statement and mathematical derivation. Discussion: Rate ∝ 1/√density and Rate ∝ 1/√molecular mass. Derive comparative expressions for two gases. Explain relationship between density and molecular mass. Practice: Identify faster diffusing gas from molecular masses.

Graham's law charts, Molecular mass tables, Mathematical derivation displays

KLB Secondary Chemistry Form 3, Pages 18-20

GAS LAWS
THE MOLE

Graham's Law - Numerical Applications and Problem Solving
Relative Mass - Introduction and Experimental Investigation

By the end of the lesson, the learner should be able to:
Solve numerical problems using Graham's law
Calculate relative rates of diffusion
Determine molecular masses from diffusion data
Compare diffusion times for equal volumes of gases
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

Worked examples: Calculate relative diffusion rates using √(M₂/M₁). Problems involving time comparisons for equal volumes. Calculate unknown molecular masses from rate data. Supervised practice: Various Graham's law calculations. Real-life applications: gas separation, gas masks.
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.

Scientific calculators, Worked example charts, Molecular mass reference tables
Different sized nails ( 5-15cm), Beam balance, Fruits of different masses, Reference charts

KLB Secondary Chemistry Form 3, Pages 20-22
KLB Secondary Chemistry Form 3, Pages 25-27

THE MOLE

Avogadro's Constant and the Mole Concept

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

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.

Beam balance, Various sized nails, Scientific calculators, Avogadro's constant charts

KLB Secondary Chemistry Form 3, Pages 27-30

THE MOLE

Interconversion of Mass and Moles for Elements

By the end of the lesson, the learner should be able to:
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

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.

Scientific calculators, Periodic table, Worked example charts, Formula triangles

KLB Secondary Chemistry Form 3, Pages 30-32

THE MOLE

Molecules and Moles - Diatomic Elements

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

KLB Secondary Chemistry Form 3, Pages 29-30

THE MOLE

Empirical Formula - Experimental Determination
Empirical Formula - Reduction Method
Empirical Formula - Percentage Composition Method

By the end of the lesson, the learner should be able to:
Define empirical formula
Determine empirical formula from experimental data
Calculate mole ratios from mass data
Express results as simplest whole number ratios
Determine empirical formula using reduction reactions
Calculate empirical formula from reduction data
Apply reduction method to copper oxides
Analyze experimental errors and sources

Experiment: Burning magnesium in air to form magnesium oxide. Measure masses before and after reaction. Calculate moles of Mg and O from mass data. Determine mole ratio and empirical formula. Safety precautions during heating.
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.

Crucible and lid, Magnesium ribbon, Bunsen burner, Beam balance, Tongs, Safety equipment
Combustion tube, Porcelain boat, Copper(II) oxide, Laboratory gas, Beam balance, Bunsen burner
Scientific calculators, Percentage composition charts, Worked example displays

KLB Secondary Chemistry Form 3, Pages 32-35
KLB Secondary Chemistry Form 3, Pages 35-37

THE MOLE

Molecular Formula - Determination from Empirical Formula

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

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.

Scientific calculators, Molecular mass charts, Worked example displays

KLB Secondary Chemistry Form 3, Pages 38-40

THE MOLE

Molecular Formula - Combustion Analysis

By the end of the lesson, the learner should be able to:
Determine molecular formula from combustion data
Calculate moles of products in combustion
Relate product moles to reactant composition
Apply combustion analysis to hydrocarbons

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, Combustion analysis charts, Molecular models of hydrocarbons

KLB Secondary Chemistry Form 3, Pages 40-41

THE MOLE

Concentration and Molarity of Solutions

By the end of the lesson, the learner should be able to:
Define concentration and molarity of solutions
Calculate molarity from mass and volume data
Convert between different concentration units
Apply molarity calculations to various solutions

Teacher exposition: Definition of molarity (moles/dm³). Worked examples: Calculate molarity from mass of solute and volume. Convert between g/dm³ and mol/dm³. Practice problems: Various salt solutions and their molarities.

Scientific calculators, Molarity charts, Various salt samples for demonstration

KLB Secondary Chemistry Form 3, Pages 41-43

THE MOLE

Preparation of Molar Solutions
Dilution of 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
Define dilution process
Apply dilution formula M₁V₁ = M₂V₂
Calculate concentrations after dilution
Prepare dilute solutions from concentrated ones

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.
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 (250, 500, 1000cm³), Sodium hydroxide pellets, Beam balance, Wash bottles, Beakers
Volumetric flasks, Hydrochloric acid (2M), Measuring cylinders, Pipettes, Safety equipment

KLB Secondary Chemistry Form 3, Pages 43-46
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

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

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.

Test tubes, Lead(II) nitrate solution, Potassium iodide solution, Burettes, Ethanol, Rulers

KLB Secondary Chemistry Form 3, Pages 53-56

THE MOLE

Stoichiometry - Gas Evolution Reactions

By the end of the lesson, the learner should be able to:
Determine stoichiometry of gas-producing reactions
Collect and measure gas volumes
Calculate mole ratios involving gases
Write equations for acid-carbonate reactions

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.

Conical flask, Thistle funnel, Plastic bags, Rubber bands, Sodium carbonate, HCl solution

KLB Secondary Chemistry Form 3, Pages 56-58

THE MOLE

Volumetric Analysis - Introduction and Apparatus
Titration - Acid-Base Neutralization

By the end of the lesson, the learner should be able to:
Define volumetric analysis and titration
Identify and use titration apparatus correctly
Explain functions of pipettes and burettes
Demonstrate proper reading techniques
Perform acid-base titrations accurately
Use indicators to determine end points
Record titration data properly
Calculate average titres from multiple readings

Practical session: Familiarization with pipettes and burettes. Practice filling and reading burettes accurately. Learn proper meniscus reading. Use pipette fillers safely. Rinse apparatus with appropriate solutions.
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.

Pipettes (10, 20, 25cm³), Burettes (50cm³), Pipette fillers, Conical flasks, Various solutions
Burettes, Pipettes, 0.1M NaOH, 0.1M HCl, Phenolphthalein indicator, Conical flasks

KLB Secondary Chemistry Form 3, Pages 58-59
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
Back Titration Method

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
Metal carbonate sample, 0.5M HCl, 0M NaOH, Phenolphthalein, Conical flasks

KLB Secondary Chemistry Form 3, Pages 65-67

THE MOLE

Redox Titrations - Principles

By the end of the lesson, the learner should be able to:
Explain principles of redox titrations
Identify color changes in redox reactions
Understand self-indicating nature of some redox reactions
Write ionic equations for redox processes

Teacher exposition: Redox titration principles. Demonstrate color changes: MnO₄⁻ (purple) → Mn²⁺ (colorless), Cr₂O₇²⁻ (orange) → Cr³⁺ (green). Discussion: Self-indicating reactions. Write half-equations and overall ionic equations.

Potassium manganate(VII), Potassium dichromate(VI), Iron(II) solutions, Color change charts

KLB Secondary Chemistry Form 3, Pages 68-70

THE MOLE

Redox Titrations - KMnO₄ Standardization
Water of Crystallization Determination

By the end of the lesson, the learner should be able to:
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
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: 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.
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.

Iron(II) ammonium sulfate, KMnO₄ solution, Dilute H₂SO₄, Pipettes, Burettes
Hydrated iron(II) salt, Standardized KMnO₄, Dilute H₂SO₄, Analytical balance

KLB Secondary Chemistry Form 3, Pages 70-72
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