SALTS

Types of salts.

By the end of the lesson, the learner should be able to:
Define a salt.
Describe various types of salts and give several examples in each case.

Descriptive approach. Teacher exposes new concepts.

text book

K.L.B. BOOK II P. 91

SALTS

Solubility of salts in water.
Solubility of bases in water.
Methods of preparing various salts.

By the end of the lesson, the learner should be able to:
To test solubility of various salts in cold water/warm water.
To describe various methods of preparing some salts.

Class experiments- Dissolve salts in 5 cc of water.
Record the solubility in a table,
Analyse the results.
Experimental and descriptive treatments of preparation of salts e.g. ZnSO4, CuSO4, NaCl and Pb(NO3)2.

Sulphates, chlorides, nitrates, carbonates of various metals.
Oxides, hydroxides, of various metals, litmus papers.
CuO, H2SO4, HCl, NaOH, PbCO3, dil HNO3.

K.L.B. BOOK II PP. 92-93
K.L.B. BOOK II pp96

SALTS

Direct synthesis of a salts.
Ionic equations.

By the end of the lesson, the learner should be able to:
To describe direct synthesis of a salt.
To write balanced equations for the reactions.

Group experiments- preparation of iron (II) sulphide by direct synthesis.
Give other examples of salts prepared by direct synthesis.
Students write down corresponding balanced equations.

Iron,
Sulphur
PbNO3, MgSO4 solutions.

K.L.B. BOOK II P. 104

SALTS

Effects of heat on carbonates.
Effects of heat on nitrates.
Effects of heat on sulphates.

By the end of the lesson, the learner should be able to:
To state effects of heat on carbonates.
To predict products resulting from heating metal carbonates.

Group experiments- To investigate effects of heat on Na2CO3, K2CO3, CaCO3, ZnCO3, PbCO3, e.t.c.
Observe various colour changes before, during and after heating.
Write equations for the reactions.

Various carbonates.
Common metal nitrates.
Common sulphates.

K.L.B. BOOK II PP. 108-109

SALTS

Hygroscopy, Deliquescence and Efflorescence.
Uses of salts.

By the end of the lesson, the learner should be able to:
To define hygroscopic deliquescent and efflorescent salts.
To give examples of hygroscopic deliquescent and efflorescent salts.

Prepare a sample of various salts.
Expose them to the atmosphere overnight.
Students classify the salts as hygroscopic, deliquescent and / or efflorescent.

K.L.B. BOOK II P. 114

EFFECTS OF AN ELECTRIC CURRENT ON SUBSTANCES.

Electrical conductivity.
Molten electrolytes.
Electrolysis.
Aqueous electrolytes. Electrodes.
Reaction on electrodes.

By the end of the lesson, the learner should be able to:
To test for electrical conductivities of substances.
To define an electrolyte
To test for electrical conductivities of electrodes.

Group experiments- to identify conductors and non-conductors.
Explain the difference in (non) conductivities.
To investigate chemical effect of an electric current.
Classify the solutions as electrolyte or non -electrolytes.
Discuss the electrical properties of the solutions.

Various solids, bulb, battery, & wires.
Molten candle wax
Sugar
Sulphur
Lead oxide.
Graphite electrodes
Battery
Various aqueous solutions switch bulb.
Various aqueous solutions switch.

K.L.B. BOOK II PP. 118-119
K.L.B. BOOK II PP.122-123

EFFECTS OF AN ELECTRIC CURRENT ON SUBSTANCES.

Binary electrolyte.
Application of electrolysis.
Electroplating.

By the end of the lesson, the learner should be able to:
To define a binary electrolyte.
To state the products of a binary electrolyte.

Completing a table of electrolysis of binary electrolytes.

text book
Silver nitrate
Iron nail
Complete circuit battery.

K.L.B. BOOK II P.127

CARBON AND SOME OF ITS COMPOUNDS.

Allotropy.
Physical and chemical properties of diamond, graphite and amorphous carbon

By the end of the lesson, the learner should be able to:
Define allotropes and allotropy.
Identify allotropes of carbon.
Represent diamond and graphite diagrammatically.

Teacher exposes new terms.
Review covalent bond.
Discuss boding in diamond and graphite.

text book
Charcoal, graphite.

K.L.B. BOOK II PP. 131-133

CARBON AND SOME OF ITS COMPOUNDS.

Burning carbon and oxygen.
Reduction properties of carbon.
Reaction of carbon with acids. Preparation of CO2.

By the end of the lesson, the learner should be able to:
Describe reaction of carbon with oxygen.

Teacher demonstration- Prepare oxygen and pass dry oxygen into a tube containing carbon. Heat the carbon. Observe effects on limewater.

Carbon, limewater, tube, limewater stand& Bunsen burner.
CuO, pounded charcoal, Bunsen burner& bottle top
Conc. HNO3, limewater.

K.L.B. BOOK II PP. 134-135

CARBON AND SOME OF ITS COMPOUNDS.

Properties of CO2.
Chemical equations for reactions involving CO2.
Uses of CO2.
Carbon monoxide lab preparation.
Chemical properties of carbon monoxide.

By the end of the lesson, the learner should be able to:
Describe properties of CO2
State uses of CO2

Simple experiments to determine properties of CO2.

Discuss the observations.

Discuss briefly the uses of CO2.

Lime water,
Magnesium ribbon,
Universal indicator,
lit candle.
text book

K.L.B. BOOK II PP.138-139
K.L.B. BOOK II PP.140-1

CARBON AND SOME OF ITS COMPOUNDS.

Carbonates and hydrogen carbonates.
Heating carbonates and hydrogen carbonates.
Extraction of sodium carbonate from trona.

By the end of the lesson, the learner should be able to:
To write chemical equations for reactions of carbonates and hydrogen carbonates with acids.

Discuss the observations above.
Write chemical equations for the reactions.

text book

K.L.B. BOOK II

CARBON AND SOME OF ITS COMPOUNDS.

Solvay process of preparing sodium carbonate.
Importance of carbon in nature. & its effects on the environment.

By the end of the lesson, the learner should be able to:
To draw schematic diagram for extraction of sodium carbonates.

Discuss each step of the process.

Write relevant equations.

text book, chart
text book

K.L.B. BOOK II

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

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
Solve numerical problems involving Boyle's law
Convert between different pressure units
Apply Boyle's law to real-life situations
Calculate volumes and pressures using P₁V₁ = P₂V₂

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.
Worked examples: Demonstrate step-by-step problem solving. Supervised practice: Students solve problems involving pressure and volume calculations. Convert units (mmHg, atm, Pa). Discuss applications in tire inflation, aerosol cans. Assignment: Additional practice problems.

Graph papers, Scientific calculators, Chart showing mathematical expressions
Scientific calculators, Worked example charts, Unit conversion tables

KLB Secondary Chemistry Form 3, Pages 3-4
KLB Secondary Chemistry Form 3, Pages 4-5

GAS LAWS

Charles's Law - Introduction and Temperature Scales

By the end of the lesson, the learner should be able to:
State Charles's law
Convert temperatures between Celsius and Kelvin scales
Define absolute zero temperature
Explain the concept of absolute temperature

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.

Round-bottomed flask, Narrow glass tube, Colored water, Rubber bung, Hot and cold water baths

KLB Secondary Chemistry Form 3, Pages 6-8

GAS LAWS

Charles's Law - Experimental Investigation and Mathematical Expression
Charles's Law - Numerical Problems and Applications

By the end of the lesson, the learner should be able to:
Investigate relationship between volume and temperature
Express Charles's law mathematically
Plot volume vs temperature graphs
Extrapolate graphs to find absolute zero

Class experiment: Volume-temperature relationship using flask and capillary tube. Record data at different temperatures. Plot graphs: volume vs temperature (°C) and volume vs absolute temperature (K). Extrapolate graph to find absolute zero. Derive V₁/T₁ = V₂/T₂ equation.

Glass apparatus, Thermometers, Graph papers, Water baths at different temperatures
Scientific calculators, Temperature conversion charts, Application examples

KLB Secondary Chemistry Form 3, Pages 8-10

GAS LAWS

Combined Gas Law and Standard Conditions

By the end of the lesson, the learner should be able to:
Derive the combined gas law equation
Apply PV/T = constant in problem solving
Define standard temperature and pressure (s.t.p)
Define room temperature and pressure (r.t.p)

Q/A: Combine Boyle's and Charles's laws. Teacher exposition: Derive P₁V₁/T₁ = P₂V₂/T₂. Define s.t.p (273K, 760mmHg) and r.t.p (298K, 760mmHg). Worked examples: Problems involving changes in all three variables. Supervised practice: Complex gas law calculations.

Scientific calculators, Combined law derivation charts, Standard conditions reference table

KLB Secondary Chemistry Form 3, Pages 12-14

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
Compare diffusion rates of different gases
Investigate factors affecting diffusion rates
Measure relative distances covered by diffusing gases
Calculate rates of diffusion using distance and time data

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.
Class experiment: Ammonia and HCl diffusion in glass tube. Insert cotton wool soaked in concentrated NH₃ and HCl at opposite ends. Time the formation of white NH₄Cl ring. Measure distances covered by each gas. Calculate rates: distance/time. Compare molecular masses of NH₃ and HCl.

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
KLB Secondary Chemistry Form 3, Pages 16-18

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

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.

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

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
Molecules and Moles - Diatomic 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
Distinguish between atoms and molecules
Define relative molecular mass
Calculate moles of molecules from given mass
Determine number of atoms in molecular compounds

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.
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.

Scientific calculators, Periodic table, Worked example charts, Formula triangles
Molecular models, Charts showing diatomic elements, Scientific calculators

KLB Secondary Chemistry Form 3, Pages 30-32
KLB Secondary Chemistry Form 3, Pages 29-30

THE MOLE

Empirical Formula - Experimental Determination
Empirical Formula - Reduction 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

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.

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

KLB Secondary Chemistry Form 3, Pages 32-35

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

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
Concentration and Molarity of Solutions

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
Define concentration and molarity of solutions
Calculate molarity from mass and volume data
Convert between different concentration units
Apply molarity calculations to various solutions

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.
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, Combustion analysis charts, Molecular models of hydrocarbons
Scientific calculators, Molarity charts, Various salt samples for demonstration

KLB Secondary Chemistry Form 3, Pages 40-41
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

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

KLB Secondary Chemistry Form 3, Pages 43-46

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
Volumetric Analysis - Introduction and Apparatus

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
Define volumetric analysis and titration
Identify and use titration apparatus correctly
Explain functions of pipettes and burettes
Demonstrate proper reading techniques

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.
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.

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 56-58
KLB Secondary Chemistry Form 3, Pages 58-59

THE MOLE

Titration - Acid-Base Neutralization
Titration - Diprotic Acids

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

KLB Secondary Chemistry Form 3, Pages 59-62

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

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

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.

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

KLB Secondary Chemistry Form 3, Pages 67-70

THE MOLE

Redox Titrations - Principles
Redox Titrations - KMnO₄ Standardization

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
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

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.
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.

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 68-70
KLB Secondary Chemistry Form 3, Pages 70-72

THE MOLE

Water of Crystallization Determination
Atomicity and Molar Gas Volume

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

KLB Secondary Chemistry Form 3, Pages 72-73

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