ACIDS, BASES AND SALTS

Definition of Acids

By the end of the lesson, the learner should be able to:
- Define an acid in terms of hydrogen ions
-Investigate reactions of magnesium and zinc carbonate with different acids
-Write equations for reactions taking place
-Explain why magnesium strip should be cleaned

Class experiment: React cleaned magnesium strips with 2M HCl, 2M ethanoic acid, 2M H₂SO₄, 2M ethanedioic acid. Record observations in table. Repeat using zinc carbonate. Write chemical equations. Discuss hydrogen ion displacement and gas evolution.

Magnesium strips, zinc carbonate, 2M HCl, 2M ethanoic acid, 2M H₂SO₄, 2M ethanedioic acid, test tubes, test tube rack

KLB Secondary Chemistry Form 4, Pages 1-3

ACIDS, BASES AND SALTS

Strength of Acids
Definition of Bases
Strength of Bases
Acid-Base Reactions
Effect of Solvent on Acids

By the end of the lesson, the learner should be able to:
- Compare strengths of acids using pH values
-Determine strengths of acids by comparing their electrical conductivity
-Classify acids as either strong or weak
-Explain complete and partial dissociation of acids
- Compare strengths of bases using pH values
-Determine strengths of bases by comparing their electrical conductivity
-Classify bases as either strong or weak
-Explain complete and partial ionization of bases

Class experiment: Test pH of 2M HCl and 2M ethanoic acid using universal indicator. Set up electrical conductivity apparatus with both acids. Record milliammeter readings. Compare results and explain in terms of hydrogen ion concentration. Discuss strong vs weak acid definitions.
Class experiment: Test pH of 2M NaOH and 2M ammonia solution using universal indicator. Test electrical conductivity of both solutions using same apparatus as acids. Compare deflections and pH values. Explain in terms of OH⁻ ion concentration and complete vs partial ionization.

2M HCl, 2M ethanoic acid, universal indicator, pH chart, electrical conductivity apparatus, milliammeter, carbon electrodes, beakers, wires
Calcium hydroxide, red litmus paper, phenolphthalein indicator, distilled water, test tubes, spatula, evaporating dish
2M NaOH, 2M ammonia solution, universal indicator, pH chart, electrical conductivity apparatus, milliammeter, carbon electrodes
Various acids and bases from previous lessons, indicators, beakers, measuring cylinders, stirring rods
HCl gas, distilled water, methylbenzene, magnesium ribbon, calcium carbonate, litmus paper, test tubes, gas absorption apparatus

KLB Secondary Chemistry Form 4, Pages 3-5
KLB Secondary Chemistry Form 4, Pages 5-7

ACIDS, BASES AND SALTS

Effect of Solvent on Bases
Amphoteric Oxides and Hydroxides
Definition of Salts and Precipitation

By the end of the lesson, the learner should be able to:
- Investigate effect of polar and non-polar solvents on ammonia gas
-Compare ammonia behavior in water vs methylbenzene
-Explain formation of ammonium hydroxide
-Write equations for ammonia dissolution in water

Class experiment: Test dry ammonia with dry litmus. Dissolve ammonia in water and test with litmus. Dissolve ammonia in methylbenzene and test with litmus. Record observations in table. Write equation for NH₃ + H₂O reaction. Explain why only aqueous ammonia shows basic properties.

Dry ammonia gas, distilled water, methylbenzene, red litmus paper, test tubes, gas collection apparatus
Al₂O₃, ZnO, PbO, Zn(OH)₂, Al(OH)₃, Pb(OH)₂, 2M HNO₃, 2M NaOH, boiling tubes, heating source
Na₂CO₃ solution, salt solutions containing various metal ions, test tubes, droppers

KLB Secondary Chemistry Form 4, Pages 9-10

ACIDS, BASES AND SALTS

Solubility of Chlorides, Sulphates and Sulphites

By the end of the lesson, the learner should be able to:
- Find out cations that form insoluble chlorides, sulphates and sulphites
-Write ionic equations for formation of insoluble salts
-Distinguish between sulphate and sulphite precipitates
-Investigate effect of warming on precipitates

Class experiment: Add NaCl, Na₂SO₄, Na₂SO₃ to solutions of Pb²⁺, Ba²⁺, Mg²⁺, Ca²⁺, Zn²⁺, Cu²⁺, Fe²⁺, Fe³⁺, Al³⁺. Warm mixtures. Record observations in table. Test sulphite precipitates with dilute HCl. List soluble and insoluble salts.

2M NaCl, 2M Na₂SO₄, 2M Na₂SO₃, 0.1M salt solutions, dilute HCl, test tubes, heating source

KLB Secondary Chemistry Form 4, Pages 14-16

ACIDS, BASES AND SALTS

Complex Ions Formation

By the end of the lesson, the learner should be able to:
- Explain formation of complex ions
-Investigate reactions with excess sodium hydroxide and ammonia
-Identify metal ions that form complex ions
-Write equations for complex ion formation

Class experiment: Add NaOH dropwise then in excess to Mg²⁺, Ca²⁺, Zn²⁺, Al³⁺, Cu²⁺, Fe²⁺, Fe³⁺, Pb²⁺ solutions. Repeat with NH₃ solution. Record observations showing precipitate formation and dissolution. Write equations for complex ion formation: [Zn(OH)₄]²⁻, [Al(OH)₄]⁻, [Pb(OH)₄]²⁻, [Zn(NH₃)₄]²⁺, [Cu(NH₃)₄]²⁺.

2M NaOH, 2M NH₃ solution, 0.5M salt solutions, test tubes, droppers

KLB Secondary Chemistry Form 4, Pages 15-16

ACIDS, BASES AND SALTS

Solubility and Saturated Solutions
Effect of Temperature on Solubility
Solubility Curves and Applications

By the end of the lesson, the learner should be able to:
- Define the term solubility
-Determine solubility of a given salt at room temperature
-Calculate mass of solute and solvent
-Express solubility in different units
- Plot solubility curves for various salts
-Use solubility curves to determine mass of crystals formed
-Apply solubility curves to practical problems
-Compare solubility patterns of different salts

Class experiment: Weigh evaporating dish and watch glass. Measure 20cm³ saturated KNO₃ solution. Record temperature. Evaporate to dryness carefully. Calculate masses of solute, solvent, and solution. Determine solubility per 100g water and in moles per litre. Discuss definition and significance.
Using data from textbook, plot solubility curves for KNO₃, KClO₃, NaCl, CaSO₄. Calculate mass of crystals deposited when saturated solutions are cooled. Work through examples: KClO₃ cooled from 70°C to 30°C. Discuss applications in salt extraction and purification.

Saturated KNO₃ solution, evaporating dish, watch glass, measuring cylinder, thermometer, balance, heating source
KClO₃, measuring cylinders, thermometer, burette, boiling tubes, heating source, graph paper
Graph paper, ruler, pencil, calculator, data tables from textbook

KLB Secondary Chemistry Form 4, Pages 16-18
KLB Secondary Chemistry Form 4, Pages 20-21

ACIDS, BASES AND SALTS

Fractional Crystallization

By the end of the lesson, the learner should be able to:
- Define fractional crystallization
-Apply knowledge of solubility curves in separation of salts
-Calculate masses of salts that crystallize
-Explain separation of salt mixtures

Work through separation problems using solubility data for KNO₃ and KClO₃ mixtures. Calculate which salt crystallizes first when cooled from 50°C to 20°C. Plot combined solubility curves. Discuss applications in Lake Magadi and Ngomeni salt works. Solve practice problems.

Calculator, graph paper, data tables, worked examples from textbook

KLB Secondary Chemistry Form 4, Pages 21-22

ACIDS, BASES AND SALTS

Hardness of Water - Investigation
Types and Causes of Water Hardness

By the end of the lesson, the learner should be able to:
- Determine the effects of various salt solutions on soap
-Identify cations that cause hardness
-Distinguish between hard and soft water
-Investigate effect of boiling on water hardness

Class experiment: Test soap lathering with distilled water, tap water, rainwater, and solutions of MgCl₂, NaCl, Ca(NO₃)₂, CaHCO₃, NaHCO₃, ZnSO₄. Record volumes of soap needed. Boil some solutions and retest. Compare results and identify hardness-causing ions.

Soap solution, burette, various salt solutions, conical flasks, distilled water, tap water, rainwater, heating source
Student books, examples from previous experiment, chalkboard for equations

KLB Secondary Chemistry Form 4, Pages 22-24

ACIDS, BASES AND SALTS

Effects of Hard Water

By the end of the lesson, the learner should be able to:
- State disadvantages of hard water
-State advantages of hard water
-Explain formation of scum and fur
-Discuss economic and health implications

Discussion based on practical experience: Soap wastage, scum formation on clothes, fur in kettles and pipes, pipe bursting in boilers. Advantages: calcium for bones, protection of lead pipes, use in brewing. Show examples of fur deposits. Calculate economic costs of hard water in households.

Samples of fur deposits, pictures of scaled pipes, calculator for cost analysis

KLB Secondary Chemistry Form 4, Pages 24-25

ACIDS, BASES AND SALTS

Methods of Removing Hardness I
Methods of Removing Hardness II

By the end of the lesson, the learner should be able to:
- Explain removal of hardness by boiling
-Explain removal by distillation
-Write equations for these processes
-Compare effectiveness of different methods
- Explain removal using sodium carbonate
-Describe ion exchange method
-Explain removal using calcium hydroxide and ammonia
-Write equations for all processes

Demonstrate boiling method: Boil hard water samples from previous experiments and test with soap. Write equations for Ca(HCO₃)₂ and Mg(HCO₃)₂ decomposition. Discuss distillation method using apparatus setup. Compare costs and effectiveness. Explain why boiling only removes temporary hardness.
Demonstrate addition of Na₂CO₃ to hard water - observe precipitation. Explain ion exchange using resin (NaX) showing Ca²⁺ + 2NaX → CaX₂ + 2Na⁺. Discuss regeneration with brine. Write equations for Ca(OH)₂ and NH₃ methods. Compare all methods for effectiveness and cost.

Hard water samples, heating source, soap solution, distillation apparatus diagram
Na₂CO₃ solution, hard water samples, ion exchange resin diagram, Ca(OH)₂, NH₃ solution

KLB Secondary Chemistry Form 4, Pages 25-26

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Endothermic and Exothermic Reactions
Enthalpy Notation and Energy Content

By the end of the lesson, the learner should be able to:
- Define endothermic and exothermic reactions using ΔH notation
-Investigate temperature changes when ammonium nitrate and sodium hydroxide dissolve in water
-Explain observations made during dissolution
-Draw energy level diagrams for endothermic and exothermic reactions

Class experiment: Wrap 250ml plastic beakers with tissue paper. Dissolve 2 spatulafuls of NH₄NO₃ in 100ml distilled water, record temperature changes. Repeat with NaOH pellets. Compare initial and final temperatures. Draw energy level diagrams showing relative energies of reactants and products.

250ml plastic beakers, tissue paper, rubber bands, NH₄NO₃, NaOH pellets, distilled water, thermometers, spatulas, measuring cylinders
Student books, calculators, worked examples from textbook, chalkboard for calculations

KLB Secondary Chemistry Form 4, Pages 29-31

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Bond Breaking and Bond Formation

By the end of the lesson, the learner should be able to:
- Explain that energy changes are due to bond breaking and bond formation
-Describe bond breaking as endothermic and bond formation as exothermic
-Investigate energy changes during melting and boiling
-Plot heating curves for pure substances

Class experiment: Heat crushed ice while stirring with thermometer. Record temperature every minute until ice melts completely, then continue until water boils. Plot temperature-time graph. Explain constant temperature during melting and boiling in terms of bond breaking. Discuss latent heat of fusion and vaporization.

Crushed pure ice, 250ml glass beakers, thermometers, heating source, stopwatch, graph paper, stirring rods

KLB Secondary Chemistry Form 4, Pages 32-35

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Latent Heat of Fusion and Vaporization

By the end of the lesson, the learner should be able to:
- Define latent heat of fusion and molar heat of fusion
-Define latent heat of vaporization and molar heat of vaporization
-Explain why temperature remains constant during phase changes
-Relate intermolecular forces to melting and boiling points

Discussion based on previous heating curve experiment. Explain energy used to overcome intermolecular forces during melting and boiling. Compare molar heats of fusion and vaporization for water and ethanol. Relate strength of intermolecular forces to magnitude of latent heats. Calculate energy required for phase changes.

Data tables showing molar heats of fusion/vaporization, calculators, heating curves from previous lesson

KLB Secondary Chemistry Form 4, Pages 32-35

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Bond Energy Calculations
Determination of Enthalpy of Solution I
Thermochemical Equations

By the end of the lesson, the learner should be able to:
- Calculate energy changes in reactions using bond energies
-Apply the formula: Heat of reaction = Bond breaking energy + Bond formation energy
-Determine whether reactions are exothermic or endothermic
-Use bond energy data to solve problems
- Write thermochemical equations including enthalpy changes
-Define molar heat of solution
-Draw energy level diagrams for dissolution reactions
-Interpret thermochemical equations correctly

Work through formation of HCl from H₂ and Cl₂ using bond energies. Calculate energy required to break H-H and Cl-Cl bonds. Calculate energy released when H-Cl bonds form. Apply formula: ΔH = Energy absorbed - Energy released. Practice with additional examples. Discuss why calculated values may differ from experimental values.
Using data from previous experiment, write thermochemical equations for NH₄NO₃ and NaOH dissolution. Show proper notation with state symbols and ΔH values. Draw corresponding energy level diagrams. Practice writing thermochemical equations for various reactions. Explain significance of molar quantities in equations.

Bond energy data tables, calculators, worked examples, practice problems
250ml plastic beakers, 2.0g samples of NH₄NO₃ and NaOH, distilled water, thermometers, measuring cylinders, analytical balance, calculators
Results from previous experiment, graph paper for energy level diagrams, practice examples

KLB Secondary Chemistry Form 4, Pages 35-36
KLB Secondary Chemistry Form 4, Pages 38-39

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Enthalpy of Solution of Concentrated Sulphuric Acid

By the end of the lesson, the learner should be able to:
- Determine heat of solution of concentrated sulphuric(VI) acid
-Apply safety precautions when handling concentrated acids
-Calculate enthalpy change considering density and purity
-Write thermochemical equation for the reaction

Teacher demonstration: Carefully add 2cm³ concentrated H₂SO₄ to 98cm³ distilled water in wrapped beaker (NEVER vice versa). Record temperature change. Calculate mass of acid using density (1.84 g/cm³) and purity (98%). Calculate molar heat of solution. Emphasize safety - always add acid to water.

Concentrated H₂SO₄, distilled water, 250ml plastic beaker, tissue paper, measuring cylinders, thermometer, safety equipment

KLB Secondary Chemistry Form 4, Pages 39-41

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Enthalpy of Combustion
Enthalpy of Displacement

By the end of the lesson, the learner should be able to:
- Define molar heat of combustion
-Determine enthalpy of combustion of ethanol experimentally
-Explain why experimental values differ from theoretical values
-Calculate molar enthalpy of combustion from experimental data

Class experiment: Burn ethanol in small bottle with wick to heat 100cm³ water in glass beaker. Record initial and final masses of bottle+ethanol and temperature change. Calculate moles of ethanol burned and heat evolved. Determine molar enthalpy of combustion. Compare with theoretical value (-1368 kJ/mol). Discuss sources of error.

Ethanol, small bottles with wicks, 250ml glass beakers, tripod stands, wire gauze, thermometers, analytical balance, measuring cylinders
Zinc powder, 0.5M CuSO₄ solution, 250ml plastic beakers, tissue paper, thermometers, analytical balance, stirring rods

KLB Secondary Chemistry Form 4, Pages 41-44

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Enthalpy of Neutralization

By the end of the lesson, the learner should be able to:
- Define molar heat of neutralization
-Determine heat of neutralization of HCl with NaOH
-Compare neutralization enthalpies of strong and weak acids/bases
-Write ionic equations for neutralization reactions

Class experiment: Mix 50cm³ of 2M HCl with 50cm³ of 2M NaOH in wrapped beaker. Record temperature changes. Calculate molar heat of neutralization. Repeat with weak acid (ethanoic) and weak base (ammonia). Compare values. Write ionic equations. Explain why strong acid + strong base gives ~57.2 kJ/mol.

2M HCl, 2M NaOH, 2M ethanoic acid, 2M ammonia solution, measuring cylinders, thermometers, 250ml plastic beakers, tissue paper

KLB Secondary Chemistry Form 4, Pages 47-49

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Standard Conditions and Standard Enthalpy Changes
Hess's Law - Introduction and Theory
Energy Cycle Diagrams

By the end of the lesson, the learner should be able to:
- Identify standard conditions for measuring enthalpy changes
-Define standard enthalpy changes using ΔH° notation
-Explain importance of standard conditions
-Use subscripts to denote different types of enthalpy changes
- State Hess's Law
-Explain the principle of energy conservation in chemical reactions
-Understand that enthalpy change is independent of reaction route
-Apply Hess's Law to simple examples

Q/A: Review previous enthalpy measurements. Introduce standard conditions: 25°C (298K) and 1 atmosphere pressure (101.325 kPa). Explain ΔH° notation and subscripts (ΔH°c for combustion, ΔH°f for formation, etc.). Discuss why standard conditions are necessary for comparison. Practice using correct notation.
Introduce Hess's Law: "The energy change in converting reactants to products is the same regardless of the route by which the chemical change occurs." Use methane formation example to show two routes giving same overall energy change. Draw energy cycle diagrams. Explain law of conservation of energy application.

Student books, examples of standard enthalpy data, notation practice exercises
Energy cycle diagrams for methane formation, chalkboard illustrations, worked examples from textbook
Graph paper, energy cycle templates, combustion data tables, calculators

KLB Secondary Chemistry Form 4, Pages 49
KLB Secondary Chemistry Form 4, Pages 49-52

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Hess's Law Calculations

By the end of the lesson, the learner should be able to:
- Solve complex problems using Hess's Law
-Apply energy cycles to multi-step reactions
-Calculate enthalpy of formation from combustion data
-Use thermochemical equations in Hess's Law problems

Work through detailed calculation for ethanol formation: 2C(s) + 3H₂(g) + ½O₂(g) → C₂H₅OH(l). Use combustion enthalpies of carbon (-393 kJ/mol), hydrogen (-286 kJ/mol), and ethanol (-1368 kJ/mol). Calculate ΔH°f(ethanol) = -278 kJ/mol. Practice with propane and other compounds.

Worked examples, combustion data, calculators, step-by-step calculation sheets

KLB Secondary Chemistry Form 4, Pages 54-56

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Lattice Energy and Hydration Energy

By the end of the lesson, the learner should be able to:
- Define lattice energy and hydration energy
-Explain relationship between heat of solution, lattice energy and hydration energy
-Draw energy cycles for dissolution of ionic compounds
-Calculate heat of solution using Born-Haber type cycles

Explain dissolution of NaCl: first lattice breaks (endothermic), then ions hydrate (exothermic). Define lattice energy as energy to form ionic solid from gaseous ions. Define hydration energy as energy when gaseous ions become hydrated. Draw energy cycle: ΔH(solution) = ΔH(lattice) + ΔH(hydration). Calculate for NaCl.

Energy cycle diagrams, lattice energy and hydration energy data tables, calculators

KLB Secondary Chemistry Form 4, Pages 54-56

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Factors Affecting Lattice and Hydration Energies
Definition and Types of Fuels

By the end of the lesson, the learner should be able to:
- Explain factors affecting lattice energy
-Explain factors affecting hydration energy
-Use data tables to identify trends
-Calculate enthalpies of solution for various ionic compounds

Analyze data tables showing lattice energies (Table 2.7) and hydration energies (Table 2.6). Identify trends: smaller ions and higher charges give larger lattice energies and hydration energies. Calculate heat of solution for MgCl₂ using: ΔH(solution) = +2489 + (-1891 + 2×(-384)) = -170 kJ/mol. Practice with other compounds.

Data tables from textbook, calculators, trend analysis exercises
Examples of different fuels, classification charts, pictures of fuel types

KLB Secondary Chemistry Form 4, Pages 54-56

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Heating Values of Fuels
Factors in Fuel Selection

By the end of the lesson, the learner should be able to:
- Define heating value of a fuel
-Calculate heating values from molar enthalpies of combustion
-Compare heating values of different fuels
-Explain units of heating value (kJ/g)
- State factors that influence choice of fuel
-Explain why different fuels are chosen for different purposes
-Compare advantages and disadvantages of various fuels
-Apply selection criteria to real situations

Calculate heating value of ethanol: ΔH°c = -1360 kJ/mol, Molar mass = 46 g/mol, Heating value = 1360/46 = 30 kJ/g. Compare heating values from Table 2.8: methane (55 kJ/g), fuel oil (45 kJ/g), charcoal (33 kJ/g), wood (17 kJ/g). Discuss significance of these values for fuel selection.
Discuss seven factors: heating value, ease of combustion, availability, transportation, storage, environmental effects, cost. Compare wood/charcoal for domestic use vs methylhydrazine for rockets. Analyze why each is suitable for its purpose. Students suggest best fuels for cooking, heating, transport in their area.

Heating value data table, calculators, fuel comparison charts
Fuel comparison tables, local fuel availability data, cost analysis sheets

KLB Secondary Chemistry Form 4, Pages 56-57
KLB Secondary Chemistry Form 4, Pages 57

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Environmental Effects of Fuels
Fuel Safety and Precautions

By the end of the lesson, the learner should be able to:
- Identify environmental effects of burning fuels
-Explain formation and effects of acid rain
-Describe contribution to global warming
-State measures to reduce pollution from fuels

Discuss pollutants from fossil fuels: SO₂, SO₃, CO, NO₂ causing acid rain. Effects: damage to buildings, corrosion, acidification of lakes, soil leaching. CO₂ and hydrocarbons cause global warming leading to ice melting, climate change. Pollution reduction measures: catalytic converters, unleaded petrol, zero emission vehicles, alternative fuels.

Pictures of environmental damage, pollution data, examples of clean technology
Safety guideline charts, examples of fuel accidents, local safety case studies

KLB Secondary Chemistry Form 4, Pages 57-58

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Endothermic and Exothermic Reactions

By the end of the lesson, the learner should be able to:
- Define endothermic and exothermic reactions using the ΔH notation
-Investigate what happens when ammonium nitrate and sodium hydroxide are separately dissolved in water
-Define enthalpy and enthalpy change
-Calculate enthalpy changes using ΔH = H(products) - H(reactants)

Class experiment: Dissolve NH₄NO₃ and NaOH separately in water, record temperature changes in Table 2.1. Explain heat absorption vs evolution. Introduce enthalpy (H) and enthalpy change (ΔH). Calculate enthalpy changes from experimental data. Draw energy level diagrams showing relative energies.

250ml plastic beakers, tissue paper, NH₄NO₃, NaOH pellets, distilled water, thermometers, calculators

KLB Secondary Chemistry Form 4, Pages 29-32

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES

Bond Breaking, Formation and Phase Changes

By the end of the lesson, the learner should be able to:
- Explain that energy changes are due to bond breaking and bond formation
-Investigate energy changes when solids and liquids are heated
-Define latent heat of fusion and vaporization
-Calculate energy changes using bond energies

Class experiment: Heat ice to melting then boiling, record temperature every minute. Plot heating curve. Explain constant temperature periods. Define latent heat of fusion/vaporization. Calculate energy changes in H₂ + Cl₂ → 2HCl using bond energies. Apply formula: ΔH = Energy absorbed - Energy released.

Ice, glass beakers, thermometers, heating source, graph paper, bond energy data tables

KLB Secondary Chemistry Form 4, Pages 32-36

ENERGY CHANGES IN PHYSICAL AND CHEMICAL PROCESSES
REACTION RATES AND REVERSIBLE REACTIONS

Determination of Enthalpy of Solution
Definition of Reaction Rate and Collision Theory
Effect of Concentration on Reaction Rate

By the end of the lesson, the learner should be able to:
- Carry out experiments to determine enthalpy changes of solution
-Calculate enthalpy change using ΔH = mcΔT
-Write correct thermochemical equations
-Define molar heat of solution
- Explain the effect of concentration on reaction rates
-Investigate reaction of magnesium with different concentrations of sulphuric acid
-Illustrate reaction rates graphically and interpret experimental data
-Calculate concentrations and plot graphs of concentration vs time

Class experiment: Dissolve exactly 2.0g NH₄NO₃ and 2.0g NaOH separately in 100ml water. Record temperature changes. Calculate enthalpy changes using ΔH = mcΔT. Calculate moles and molar heat of solution. Write thermochemical equations: NH₄NO₃(s) + aq → NH₄NO₃(aq) ΔH = +25.2 kJ mol⁻¹.
Class experiment: Label 4 conical flasks A-D. Add 40cm³ of 2M H₂SO₄ to A, dilute others with water (30+10, 20+20, 10+30 cm³). Drop 2cm magnesium ribbon into each, time complete dissolution. Record in Table 3.1. Calculate concentrations, plot graph. Explain: higher concentration → more collisions → faster reaction.

2.0g samples of NH₄NO₃ and NaOH, plastic beakers, thermometers, analytical balance, calculators
Examples of fast/slow reactions, energy diagram templates, chalk/markers for diagrams
4 conical flasks, 2M H₂SO₄, distilled water, magnesium ribbon, stopwatch, measuring cylinders, graph paper

KLB Secondary Chemistry Form 4, Pages 36-39
KLB Secondary Chemistry Form 4, Pages 65-67

REACTION RATES AND REVERSIBLE REACTIONS

Change of Reaction Rate with Time

By the end of the lesson, the learner should be able to:
- Describe methods used to measure rate of reaction
-Investigate how reaction rate changes as reaction proceeds
-Plot graphs of volume of gas vs time
-Calculate average rates at different time intervals

Class experiment: React 2cm magnesium ribbon with 100cm³ of 0.5M HCl in conical flask. Collect H₂ gas in graduated syringe as in Fig 3.4. Record gas volume every 30 seconds for 5 minutes in Table 3.2. Plot volume vs time graph. Calculate average rates between time intervals. Explain why rate decreases as reactants are consumed.

0.5M HCl, magnesium ribbon, conical flask, gas collection apparatus, graduated syringe, stopwatch, graph paper

KLB Secondary Chemistry Form 4, Pages 67-70

REACTION RATES AND REVERSIBLE REACTIONS

Effect of Temperature on Reaction Rate
Effect of Surface Area on Reaction Rate

By the end of the lesson, the learner should be able to:
- Explain the effect of temperature on reaction rates
-Investigate temperature effects using sodium thiosulphate and HCl
-Plot graphs of time vs temperature and 1/time vs temperature
-Apply collision theory to explain temperature effects

Class experiment: Place 30cm³ of 0.15M Na₂S₂O₃ in flasks at room temp, 30°C, 40°C, 50°C, 60°C. Mark cross on paper under flask. Add 5cm³ of 2M HCl, time until cross disappears. Record in Table 3.4. Plot time vs temperature and 1/time vs temperature graphs. Explain: higher temperature → more kinetic energy → more effective collisions.

0.15M Na₂S₂O₃, 2M HCl, conical flasks, water baths at different temperatures, paper with cross marked, stopwatch, thermometers
Marble chips, marble powder, 1M HCl, gas collection apparatus, balance, conical flasks, measuring cylinders, graph paper

KLB Secondary Chemistry Form 4, Pages 70-73

REACTION RATES AND REVERSIBLE REACTIONS

Effect of Catalysts on Reaction Rate

By the end of the lesson, the learner should be able to:
- Explain effects of suitable catalysts on reaction rates
-Investigate decomposition of hydrogen peroxide with and without catalyst
-Define catalyst and explain how catalysts work
-Compare activation energies in catalyzed vs uncatalyzed reactions

Class experiment: Decompose 5cm³ of 20-volume H₂O₂ in 45cm³ water without catalyst, collect O₂ gas. Repeat adding 2g MnO₂ powder. Record gas volumes as in Fig 3.12. Compare rates and final mass of MnO₂. Write equation: 2H₂O₂ → 2H₂O + O₂. Define catalyst and explain how it lowers activation energy. Show energy diagrams for both pathways.

20-volume H₂O₂, MnO₂ powder, gas collection apparatus, balance, conical flasks, filter paper, measuring cylinders

KLB Secondary Chemistry Form 4, Pages 76-78

REACTION RATES AND REVERSIBLE REACTIONS

Effect of Light and Pressure on Reaction Rate
Reversible Reactions

By the end of the lesson, the learner should be able to:
- Identify reactions affected by light
-Investigate effect of light on silver bromide decomposition
-Explain effect of pressure on gaseous reactions
-Give examples of photochemical reactions
- State examples of simple reversible reactions
-Investigate heating of hydrated copper(II) sulphate
-Write equations for reversible reactions using double arrows
-Distinguish between reversible and irreversible reactions

Teacher demonstration: Mix KBr and AgNO₃ solutions to form AgBr precipitate. Divide into 3 test tubes: place one in dark cupboard, one on bench, one in direct sunlight. Observe color changes after 10 minutes. Write equations. Discuss photochemical reactions: photography, Cl₂ + H₂, photosynthesis. Explain pressure effects on gaseous reactions through compression.
Class experiment: Heat CuSO₄·5H₂O crystals in boiling tube A, collect liquid in tube B as in Fig 3.15. Observe color changes: blue → white + colorless liquid. Pour liquid back into tube A, observe return to blue. Write equation with double arrows: CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O. Give other examples: NH₄Cl ⇌ NH₃ + HCl. Compare with irreversible reactions.

0.1M KBr, 0.05M AgNO₃, test tubes, dark cupboard, direct light source, examples of photochemical reactions
CuSO₄·5H₂O crystals, boiling tubes, delivery tube, heating source, test tube holder

KLB Secondary Chemistry Form 4, Pages 78-80

REACTION RATES AND REVERSIBLE REACTIONS

Chemical Equilibrium
Le Chatelier's Principle and Effect of Concentration

By the end of the lesson, the learner should be able to:
- Explain chemical equilibrium
-Define dynamic equilibrium
-Investigate acid-base equilibrium using indicators
-Explain why equilibrium appears static but is actually dynamic

Experiment: Add 0.5M NaOH to 2cm³ in boiling tube with universal indicator. Add 0.5M HCl dropwise until green color (neutralization point). Continue adding base then acid alternately, observe color changes. Explain equilibrium as state where forward and backward reaction rates are equal. Use NH₄Cl ⇌ NH₃ + HCl example to show dynamic nature. Introduce equilibrium symbol ⇌.

0.5M NaOH, 0.5M HCl, universal indicator, boiling tubes, droppers, examples of equilibrium systems
Bromine water, 2M NaOH, 2M HCl, beakers, chromate/dichromate solutions for demonstration

KLB Secondary Chemistry Form 4, Pages 80-82

REACTION RATES AND REVERSIBLE REACTIONS

Effect of Pressure and Temperature on Equilibrium

By the end of the lesson, the learner should be able to:
- Explain effect of pressure changes on equilibrium
-Explain effect of temperature changes on equilibrium
-Investigate NO₂/N₂O₄ equilibrium with temperature
-Apply Le Chatelier's Principle to industrial processes

Teacher demonstration: React copper turnings with concentrated HNO₃ to produce NO₂ gas in test tube. Heat and cool the tube, observe color changes: brown ⇌ pale yellow representing 2NO₂ ⇌ N₂O₄. Explain pressure effects using molecule count. Show Table 3.7 with pressure effects. Discuss temperature effects: heating favors endothermic direction, cooling favors exothermic direction. Use Table 3.8.

Copper turnings, concentrated HNO₃, test tubes, heating source, ice bath, gas collection apparatus, safety equipment

KLB Secondary Chemistry Form 4, Pages 84-87

REACTION RATES AND REVERSIBLE REACTIONS

Industrial Applications - Haber Process

By the end of the lesson, the learner should be able to:
- Apply equilibrium principles to Haber Process
-Explain optimum conditions for ammonia manufacture
-Calculate effect of temperature and pressure on yield
-Explain role of catalysts in industrial processes

Analyze Haber Process: N₂ + 3H₂ ⇌ 2NH₃ ΔH = -92 kJ/mol. Apply Le Chatelier's Principle: high pressure favors forward reaction (4 molecules → 2 molecules), low temperature favors exothermic forward reaction but slows rate. Explain optimum conditions: 450°C temperature, 200 atmospheres pressure, iron catalyst. Discuss removal of NH₃ to shift equilibrium right. Economic considerations.

Haber Process flow diagram, equilibrium data showing temperature/pressure effects on NH₃ yield, industrial catalyst information

KLB Secondary Chemistry Form 4, Pages 87-89

REACTION RATES AND REVERSIBLE REACTIONS
ELECTROCHEMISTRY
ELECTROCHEMISTRY

Industrial Applications - Contact Process
Redox Reactions and Oxidation Numbers
Oxidation Numbers in Naming and Redox Identification
Displacement Reactions - Metals and Halogens
Electrochemical Cells and Cell Diagrams
Standard Electrode Potentials

By the end of the lesson, the learner should be able to:
- Apply equilibrium principles to Contact Process
-Explain optimum conditions for sulphuric acid manufacture
-Compare different industrial equilibrium processes
-Evaluate economic factors in industrial chemistry
Explain displacement reactions using electron transfer
- Arrange metals and halogens by reactivity
- Predict displacement reactions
- Compare oxidizing powers of halogens

Analyze Contact Process: 2SO₂ + O₂ ⇌ 2SO₃ ΔH = -197 kJ/mol. Apply principles: high pressure favors forward reaction (3 molecules → 2 molecules), low temperature favors exothermic reaction. Explain optimum conditions: 450°C, atmospheric pressure, V₂O₅ catalyst, 96% conversion. Compare with Haber Process. Discuss catalyst choice and economic factors.
Experiment 4.3: Metal displacement reactions - systematic testing
- Experiment 4.4: Halogen displacement (FUME CUPBOARD)
- Tabulate results and arrange by reactivity

Contact Process flow diagram, comparison table with Haber Process, catalyst effectiveness data
Iron filings, 1M CuSO₄, 1M FeSO₄, 2M NaOH, 20V H₂O₂, test tubes
Compound charts, calculators, student books, practice exercises
Various metals (Ca, Mg, Zn, Fe, Pb, Cu), metal salt solutions, halogens (Cl₂, Br₂, I₂), halide solutions
Metal electrodes, 1M metal salt solutions, voltmeters, salt bridges, connecting wires
Standard electrode potential table, diagrams, charts showing standard conditions

KLB Secondary Chemistry Form 4, Pages 89
KLB Secondary Chemistry Form 4, Pages 116-122

ELECTROCHEMISTRY

Calculating Cell EMF and Predicting Reactions
Types of Electrochemical Cells

By the end of the lesson, the learner should be able to:
Calculate EMF using standard electrode potentials
- Predict reaction spontaneity using EMF
- Solve numerical problems on cell EMF
- Apply EMF calculations practically

Worked examples: Calculate EMF for various cells
- Practice EMF calculations
- Exercise 4.2 & 4.3: Cell EMF and reaction feasibility problems
- Distinguish spontaneous from non-spontaneous reactions

Calculators, electrode potential data, worked examples, practice problems
Cell diagrams, sample batteries, charts showing cell applications

KLB Secondary Chemistry Form 4, Pages 133-137

ELECTROCHEMISTRY

Electrolysis of Aqueous Solutions I
Electrolysis of Aqueous Solutions II

By the end of the lesson, the learner should be able to:
Define electrolysis and preferential discharge
- Investigate electrolysis of dilute sodium chloride
- Compare dilute vs concentrated solution effects
- Test products formed

Experiment 4.6(a): Electrolysis of dilute NaCl
- Experiment 4.6(b): Electrolysis of brine
- Test gases evolved
- Compare results and explain differences

Dilute and concentrated NaCl solutions, carbon electrodes, gas collection tubes, test equipment
U-tube apparatus, 2M H₂SO₄, 0.5M MgSO₄, platinum/carbon electrodes, gas syringes

KLB Secondary Chemistry Form 4, Pages 141-146

ELECTROCHEMISTRY

Effect of Electrode Material on Electrolysis
Factors Affecting Electrolysis

By the end of the lesson, the learner should be able to:
Compare inert vs reactive electrodes
- Investigate electrode dissolution
- Explain electrode selection importance
- Analyze copper purification process

Experiment 4.9: Electrolysis of CuSO₄ with carbon vs copper electrodes
- Weigh electrodes before/after
- Observe color changes
- Discussion on electrode effects

Copper and carbon electrodes, 3M CuSO₄ solution, accurate balance, beakers, connecting wires
Electrochemical series chart, summary tables, practice exercises, student books

KLB Secondary Chemistry Form 4, Pages 141-148

ELECTROCHEMISTRY

Applications of Electrolysis I
Applications of Electrolysis II

By the end of the lesson, the learner should be able to:
Describe electrolytic extraction of reactive metals
- Explain electroplating process
- Apply electrolysis principles to metal coating
- Design electroplating setup
Describe manufacture of NaOH and Cl₂ from brine
- Explain mercury cell operation
- Analyze industrial electrolysis processes
- Discuss environmental considerations

Discussion: Extraction of Na, Mg, Al by electrolysis
- Practical: Electroplate iron nail with copper
- Calculate plating requirements
- Industrial applications
Study mercury cell for NaOH production
- Flow chart analysis of industrial processes
- Discussion on applications and environmental impact
- Purification of metals

Iron nails, copper electrodes, CuSO₄ solution, power supply, industrial process diagrams
Flow charts, mercury cell diagrams, environmental impact data, industrial case studies

KLB Secondary Chemistry Form 4, Pages 155-157

ELECTROCHEMISTRY

Faraday's Laws and Quantitative Electrolysis
Electrolysis Calculations I

By the end of the lesson, the learner should be able to:
State Faraday's laws of electrolysis
- Define Faraday constant
- Calculate mass deposited in electrolysis
- Relate electricity to amount of substance

Experiment 4.10: Quantitative electrolysis of CuSO₄
- Measure mass vs electricity passed
- Calculate Faraday constant
- Verify Faraday's laws

Accurate balance, copper electrodes, CuSO₄ solution, ammeter, timer, calculators
Calculators, worked examples, practice problems, gas volume data, Faraday constant

KLB Secondary Chemistry Form 4, Pages 161-164

ELECTROCHEMISTRY

Electrolysis Calculations II
Advanced Applications and Problem Solving

By the end of the lesson, the learner should be able to:
Determine charge on ions from electrolysis data
- Calculate current-time relationships
- Solve complex multi-step problems
- Apply concepts to industrial situations

Complex problems: Determine ionic charges
- Current-time-mass relationships
- Multi-step calculations
- Industrial calculation examples

Calculators, complex problem sets, industrial data, student books
Past papers, comprehensive problem sets, industrial case studies, calculators

KLB Secondary Chemistry Form 4, Pages 161-164