Chemistry Notes for Senior High Schools

MINISTRY OF EDUCATION GHANA

  - Chemistry for Senior High Schools
    - Authors: Alhassan Kuvidana Mustapha, Christian Dzikunu, Delali Robert Akplai, Yirwelle Santus
    - Year 2

Table of Contents

• Foreword

• Section 1: Energy Changes
    - Physical Chemistry
    - Matter and Its Properties

• Section 2: Chemical Kinetics
    - Physical Chemistry
    - Matter and Its Properties

• Section 3: Dynamic Equilibrium
    - Physical Chemistry
    - Equilibria

• Section 4: Acids, Bases and Salts
    - Physical Chemistry
    - Equilibria

• Section 5: Trends of Chemical and Physical Properties of Elements and Their Compounds in the Periodic Table
    - Systematic Chemistry of the Elements
    - Periodicity

• Section 6: Physical and Chemical Properties of the Halogens

• Section 7: Structure, Chemical Bonding and Properties of Molecular Compounds

• Section 8: Organic Compounds

• Bibliography

• Glossary

Foreword

• Ghana’s new Senior High School Curriculum aims to ensure learners achieve their potential by imparting 21st Century skills and education in line with national values.

• Materials developed through partnerships, emphasizing quality education.

• Objectives include preparing learners for responsible adulthood and contributing to national development.

Section 1: Energy Changes

Physical Chemistry

Matter and Its Properties

Introduction

• Learning about energy changes during chemical reactions (enthalpy).

• Investigating Hess’s Law, real-world applications, problem-solving, scientific communication.

Key Ideas

• Atomisation Energy: Energy to break all bonds in a mole to form gaseous atoms.

• Bond Dissociation Energy: Energy to break a specific bond in a compound.

• Bond Enthalpy: Average energy to split one mole of covalent bonds in gaseous molecules.

• Enthalpy Change (ΔH): Difference in enthalpy between products and reactants.

• Hess’s Law: Total enthalpy change equals sum of individual enthalpy changes.

• Lattice Energy: Energy to separate one mole of ionic solid into gaseous ions.

Enthalpy Change

• System Definition: The part of the universe focused on in an experiment, surroundings encompass everything else.

• Types of Systems:
    1. Open: Exchanges heat and matter (e.g., heating water in open beaker).
    2. Closed: Exchanges heat, not matter (e.g., covered beaker).
    3. Isolated: Exchanges neither heat nor matter (e.g., calorimeter).

• Enthalpy (H) Equation:
    $H = U + PV$ where $U$ is internal energy, $P$ is pressure, and $V$ is volume.

• When pressure is constant, enthalpy represents heat exchange during reactions.

• Enthalpy Change (ΔH) Formula:
  $ext{ΔH} = H_{ ext{products}} - H_{ ext{reactants}}$

Endothermic and Exothermic Reactions

• Exothermic Reactions:
    - Release heat to surroundings (ΔH < 0).
    - Example: $ext{CH}_4(g) + 2 ext{O}_2(g) <br>ightarrow ext{CO}_2(g) + 2 ext{H}_2 ext{O}(l) ext{ with } ΔH = -890.4 kJ/mol$.

• Endothermic Reactions:
    - Absorb heat from surroundings (ΔH > 0).
    - Example: $ext{CaCO}_3(s) <br>ightarrow ext{CaO}(s) + ext{CO}_2(g) ext{ with } ΔH = +177.8 kJ/mol$.

Visual Representation

• Energy Profile Diagrams depict energy changes in both types of reactions.

Activity 1.1: Demonstrating Chemical Systems

• Materials: 7 plastic bottles, hot water, thermometer.

• Steps:
    1. Set up open, closed, and isolated systems with hot water.
    2. Measure initial temperatures and record time intervals.
    3. Predict and observe temperature changes at each system type.

Activity 1.2: Exploring Enthalpy

• Materials: Markers, chart paper, endothermic/exothermic video.

• Steps: Discuss key terms, provide real-life examples, and analyze implications.

Standard Enthalpy Changes

• Standard conditions: 1 atm pressure, 298 K temperature.

• Units: kilojoules per mole (kJ/mol).

• Standard Enthalpy Change of Reaction (ΔH°): Measured under standard conditions; dependent on stoichiometry.

Types of Standard Enthalpy Changes

1. Standard Enthalpy Change of Formation: Formation of one mole of compound from elements.

2. Standard Enthalpy Change of Combustion: Enthalpy change during complete combustion of a substance.

3. Standard Enthalpy Change of Neutralization: Enthalpy change when acids react with bases.

4. Standard Enthalpy Change of Solution: Enthalpy change on dissolving one mole in excess solvent.

5. Standard Enthalpy Change of Hydration: Enthalpy change when gaseous ions dissolve in water.

Activity 1.4: Understanding Types of Enthalpy Changes

• Define each type and provide examples.

• Calculate standard enthalpy changes using Hess’s law.

Section 2: Chemical Kinetics

Introduction

• Explores measuring reaction rates and understanding factors affecting these rates.

• Key concepts include average and instantaneous rates, average rate equations, and Catalysts (substances that increase reaction rates without being consumed).

Rate of Reaction

• Defined as the change in concentration or moles of a reactant/product per unit time.

• Generalized based on reactions: $rate = - rac{1}{a} rac{Δ[A]}{Δt} = - rac{1}{b} rac{Δ[B]}{Δt}$ etc.

Initial Rate of a Reaction

• The rate at the onset of reaction typically determined by monitoring the change over small time intervals.

Average Rate of Reaction

• Defined over larger time intervals by comparing the concentration changes over defined periods.

Activity 2.1: Defining Reaction Rate

• Define reaction rate, explore measuring methods, and rate-limiting factors.

Temperature and Reaction Rates

• Increasing temperature generally accelerates reaction rates due to increased kinetic energy, yielding more collisions.

Concentration and Surface Area Effects

• Higher concentration leads to more collisions.

• Increased surface area enhances reactivity in solid reactants.

Catalyst and Pressure Effects

• Catalysts significantly boost reaction speeds by providing alternate routes for reactions.

• In gaseous reactions, increased pressure raises reaction rates by reducing volume and increasing collision frequency.

Le Chatelier’s Principle

• Describes how changes in concentration, pressure, or temperature can shift equilibrium positions, with applications in industrial processes (e.g., Haber Process for ammonia).

Activity 2.10: Visualising Activation Energy

• Reflect on the analogy that illustrates activation energy; explore and illustrate correlated variables.

Section 3: Dynamic Equilibrium

(Continue with a similar detailed format for each section…)