06.00 Chapter Summary

1. Electrolysis Principles

General Principles of Electrolysis

• Electrolysis: A process where an electric current is passed through a molten or aqueous ionic compound, causing it to decompose into its constituent elements.
  • Molten Ionic Compounds: Conduct electricity as ions are free to move.
  • Aqueous Solutions: Conduct electricity due to the presence of free ions from both the dissolved ionic compound and water.
  • Covalent Compounds: Do not conduct electricity and thus do not undergo electrolysis.
  • Solid Ionic Compounds: Do not conduct electricity because ions are fixed in place and cannot move.

Key Components of an Electrolytic Cell

• Electrodes:
  • Anode: Positive electrode where oxidation occurs; attracts anions (negative ions).
  • Cathode: Negative electrode where reduction occurs; attracts cations (positive ions).
  • Material: Typically made of metal or graphite. Inert electrodes (e.g., graphite or platinum) are preferred to prevent side reactions.
• Electrolyte: The molten or dissolved ionic compound that allows ion movement and conducts electricity.

Charge Transfer in Electrolysis

• External Circuit: Electrons flow from the anode to the cathode via the power supply.
• Internal Circuit (Electrolyte):
  • Cations (e.g., Na⁺, Cu²⁺) move to the cathode, gain electrons, and are reduced.
  • Anions (e.g., Cl⁻, Br⁻) move to the anode, lose electrons, and are oxidized.
• Charge Carriers:
  • External Circuit: Electrons.
  • Electrolyte: Ions (cations and anions).

Mnemonics to Remember Key Concepts

• PANIC:
  • Positive = Anode
  • Negative = Cathode
• OIL RIG:
  • Oxidation Is Loss (of electrons)
  • Reduction Is Gain (of electrons)
• RED CAT & AN OX:
  • REDuction at the CAthode
  • ANode for OXidation

2. Electrolysis of Molten Compounds

Binary Ionic Compounds

• Definition: Compounds consisting of two elements bonded ionically (e.g., PbBr₂, KCl).
• Electrolysis Products:
  • Cathode: Metal (reduction of cations).
  • Anode: Non-metal (oxidation of anions).

Example: Electrolysis of Molten Lead(II) Bromide (PbBr₂)

1. Setup:
   • Electrolyte: Molten PbBr₂.
   • Electrodes: Two graphite rods connected to a power supply.
2. Process:
   • Anode (Positive Electrode):
     • Reaction: 2Br⁻ → Br₂(g) + 2e⁻
     • Observation: Bubbling of brown bromine gas.
   • Cathode (Negative Electrode):
     • Reaction: Pb²⁺ + 2e⁻ → Pb(s)
     • Observation: Deposition of grey lead metal on the cathode.

Worked Example: Electrolysis of Molten Potassium Chloride (KCl)

• Ions Present: K⁺ and Cl⁻.
• Products:
  • Anode: 2Cl⁻ → Cl₂(g) + 2e⁻ (Chlorine gas is released).
  • Cathode: K⁺ + e⁻ → K(s) (Potassium metal is deposited).

3. Electrolysis of Aqueous Solutions

Understanding Aqueous Electrolysis

• Ion Sources:
  • From the Electrolyte: e.g., Na⁺, Cl⁻ in NaCl solution.
  • From Water: H⁺ and OH⁻ ions from H₂O ⇌ H⁺ + OH⁻.
• Product Formation: Depends on the relative reactivity and concentration of ions.

Electrolysis of Aqueous Sodium Chloride (Brine)

• Electrolyte: Brine (concentrated NaCl solution).
• Ions Present: Na⁺, Cl⁻, H⁺, OH⁻.
• Products:
  • Anode (Positive Electrode):
    • Reaction: 2Cl⁻ → Cl₂(g) + 2e⁻ (Chlorine gas is released).
  • Cathode (Negative Electrode):
    • Reaction: 2H⁺ + 2e⁻ → H₂(g) (Hydrogen gas is released).
  • Remaining Solution: NaOH (Sodium hydroxide).

Electrolysis of Dilute Sulfuric Acid (H₂SO₄)

• Electrolyte: Dilute H₂SO₄ solution.
• Ions Present: H⁺, SO₄²⁻, OH⁻.
• Products:
  • Anode (Positive Electrode):
    • Reaction: 4OH⁻ → O₂(g) + 2H₂O(l) + 4e⁻ (Oxygen gas is released).
  • Cathode (Negative Electrode):
    • Reaction: 2H⁺ + 2e⁻ → H₂(g) (Hydrogen gas is released).

Effect of Ion Concentration on Product Formation

• Concentrated Solutions:
  • Anions from Electrolyte: More likely to be discharged (e.g., Cl⁻ in concentrated NaCl).
  • Example: Concentrated NaCl produces Cl₂ at the anode.
• Dilute Solutions:
  • Water Ions: More likely to be discharged (e.g., OH⁻ in dilute NaCl produces O₂ at the anode).
  • Example: Dilute NaCl produces O₂ at the anode.

Electrolysis of Aqueous Copper(II) Sulfate (CuSO₄)

• Using Graphite Electrodes:
  • Anode (Positive Electrode):
    • Reaction: 4OH⁻ → O₂(g) + 2H₂O(l) + 4e⁻ (Oxygen gas is released).
  • Cathode (Negative Electrode):
    • Reaction: Cu²⁺ + 2e⁻ → Cu(s) (Copper metal is deposited).
• Using Copper Electrodes:
  • Anode (Positive Electrode):
    • Reaction: Cu(s) → Cu²⁺ + 2e⁻ (Copper atoms are oxidized into Cu²⁺ ions).
  • Cathode (Negative Electrode):
    • Reaction: Cu²⁺ + 2e⁻ → Cu(s) (Copper metal is deposited).
  • Observation:
    • Cathode: Increases in mass due to copper deposition.
    • Anode: Decreases in mass due to copper oxidation.
    • Solution: Concentration of Cu²⁺ remains constant.

Gas Tests for Identifying Products

• Oxygen: Relights a glowing splint.
• Hydrogen: Produces a squeaky pop when a lit splint is introduced.
• Chlorine: Turns damp litmus paper red, then bleaches it white.

4. Ionic Half Equations

Basic Concepts

• Oxidation: Loss of electrons; occurs at the anode.
• Reduction: Gain of electrons; occurs at the cathode.
• Ionic Half Equations:
  • Oxidation Half-Equation: Shows loss of electrons.
  • Reduction Half-Equation: Shows gain of electrons.
• Balancing: Ensure atoms and charges are balanced in each half-equation.

Examples of Ionic Half Equations

• Metals (Reduction):
  • Lithium: Li⁺ + e⁻ → Li(s)
  • Copper: Cu²⁺ + 2e⁻ → Cu(s)
  • Aluminum: Al³⁺ + 3e⁻ → Al(s)
• Non-Metals (Oxidation and Reduction):
  • Hydrogen:
    • Reduction: 2H⁺ + 2e⁻ → H₂(g)
  • Halides (e.g., Chloride):
    • Oxidation: 2Cl⁻ → Cl₂(g) + 2e⁻
  • Hydroxide:
    • Oxidation: 4OH⁻ → O₂(g) + 2H₂O(l) + 4e⁻

Table of Common Ionic Half Equations

5. Electroplating

Process Overview

• Definition: Electroplating involves coating the surface of one metal (object) with a layer of another metal using electrolysis.
• Components:
  • Cathode: The object to be plated (negative electrode).
  • Anode: Pure metal that will be deposited onto the cathode (positive electrode).
  • Electrolyte: Aqueous solution of a soluble salt of the pure metal (e.g., SnCl₂ for tin plating).

Example: Electroplating Iron with Tin

1. Setup:
   • Electrolyte: Tin(II) chloride (SnCl₂) solution.
   • Anode (Positive Electrode): Pure tin (Sn).
   • Cathode (Negative Electrode): Iron strip (Fe).
2. Process:
   • Anode Reaction (Oxidation): Sn(s) → Sn²⁺(aq) + 2e⁻
   • Cathode Reaction (Reduction): Sn²⁺(aq) + 2e⁻ → Sn(s)
   • Outcome: Tin atoms deposit onto the iron strip, coating it with tin.
3. Observations:
   • Cathode: Gains mass as tin is deposited.
   • Anode: Loses mass as tin dissolves into the electrolyte.

Applications of Electroplating

• Corrosion Resistance:
  • Examples: Chromium or nickel plating to prevent rust.
• Aesthetic Enhancement:
  • Examples: Silver plating on cutlery and jewellery for a shiny appearance.
• Protection:
  • Examples: Galvanizing iron with zinc to protect against corrosion.

Examiner Tips for Electroplating

• Writing Ionic Half-Equations: Be prepared to write balanced half-equations for both anode and cathode reactions.
  • Example:
    • Anode: Sn(s) → Sn²⁺(aq) + 2e⁻
    • Cathode: Sn²⁺(aq) + 2e⁻ → Sn(s)
• Understanding Mass Changes: Know that the mass gained at the cathode equals the mass lost at the anode, maintaining ion concentration.

6. Hydrogen Fuel Cells

Functionality

• Definition: A hydrogen fuel cell is an electrochemical cell that converts hydrogen and oxygen into water, producing electricity in the process.
• Operation:
  • Anode Reaction: H₂ → 2H⁺ + 2e⁻
  • Cathode Reaction: O₂ + 4e⁻ → 2O²⁻
  • Overall Reaction: 2H₂ + O₂ → 2H₂O
• Setup:
  • Fuel (Hydrogen): Enters the anode side.
  • Oxidizer (Oxygen): Enters the cathode side.
  • Product: Water is the only chemical byproduct.

Advantages of Hydrogen Fuel Cells

• Renewable Energy Source: Hydrogen can be produced from water via electrolysis.
• Environmentally Friendly: The only byproduct is water, eliminating harmful emissions.
• High Energy Yield: Hydrogen has a higher energy content per kilogram compared to petrol or diesel.
• Efficient Transmission: No power loss through moving parts, enhancing efficiency.
• Quiet Operation: Reduces noise pollution compared to internal combustion engines.

Disadvantages of Hydrogen Fuel Cells

• Production Challenges:
  • Fossil Fuel Dependency: Traditional hydrogen production methods release CO₂.
  • Energy-Intensive: Electrolysis requires substantial electricity input.
• Storage Issues:
  • Safety: Hydrogen is highly flammable and explosive under pressure.
  • Cost: Storing hydrogen safely is expensive.
• Material Costs: Components of fuel cells (e.g., platinum catalysts) are expensive.
• Infrastructure Limitations: Limited number of hydrogen refueling stations available.
• Performance at Low Temperatures: Efficiency decreases in cold environments.

Examiner Tips for Hydrogen Fuel Cells

• Advantages and Disadvantages: Be prepared to compare hydrogen fuel cells with traditional petrol or diesel engines.
• Balanced Understanding: Recognize both the environmental benefits and the practical challenges involved in hydrogen fuel cell technology.

7. Endothermic & Exothermic Reactions

Basic Concepts

• System: The reacting chemicals involved in the reaction.
• Surroundings: Everything outside the system.
• Energy Changes: Result from the breaking and forming of chemical bonds.

Exothermic Reactions

• Definition: Reactions that release thermal energy to the surroundings.
• Energy Change: Negative (energy of the system decreases).
• Temperature Effect: Surroundings increase in temperature.
• Characteristics:
  • Energy Transfer: From the system to the surroundings.
  • Common Examples:
    • Combustion: Burning fuels (e.g., burning methane).
    • Oxidation: Rusting of iron.
    • Neutralization: Mixing acids and bases.
    • Practical Uses: Hand warmers, self-heating cans.
• Diagram Representation:
  • Arrows indicating heat moving from the system to the surroundings.

Endothermic Reactions

• Definition: Reactions that absorb thermal energy from the surroundings.
• Energy Change: Positive (energy of the system increases).
• Temperature Effect: Surroundings decrease in temperature.
• Characteristics:
  • Energy Transfer: From the surroundings to the system.
  • Common Examples:
    • Electrolysis: Decomposition of compounds using electricity.
    • Thermal Decomposition: Breaking down compounds with heat.
    • Photosynthesis: Conversion of CO₂ and H₂O into glucose and O₂.
    • Practical Uses: Cold packs for injuries.
• Diagram Representation:
  • Arrows indicating heat moving from the surroundings to the system.

Worked Example: Identifying Reaction Types Based on Temperature Change

Explanation:

• Reactions 1, 3, and 4: Temperature increased → Exothermic
• Reaction 2: Temperature decreased → Endothermic

Examiner Tips for Endo/Exothermic Reactions

• Understanding Energy Flow: Clearly distinguish whether energy is absorbed or released.
• Real-Life Examples: Be familiar with practical applications like hand warmers (exothermic) and cold packs (endothermic).
• Reaction Pathway Diagrams: Know how to interpret and draw diagrams showing energy changes.

8. Reaction Pathway Diagrams

Definition

• Reaction Pathway Diagrams: Graphical representations showing the relative energies of reactants and products during a chemical reaction.

Components of the Diagram

• X-Axis: Progress of the reaction from reactants to products.
• Y-Axis: Energy level.
• Reactants: Starting materials shown at their initial energy level.
• Products: Final materials shown at their final energy level.
• Activation Energy: The energy barrier that must be overcome for the reaction to proceed.
• Overall Energy Change: Difference in energy between reactants and products.
  • Exothermic: Products are at a lower energy level than reactants.
  • Endothermic: Products are at a higher energy level than reactants.

Interpreting the Diagram

• Exothermic Reaction:
  • Energy decreases from reactants to products.
  • The diagram slopes downward from reactants to products.
• Endothermic Reaction:
  • Energy increases from reactants to products.
  • The diagram slopes upward from reactants to products.

9. Exam Tips and Tricks

Key Mnemonics

• PANIC: Positive = Anode, Negative = Cathode.
• OIL RIG: Oxidation Is Loss, Reduction Is Gain.
• RED CAT & AN OX:
  • REDuction at the CAthode.
  • ANode for OXidation.

Gas Identification Tests

• Oxygen:
  • Test: Relights a glowing splint.
• Hydrogen:
  • Test: Produces a squeaky pop when a lit splint is introduced.
• Chlorine:
  • Test: Turns damp litmus paper red, then bleaches it white.

Electrode Selection

• Inert Electrodes: Use graphite or platinum to prevent side reactions.
• Active Electrodes: Metals like copper can participate in reactions (e.g., copper ions are dissolved or deposited).

Balancing Ionic Half Equations

• Ensure Both Atoms and Charges Are Balanced:
  • Count the number of each type of atom on both sides.
  • Ensure the total charge is equal on both sides of the equation.

Understanding Reaction Pathways

• Energy Levels: Know how to interpret energy changes in reaction pathway diagrams.
• Activation Energy: Recognize the energy barrier required for reactions to proceed.

Practical Applications

• Electroplating: Understand the purpose (e.g., corrosion resistance, aesthetic enhancement) and the process (anode and cathode reactions).
• Hydrogen Fuel Cells: Be familiar with their operation, advantages, and disadvantages compared to traditional engines.

Additional Tips

• Draw Diagrams: When required, draw clear and labeled diagrams to illustrate concepts like electrolytic cells and fuel cells.
• Stay Organized: Structure answers logically, especially for multi-part questions involving reactions and equations.
• Practice Past Papers: Familiarize yourself with the format and types of questions typically asked in exams.