Chapter 8 – Redox Reactions
1. Introduction to Redox Reactions
Definition
Redox reactions are chemical reactions where oxidation (loss of electrons) and reduction (gain of electrons) occur simultaneously.
Why We Study This
Redox reactions explain energy transfer in batteries, metal corrosion, and biological processes like respiration. Without understanding redox, we couldn’t design fuels or prevent rust!
Detailed Explanation
In cellular respiration: C6H12O6 + 6O2 → 6CO2 + 6H2O, glucose is oxidized (loses H) while oxygen is reduced (gains H).
2. Oxidation and Reduction
Definition
Oxidation = Loss of electrons
Reduction = Gain of electrons
Why We Need This Concept
To identify which substances act as oxidizing/reducing agents in reactions. This helps predict reaction products.
Detailed Explanation
Example: Zn + CuSO4 → ZnSO4 + Cu
• Zn oxidizes (loses 2e⁻ → Zn²⁺)
• Cu²⁺ reduces (gains 2e⁻ → Cu)
3. Rules to Assign Oxidation Number
Definition
The oxidation number is the hypothetical charge an atom would have if all bonds were ionic.
Why We Calculate It
To track electron transfer in reactions and balance complex equations.
Detailed Rules
1. Free elements: O.N. = 0
2. Oxygen: Usually -2 (except peroxides)
3. Hydrogen: +1 (except in metal hydrides)
Example: In HNO3, H=+1, O=-2 ⇒ N=+5
4. Balancing Redox Reactions
Definition
A method to ensure atoms and charges are conserved in redox equations.
Why It’s Essential
Unbalanced equations give incorrect stoichiometry, leading to errors in lab calculations.
Step-by-Step Method
In Acidic Medium:
1. Split into half-reactions
2. Balance O with H2O
3. Balance H with H⁺
4. Balance charges with electrons
Example: Cr2O7²⁻ → Cr³⁺ (balanced with 14H⁺ and 6e⁻)
5. Types of Redox Reactions
Definition
Redox reactions are categorized based on how electrons are transferred between substances.
Why Classify Them?
Classification helps predict reaction behavior and products, which is vital for industrial processes like metal extraction.
Detailed Types
1. Combination:
Two substances combine: 2Mg + O2 → 2MgO
2. Decomposition:
One compound splits: 2HgO → 2Hg + O2
3. Displacement:
• Metal: Zn + CuSO4 → ZnSO4 + Cu
• Non-metal: Cl2 + 2NaBr → 2NaCl + Br2
4. Disproportionation:
Same element is oxidized and reduced: 3Cl2 + 6NaOH → 5NaCl + NaClO3 + 3H2O
6. Redox Couple
Definition
A pair of substances where one is the oxidized form and the other is the reduced form (e.g., Fe3+/Fe2+).
Why It Matters
Redox couples determine the direction of electron flow in electrochemical cells and help calculate cell potential.
Detailed Explanation
In the Zn-Cu cell:
• Zn/Zn2+ (oxidation couple)
• Cu2+/Cu (reduction couple)
Electrons flow from Zn (higher tendency to oxidize) to Cu²⁺.
7. Electrochemical Cell
Definition
A device that converts chemical energy into electrical energy using spontaneous redox reactions.
Why We Use It
To generate electricity (e.g., batteries) or study reaction thermodynamics (e.g., corrosion prevention).
Components & Working
Anode (Oxidation): Zn → Zn2+ + 2e-
Cathode (Reduction): Cu2+ + 2e- → Cu
Salt Bridge: Maintains ionic balance using KNO3.
8. Electrode Potential
Definition
The voltage difference between an electrode and its electrolyte when no current flows.
Why Measure It?
To predict if a reaction is spontaneous (E°cell > 0) and compare oxidizing/reducing strengths.
Detailed Concept
Standard Electrode Potential (E°):
Measured under standard conditions (25°C, 1M, 1atm).
Formula: E°cell = E°cathode – E°anode
Example: For Zn²⁺/Zn (-0.76V) and Cu²⁺/Cu (+0.34V), E°cell = 1.10V.
9. Standard Hydrogen Electrode (SHE)
Definition
A reference electrode where hydrogen gas is oxidized or reduced at a platinum surface.
Why It’s Important
SHE acts as a universal reference (E° = 0V) to measure other electrodes’ potentials.
Setup & Applications
Construction:
• Pt electrode dipped in 1M H+
• H2 gas at 1 atm
Use Case: To find E° of Zn: Connect Zn half-cell to SHE and measure voltage.
10. Electrochemical Series
Definition
A list of elements/elements’ ions arranged by their standard reduction potentials.
Why It’s Useful
To identify strong oxidizing agents (top) vs. reducing agents (bottom) and predict reaction feasibility.
Key Features
• Metals like Li/K (strong reducers) are at the top.
• Non-metals like F2/Cl2 (strong oxidizers) are at the bottom.
Application: Iron (Fe) corrodes because it is above H2 in the series.
11. Redox Titrations
Definition
A volumetric analysis method where a redox reaction determines the concentration of an analyte.
Why Use Them?
For precise quantification of substances like Fe²⁺ in water or vitamin C in food.
Types & Process
1. Permanganometry:
• Titrant: KMnO4 (self-indicator, pink → colorless)
• Reaction: 5Fe2+ + MnO4- + 8H+ → 5Fe3+ + Mn2+ + 4H2O
2. Iodometry:
• Uses starch indicator (blue-black endpoint)
• Example: Determining Cu²⁺ in solution.
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