Course Content
Module 1: Introduction to Electrochemistry
Overview of electrochemistry, its applications, and relevance in daily life and industrial processes.
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Definition and Importance of Electrochemistry
Module 2: Redox Reactions
Understanding oxidation, reduction, oxidation numbers, and balancing redox equations.
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Oxidation and Reduction Concepts
Oxidation Numbers
Balancing Redox Equations
📘 Module 3: Electrolysis
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Introduction to Electrolysis
Faraday’s Laws of Electrolysis
Industrial Applications of Electrolysis
Electrolysis Calculations
📘 Module 4: Redox Titrations
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Introduction and Indicators
KMnO₄ and Fe²⁺ Titration – Theory and Procedure
Oxalic Acid vs KMnO₄ Titration – Theory and Procedure
Iodometric Titrations – Iodine vs Thiosulfate (S₂O₃²⁻)
Electrode Potentials
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Standard Electrode Potentials (E° Values)
Electrochemical Cells and E° Cell Calculations
The Nernst Equation – Effect of Concentration on E°cell
📘 Module 6: Corrosion
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🧪 Lesson 1: Causes, Mechanism, and Prevention of Corrosion
📘 Module 7: Conductivity of Electrolyte Solutions
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🧪 Lesson 1: Electrical Conductivity and Factors Affecting It
📘 Module 8: Electrolytic Cells
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🧪 Lesson 1: Electrolysis Mechanism and Product Prediction
🧪 Lesson 2: Electrolysis of Water and Brine
📘 Module 9: Applications of Electrolysis in Industry
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🧪 Lesson 1: Electrolysis in Metal Extraction, Refining, and Electroplating
📘 Module 10: Redox Titrations and Calculations
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🧪 Lesson 1: Introduction, Procedure, and Redox Calculations
Advanced Chemistry: Electrochemistry

Introduction
An electrochemical cell is a system that converts chemical energy into electrical energy through redox reactions. The cell potential (E°cell) is calculated using standard electrode potentials and helps determine whether a redox reaction is feasible.

🔹 1. What is an Electrochemical Cell?

An electrochemical (galvanic) cell consists of two half-cells connected externally by a wire and internally by a salt bridge.

  • Anode: Where oxidation occurs (electrons released)

  • Cathode: Where reduction occurs (electrons received)

📌 Mnemonic: “AN OX, RED CAT”

  • ANode = OXidation

  • REDuction = CAThode

🔹 2. Cell Notation

Standard shorthand for an electrochemical cell:

 
Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

  • Left side: anode (oxidation half-cell)

  • Right side: cathode (reduction half-cell)

  • || = salt bridge

  • | = phase boundary

🔹 3. Cell Diagram Example

Zinc-Copper Cell:

  • Anode: Zn → Zn²⁺ + 2e⁻ (E° = –0.76 V)

  • Cathode: Cu²⁺ + 2e⁻ → Cu (E° = +0.34 V)

📌 Overall reaction:

scss
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

🔹 4. Calculating Standard Cell Potential (E°cell)

Formula:

mathematica
E°cell = E°cathode – E°anode

Example:

mathematica
E°cell = (+0.34 V) – (–0.76 V) = +1.10 V

✅ Positive E°cell → the reaction is spontaneous
❌ Negative E°cell → the reaction is non-spontaneous

🧪 Worked Example:

Q: Calculate E°cell for a cell composed of Fe²⁺/Fe and Cu²⁺/Cu.

Given:

  • Cu²⁺ + 2e⁻ → Cu; E° = +0.34 V

  • Fe²⁺ + 2e⁻ → Fe; E° = –0.44 V

Solution:

mathematica
E°cell = E°cathode – E°anode
= 0.34 – (–0.44) = 0.78 V

✅ Answer: E°cell = +0.78 V (reaction is feasible)

🔹 5. Salt Bridge Function

  • Completes the circuit

  • Maintains electrical neutrality by allowing ion flow

  • Often made of KNO₃ or NaCl in agar gel

🔹 6. Real-Life Applications

  • Dry cells (batteries)

  • Fuel cells (H₂/O₂ cell)

  • Electrochemical sensors (glucose, pH meters)

  • Corrosion prediction and protection

🧠 NECTA Tips

  • Label anode and cathode clearly

  • Always write oxidation and reduction half-equations

  • Show full cell notation

  • Round E° values to 2 decimal places unless instructed

✅ Summary

  • Electrochemical cells generate electricity through redox reactions

  • Anode = oxidation, Cathode = reduction

  • E°cell = E°cathode – E°anode

  • Positive E°cell → spontaneous reaction

  • Cell design and E° values are common in NECTA theory

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