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
In real-life conditions, concentrations of ions are often not standard (1 mol/dm³), so the actual electrode potential differs from the standard value. The Nernst Equation allows us to calculate the electrode potential under non-standard conditions.

🔹 1. Nernst Equation (Simplified Form)

For a redox half-reaction:

perl
aA + ne⁻ ⇌ bB

The Nernst equation is:

ini
E = E° – (0.0591 / n) × log([Products] / [Reactants])

Where:

  • E = actual electrode potential

  • E° = standard electrode potential

  • n = number of electrons transferred

  • log = base 10 logarithm

  • Concentrations are in mol/dm³

🔹 2. Full Cell Version of Nernst Equation

For a full cell:

ini
Ecell = E°cell – (0.0591 / n) × log([Reduced form] / [Oxidized form])

🧪 Example 1: Effect of concentration on Zn/Cu cell

Given:

  • Cu²⁺ + 2e⁻ → Cu (E° = +0.34 V)

  • Zn²⁺ + 2e⁻ → Zn (E° = –0.76 V)

  • [Zn²⁺] = 1.0 M, [Cu²⁺] = 0.01 M

Step 1: Calculate E°cell

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

Step 2: Apply Nernst equation

ini
Ecell = 1.10 – (0.0591 / 2) × log([Zn²⁺] / [Cu²⁺])
Ecell = 1.10 – (0.02955) × log(1 / 0.01)
Ecell = 1.10 – 0.02955 × 2 = 1.10 – 0.0591 = 1.0409 V

✅ Answer: Ecell = 1.04 V

🔹 3. Observations from the Nernst Equation

  • Increasing [oxidized species] (e.g. Cu²⁺) increases cell potential

  • Increasing [reduced species] (e.g. Zn²⁺) decreases cell potential

  • When concentrations are equal, E = E°

🔹 4. Conditions that Affect Ecell

Factor Effect on Ecell
[Ion] concentration Shifts potential up or down
Temperature Affects reaction rate and equilibrium (not included in simplified Nernst)
Electrode surface Affects kinetics, not potential directly

🧠 NECTA Tips

  • Always state n (number of electrons in half-equation)

  • Use 2 decimal places in final answers

  • Write units clearly (V or mV)

  • Use log values accurately (log 10 = 1, log 0.01 = –2)

✅ Summary

  • The Nernst Equation corrects standard E° values for real concentrations

  • E decreases when oxidized species is diluted

  • E increases when reduced species is diluted

  • Useful in practical cell design, battery performance, and NECTA theory questions

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