JEE Chemistry Redox Reactions Complete Guide

Redox reactions, fundamental in JEE Chemistry, involve the transfer of electrons between chemical species. Mastery of redox concepts is crucial for balancing reactions, understanding electrochemical cells, and solving numerical problems efficiently. This comprehensive guide covers all essential topics, balancing methods, key formulas, and tips for exam success.

1. Understanding Redox Reactions

A redox (reduction-oxidation) reaction is one in which electrons are transferred from one reactant to another. Oxidation refers to the loss of electrons, while reduction is the gain of electrons.

\text{Oxidation: } \mathrm{A} \rightarrow \mathrm{A}^{n+} + n e^-

\text{Reduction: } \mathrm{B}^{m+} + m e^- \rightarrow \mathrm{B}

The species losing electrons is called the reducing agent and the one gaining electrons is the oxidizing agent.

2. Oxidation Number (Oxidation State)

Oxidation number helps identify which atoms undergo oxidation or reduction. It is a hypothetical charge assigned to an atom assuming ionic bonding.

2.1 Rules to Assign Oxidation Numbers

• Elemental form: 0 (e.g., \( O_2, H_2, N_2 \))
• Monoatomic ions: charge of the ion (e.g., \( Na^+ \) is +1)
• Oxygen: usually -2 (except in peroxides -1)
• Hydrogen: +1 with non-metals, -1 with metals
• Sum of oxidation numbers in a neutral molecule = 0
• Sum in a polyatomic ion = charge of the ion

3. Identifying Oxidation and Reduction

During a redox reaction, oxidation number of an element increases (oxidation) or decreases (reduction). Look for changes in oxidation states in reactants and products.

4. Balancing Redox Reactions

Balancing redox reactions ensures mass and charge conservation. Two main methods:

• Oxidation Number Method
• Ion-Electron Method (Half-Reaction Method)

4.1 Oxidation Number Method Steps

1. Assign oxidation numbers to all elements.
2. Identify oxidized and reduced species.
3. Calculate change in oxidation numbers and balance electron transfer.
4. Balance other atoms except hydrogen and oxygen.
5. Balance oxygen using \( \mathrm{H_2O} \) and hydrogen using \( \mathrm{H^+} \) (in acidic medium) or \( \mathrm{OH^-} \) (in basic medium).
6. Balance charge by adding electrons.
7. Verify mass and charge balance.

4.2 Ion-Electron Method Steps

1. Separate the reaction into oxidation and reduction half-reactions.
2. Balance all elements except \( \mathrm{H} \) and \( \mathrm{O} \).
3. Balance oxygen atoms by adding \( \mathrm{H_2O} \).
4. Balance hydrogen atoms by adding \( \mathrm{H^+} \) ions (acidic medium) or \( \mathrm{OH^-} \) ions (basic medium).
5. Balance charge by adding electrons.
6. Multiply half-reactions to equalize electrons and add them.
7. Cancel out common species and verify.

5. Electrochemical Cells

Electrochemical cells convert chemical energy to electrical energy via redox reactions. Two main types:

• Galvanic (Voltaic) Cells: spontaneous redox reactions producing electric current.
• Electrolytic Cells: use external current to drive non-spontaneous reactions.

5.1 Galvanic Cell Components

• Anode: site of oxidation (loss of electrons)
• Cathode: site of reduction (gain of electrons)
• Salt Bridge: maintains charge neutrality by ion flow

5.2 Standard Electrode Potential (\(E^\circ\))

The electrode potential measured under standard conditions (1 M, 1 atm, 25°C). It indicates tendency to gain electrons.

E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}

5.3 Nernst Equation

Calculates cell potential under non-standard conditions:

E = E^\circ - \frac{RT}{nF} \ln Q

At 25°C, simplified as:

E = E^\circ - \frac{0.0592}{n} \log Q

Where,

• \( E \): Cell potential
• \( E^\circ \): Standard cell potential
• \( R \): Gas constant
• \( T \): Temperature in Kelvin
• \( n \): Number of electrons transferred
• \( F \): Faraday constant
• \( Q \): Reaction quotient

6. Important Electrochemical Concepts

6.1 Electrochemical Series

A list of elements or ions arranged according to their standard electrode potentials. Helps predict redox behavior, reaction spontaneity, and oxidizing/reducing strength.

6.2 Faraday’s Laws of Electrolysis

1. Mass of substance deposited or liberated at an electrode is proportional to the quantity of electricity passed.
2. Mass of substances deposited or liberated by the same quantity of electricity are proportional to their equivalent weights.

m = \frac{Q \times M}{n \times F}

Where,

• m = mass of substance
• Q = total charge passed
• M = molar mass
• n = number of electrons
• F = Faraday constant (96485 C/mol)

7. Redox Titrations

Redox titrations involve reactions where the titrant or analyte undergoes oxidation or reduction. Common examples:

• Permanganometry (using \( \mathrm{KMnO_4} \))
• Iodometry (using iodine/iodide)
• Chromatometry (using dichromate ions)

7.1 Example: Permanganometry

In acidic medium, permanganate ion acts as a strong oxidizing agent:

\mathrm{MnO_4^-} + 8 \mathrm{H^+} + 5 e^- \rightarrow \mathrm{Mn^{2+}} + 4 \mathrm{H_2O}

8. Common Redox Reaction Types

• Metal displacement reactions
• Reaction of metals with acids
• Reactions involving halogens
• Reactions in acidic and basic medium
• Reaction involving free radicals

9. Practice Tips for JEE Redox Questions

• Always assign oxidation numbers first.
• Practice balancing redox reactions in both acidic and basic media.
• Memorize common electrode potentials and electrochemical series.
• Understand Faraday’s laws for electrolysis numericals.
• Attempt redox titration problems and electrochemical cell problems.
• Use the Nernst equation confidently for potential calculations.
• Focus on conceptual clarity over rote memorization.

10. Summary Table of Important Formulas

11. Conclusion

Redox reactions form an essential part of JEE Chemistry syllabus. With clear understanding of electron transfer, oxidation states, balancing techniques, and electrochemical principles, you can master this topic effectively. Consistent practice, solving numerical problems, and memorizing key concepts will ensure excellent exam performance.