Chapter 7: Redox Reactions

7.1 Classical Concept of Oxidation and Reduction

Example:

\[ \text{CuO}(s) + \text{H}_2(g) \rightarrow \text{Cu}(s) + \text{H}_2\text{O}(l) \]

• CuO is reduced (loss of oxygen)
• H₂ is oxidised (gain of oxygen)

7.2 Redox Reactions in Terms of Electron Transfer

Remember: OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons)

7.3 Oxidation Number (Oxidation State)

The oxidation number is the charge an atom would have if all bonds were ionic.

Rules for Assigning Oxidation Numbers

1. Free elements: Oxidation number = 0 (e.g., O₂, Na, Fe)
2. Monoatomic ions: Oxidation number = charge on ion (e.g., Na⁺ = +1, Cl⁻ = −1)
3. Hydrogen: Usually +1, except in metal hydrides (NaH, CaH₂) where it is −1
4. Oxygen: Usually −2, except in:
   • Peroxides (H₂O₂, Na₂O₂): −1
   • Superoxides (KO₂): −1/2
   • OF₂: +2
5. Fluorine: Always −1
6. Sum of oxidation numbers in a neutral molecule = 0; in an ion = charge on ion

Examples:

In \(\text{KMnO}_4\): Let Mn = \(x\)

\[ (+1) + x + 4(-2) = 0 \implies x = +7 \]

In \(\text{Cr}_2\text{O}_7^{2-}\): Let Cr = \(x\)

\[ 2x + 7(-2) = -2 \implies x = +6 \]

Redox Reaction Identification Using Oxidation Numbers

• Oxidation: Increase in oxidation number
• Reduction: Decrease in oxidation number
• Disproportionation: Same element is simultaneously oxidised and reduced

Example of disproportionation:

\[ 2\text{H}_2\text{O}_2 \rightarrow 2\text{H}_2\text{O} + \text{O}_2 \]

Oxygen in H₂O₂ is −1. In H₂O it becomes −2 (reduced), in O₂ it becomes 0 (oxidised).

7.4 Balancing Redox Reactions

Method 1: Oxidation Number Method

1. Write the skeleton equation
2. Assign oxidation numbers to all atoms
3. Identify atoms whose oxidation number changes
4. Equalize the increase and decrease in oxidation number
5. Balance the remaining atoms (O using H₂O, H using H⁺ in acidic medium or OH⁻ in basic)
6. Balance charge using electrons

Method 2: Half-Reaction (Ion-Electron) Method

Steps (in acidic medium):

1. Separate into oxidation and reduction half-reactions
2. Balance atoms other than O and H
3. Balance O by adding H₂O
4. Balance H by adding H⁺
5. Balance charge by adding electrons
6. Equalize electrons in both half-reactions
7. Add the half-reactions

Steps (in basic medium): Follow steps 1–7 for acidic medium, then add OH⁻ to both sides to neutralize H⁺.

Worked Example: Balance \(\text{Fe}^{2+} + \text{MnO}_4^- \rightarrow \text{Fe}^{3+} + \text{Mn}^{2+}\) in acidic medium.

Oxidation half-reaction:

\[ \text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + e^- \]

Reduction half-reaction:

\[ \text{MnO}_4^- + 8\text{H}^+ + 5e^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} \]

Multiply oxidation by 5:

\[ 5\text{Fe}^{2+} \rightarrow 5\text{Fe}^{3+} + 5e^- \]

Balanced equation:

\[ \text{MnO}_4^- + 5\text{Fe}^{2+} + 8\text{H}^+ \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O} \]

7.5 Types of Redox Reactions

7.6 Applications of Redox Reactions

1. Electrochemical cells (batteries, fuel cells)
2. Corrosion of metals (rusting of iron)
3. Extraction of metals from ores
4. Biological processes (photosynthesis, respiration)
5. Quantitative analysis (titrations with KMnO₄, K₂Cr₂O₇)

Practice Questions

Multiple Choice Questions (MCQs)

1. The oxidation number of Cr in \(\text{Cr}_2\text{O}_7^{2-}\) is:

 (a) +3

 (b) +4

 (c) +6

 (d) +7

2. In which of the following, hydrogen has an oxidation number of −1?

 (a) H₂O

 (b) HF

 (c) NaH

 (d) H₂O₂

3. Which of the following is a disproportionation reaction?

 (a) \(2\text{KClO}_3 \rightarrow 2\text{KCl} + 3\text{O}_2\)

 (b) \(\text{Cl}_2 + 2\text{NaOH} \rightarrow \text{NaCl} + \text{NaOCl} + \text{H}_2\text{O}\)

 (c) \(\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2\)

 (d) \(\text{Fe}_2\text{O}_3 + 3\text{CO} \rightarrow 2\text{Fe} + 3\text{CO}_2\)

4. The number of electrons involved in the reduction of \(\text{MnO}_4^-\) to \(\text{Mn}^{2+}\) is:

 (a) 1

 (b) 3

 (c) 5

 (d) 7

5. The oxidation state of S in \(\text{Na}_2\text{S}_2\text{O}_3\) is:

 (a) +2

 (b) −2

 (c) +4

 (d) +6

Short Answer Questions (2–3 Marks)

6. Assign oxidation numbers to all atoms in: (i) H₂SO₄ (ii) KMnO₄ (iii) Na₂Cr₂O₇

7. Balance the following reaction in acidic medium using the ion-electron method:

\[ \text{Cr}_2\text{O}_7^{2-} + \text{I}^- \rightarrow \text{Cr}^{3+} + \text{I}_2 \]

8. Identify the oxidising and reducing agents in: \(2\text{FeCl}_3 + \text{H}_2\text{S} \rightarrow 2\text{FeCl}_2 + \text{S} + 2\text{HCl}\)

9. What is a disproportionation reaction? Give an example.

10. The oxidation number of phosphorus in \(\text{Ba}(\text{H}_2\text{PO}_4)_2\) is?

Long Answer Questions (5 Marks)

11. (a) Explain the concept of oxidation number with examples.

 (b) Balance the following reaction in basic medium:

\[ \text{MnO}_4^- + \text{C}_2\text{O}_4^{2-} \rightarrow \text{MnO}_2 + \text{CO}_3^{2-} \]

12. (a) Classify the following reactions as combination, decomposition, displacement, or disproportionation:

 (i) \(2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}\)

 (ii) \(\text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2\)

 (iii) \(2\text{Cu}_2\text{O} \rightarrow 4\text{Cu} + \text{O}_2\) (at high temperature)

 (b) Give two applications of redox reactions in daily life.

Assertion-Reason Questions

13. Assertion (A): Fluorine always shows an oxidation state of −1.

Reason (R): Fluorine is the most electronegative element.

14. Assertion (A): In H₂O₂, the oxidation state of oxygen is −1.

Reason (R): H₂O₂ has an O–O bond.

Answer Key