Redox Reactions: Complete Guide for Class 11 NCERT Chemistry

What Are Redox Reactions?

Redox Reactions are chemical reactions where oxidation and reduction occur simultaneously. Oxidation means loss of electrons, while reduction means gain of electrons. These reactions are fundamental in processes like combustion, corrosion, and cellular respiration.

In Class 11 NCERT Chemistry, understanding redox reactions helps explain how substances interact by exchanging electrons. The term "redox" is a combination of "reduction" and "oxidation". For example, in the reaction:

$$\text{Zn} + \text{CuSO}_4 \rightarrow \text{ZnSO}_4 + \text{Cu}$$

Zinc loses electrons (oxidised) and copper ions gain electrons (reduced). This electron transfer is the core of redox reactions.

Types of Redox Reactions Explained

Redox reactions are classified into four main types based on how substances react:

1. Combination Reactions: Two or more substances combine to form one product. At least one reactant is elemental.

• Example: $\text{C}(s) + \text{O}_2(g) \rightarrow \text{CO}_2(g)$

2. Decomposition Reactions: A compound breaks down into two or more substances, including at least one element.

• Example: $2\text{H}_2\text{O}(l) \rightarrow 2\text{H}_2(g) + \text{O}_2(g)$

3. Displacement Reactions: An element displaces another from its compound.

• Metal displacement: $\text{CuSO}_4(aq) + \text{Zn}(s) \rightarrow \text{Cu}(s) + \text{ZnSO}_4(aq)$
• Non-metal displacement: $\text{Cl}_2(g) + 2\text{KBr}(aq) \rightarrow 2\text{KCl}(aq) + \text{Br}_2(l)$

4. Disproportionation Reactions: A single element undergoes both oxidation and reduction.

• Example: $2\text{H}_2\text{O}_2(aq) \rightarrow 2\text{H}_2\text{O}(l) + \text{O}_2(g)$

Understanding these types helps in identifying and balancing redox reactions.

Combination and Decomposition Reactions in Detail

Combination Reactions

In these reactions, two or more reactants form a single product. At least one reactant is in elemental form, which undergoes oxidation or reduction.

Examples:

• $3\text{Mg}(s) + \text{N}_2(g) \rightarrow \text{Mg}_3\text{N}_2(s)$
• $\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(l)$

These reactions are common in synthesis and combustion.

Decomposition Reactions

Here, a compound breaks down into simpler substances, at least one being elemental.

Examples:

• $2\text{NaH}(s) \rightarrow 2\text{Na}(s) + \text{H}_2(g)$
• $2\text{KClO}_3(s) \rightarrow 2\text{KCl}(s) + 3\text{O}_2(g)$

Note: Not all decomposition reactions are redox. For instance, the decomposition of calcium carbonate ($\text{CaCO}_3$) is not a redox reaction because oxidation states remain unchanged.

These reactions illustrate how redox processes can both build and break chemical bonds.

Displacement Reactions: Metal and Non-Metal Examples

Displacement reactions occur when an element displaces another from its compound. These are important in metal extraction and halogen chemistry.

Metal Displacement:

• A more reactive metal displaces a less reactive metal from its salt solution.
• Example:

$$\text{CuSO}_4(aq) + \text{Zn}(s) \rightarrow \text{Cu}(s) + \text{ZnSO}_4(aq)$$

• Other examples include:
• $\text{V}_2\text{O}_5(s) + 5\text{Ca}(s) \rightarrow 2\text{V}(s) + 5\text{CaO}(s)$
• $\text{Cr}_2\text{O}_3(s) + 2\text{Al}(s) \rightarrow \text{Al}_2\text{O}_3(s) + 2\text{Cr}(s)$

Non-Metal Displacement:

• Halogens displace other halide ions due to higher reactivity.
• Examples:
• $\text{Cl}_2(g) + 2\text{KBr}(aq) \rightarrow 2\text{KCl}(aq) + \text{Br}_2(l)$
• $\text{Cl}_2(g) + 2\text{KI}(aq) \rightarrow 2\text{KCl}(aq) + \text{I}_2(s)$

Hydrogen Displacement:

• Metals can displace hydrogen from water or acids.
• Examples:
• $2\text{Na}(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{NaOH}(aq) + \text{H}_2(g)$
• $\text{Zn}(s) + 2\text{HCl}(aq) \rightarrow \text{ZnCl}_2(aq) + \text{H}_2(g)$

These reactions demonstrate electron transfer and reactivity trends in the periodic table.

Understanding Disproportionation Reactions

Disproportionation reactions are unique redox processes where a single element is simultaneously oxidised and reduced, producing two different products.

Key Examples:

• Hydrogen peroxide decomposition:

$$2\text{H}_2\text{O}_2(aq) \rightarrow 2\text{H}_2\text{O}(l) + \text{O}_2(g)$$ Here, oxygen in $\text{H}_2\text{O}_2$ is both oxidised to $\text{O}_2$ and reduced to $\text{H}_2\text{O}$.

• Chlorine in alkaline medium:

$$\text{Cl}_2(g) + 2\text{OH}^-(aq) \rightarrow \text{ClO}^-(aq) + \text{Cl}^-(aq) + \text{H}_2\text{O}(l)$$

• Phosphorus and sulfur disproportionation also occur in basic solutions forming multiple oxidation state products.

Important Note: Fluorine does not undergo disproportionation due to its high electronegativity and inability to form positive oxidation states.

Disproportionation reactions highlight the versatility of elements in redox chemistry.

Balancing Redox Reactions: The Basics

Balancing redox reactions is crucial for solving Class 11 NCERT problems. It ensures mass and charge conservation.

Steps to Balance Redox Reactions: 1. Write the unbalanced equation. 2. Separate into oxidation and reduction half-reactions. 3. Balance atoms other than O and H. 4. Balance oxygen atoms by adding $\text{H}_2\text{O}$. 5. Balance hydrogen atoms by adding $\text{H}^+$ (in acidic medium) or $\text{OH}^-$ (in basic medium). 6. Balance charges by adding electrons ($e^-$). 7. Equalize electrons in both half-reactions. 8. Add the half-reactions and simplify.

Example: Balance the reaction between permanganate ion and iron(II) ion in acidic medium:

Unbalanced: $$\text{MnO}_4^- + \text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + \text{Fe}^{3+}$$

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 half by 5 and add: $$5\text{Fe}^{2+} + \text{MnO}_4^- + 8\text{H}^+ \rightarrow 5\text{Fe}^{3+} + \text{Mn}^{2+} + 4\text{H}_2\text{O}$$

Balancing redox reactions is essential for chemical equations and electrochemistry topics.

Comparison Table: Types of Redox Reactions