Comprehensive NEET Chemistry Notes

Chapter: Electrochemistry and Redox Reactions

1. Introduction to Redox Reactions

Redox reactions involve the transfer of electrons between two species. The term "redox" comes from the concepts of reduction and oxidation. In these reactions, one reactant is oxidized (loses electrons) while the other is reduced (gains electrons).

1.1 Oxidation and Reduction

• Oxidation: The process where an atom, ion, or molecule loses electrons. Example: $\text{Example:}\text{Mg}(s)+{\text{O}}_2(g)\to \text{MgO}(s)$
• Reduction: The process where an atom, ion, or molecule gains electrons. Example: $\text{Example:}\text{CuO}(s)+{\text{H}}_2(g)\to \text{Cu}(s)+{\text{H}}_2\text{O}(l)$

Mnemonic: OIL RIG - Oxidation Is Loss, Reduction Is Gain

NEET Tip:

Always identify the species getting oxidized and reduced by tracking electron transfer.

2. Types of Redox Reactions

2.1 Combination Reactions

• These reactions involve the combination of two or more elements or compounds to form a single product. $\text{A}+\text{B}\to \text{C}$
• Example: ${\text{N}}_2(g)+3{\text{H}}_2(g)\to 2{\text{NH}}_3(g)$

2.2 Decomposition Reactions

• These reactions involve the breakdown of a compound into two or more simpler substances. $\text{AB}\to \text{A}+\text{B}$
• Example: $2{\text{H}}_2{\text{O}}_2(l)\to 2{\text{H}}_2\text{O}(l)+{\text{O}}_2(g)$

Real-life Application:

Hydrogen peroxide decomposition is used in rocketry for propelling spacecraft.

3. Electrochemical Reactions

3.1 Galvanic Cells

• In a galvanic cell, a spontaneous redox reaction generates electrical energy.
• Example: Daniell Cell
• • Zinc rod in ZnSO4 solution (Anode)
  • Copper rod in CuSO4 solution (Cathode)
  • $\text{Zn}(s)\to {\text{Zn}}^{2+}(aq)+2e^{-}$
  • ${\text{Cu}}^{2+}(aq)+2e^{-}\to \text{Cu}(s)$

NEET Problem-Solving Strategy:

When solving electrochemical problems, always write the oxidation and reduction half-reactions separately before combining them.

3.2 Standard Electrode Potentials

• The standard electrode potential (E°) is a measure of the tendency of a chemical species to be reduced.
• Standard Hydrogen Electrode (SHE): ${\text{2H}}^{+}(aq)+2e^{-}\to {\text{H}}_2(g)E°=0.00\text{V}$

4. Applications of Redox Reactions

4.1 Electrolysis

• Electrolysis involves driving a non-spontaneous redox reaction using electrical energy.
• Example: Electrolysis of water ${\text{2H}}_2\text{O}(l)\to 2{\text{H}}_2(g)+{\text{O}}_2(g)$

Did You Know?

Electrolysis is used in the production of chlorine and caustic soda.

4.2 Batteries

• Batteries are devices that convert chemical energy into electrical energy through redox reactions.
• Example: Lead-acid battery $\text{Pb}+{\text{PbO}}_2+2{\text{H}}_2{\text{SO}}_4\to 2{\text{PbSO}}_4+2{\text{H}}_2\text{O}$

5. Balancing Redox Reactions

5.1 Oxidation Number Method

1. Assign oxidation numbers.
2. Identify the atoms that change oxidation states.
3. Equalize the increase and decrease in oxidation numbers.
4. Balance the atoms and charges.

5.2 Half-Reaction Method

1. Separate the reaction into oxidation and reduction half-reactions.
2. Balance each half-reaction for atoms and charge.
3. Combine the half-reactions, ensuring that electrons are balanced.

Example: ${\text{MnO}}_4^{-}+8{\text{H}}^{+}+5{\text{Fe}}^{2+}\to {\text{Mn}}^{2+}+5{\text{Fe}}^{3+}+4{\text{H}}_2\text{O}$

Quick Recap

• Oxidation involves the loss of electrons, reduction involves the gain.
• Types of redox reactions: combination, decomposition, displacement, and disproportionation.
• Galvanic cells convert chemical energy to electrical energy.
• Electrolysis and batteries are practical applications of redox reactions.
• Balancing redox reactions can be done using the oxidation number method or half-reaction method.

Concept Connection

Link to Biology:

Redox reactions play a crucial role in cellular respiration and photosynthesis.

Link to Physics:

Electrochemical cells relate to the concepts of current, voltage, and resistance in circuits.

Practice Questions

1. Identify the oxidizing and reducing agents in the reaction: $\text{2Na}+{\text{Cl}}_2\to 2\text{NaCl}$
2. • Solution: Na is oxidized (reducing agent), Cl2 is reduced (oxidizing agent).
2. Balance the following redox reaction using the half-reaction method: ${\text{MnO}}_4^{-}+{\text{C}}_2{\text{O}}_4^{2-}\to {\text{Mn}}^{2+}+{\text{CO}}_2$
3. Write the cell notation for a Daniell cell.
4. Calculate the standard cell potential for the reaction: $\text{Zn}(s)+{\text{CuSO}}_4(aq)\to {\text{ZnSO}}_4(aq)+\text{Cu}(s)$
5. Explain the process of electrolysis in the chlor-alkali industry.

Solution:

• Identify oxidation and reduction half-reactions.
• Use electrode potentials to calculate the cell potential.