Class 12 Chemistry Chapter – Electrochemistry is one of the most interesting and scoring topics in Physical Chemistry. It connects electricity with chemical reactions and explains how redox reactions help generate electric current or use it to carry out chemical transformations. The chapter is extremely important for board exams and entrance tests like JEE Main, NEET, and CUET.
Electrochemistry explains the working of electrochemical cells, electrolytic cells, batteries, and corrosion processes. It also introduces students to vital concepts like electrode potential, EMF of a cell, and conductance of Class 11 Chemistry Solutions. These concepts form the foundation for several technological and industrial applications.
This detailed guide covers the NCERT-based explanation of each section, supported by tables, equations, and examples. It helps students understand theoretical and numerical problems related to galvanic and electrolytic cells with absolute clarity.
Table of Contents
- Introduction to Electrochemistry
- Electrochemical and Electrolytic Cells
- Conductance of Electrolytic Solutions
- Nernst Equation and EMF of Cell
- Applications and Corrosion
- FAQs
Introduction to Electrochemistry
Basic Concepts and Redox Reactions
| Concept | Definition | Example |
|---|---|---|
| Oxidation | Loss of electrons by a substance | \(Zn ightarrow Zn^{2+} + 2e^-\) |
| Reduction | Gain of electrons by a substance | \(Cu^{2+} + 2e^- ightarrow Cu\) |
| Redox Reaction | Simultaneous oxidation and reduction | \(Zn + Cu^{2+} ightarrow Zn^{2+} + Cu\) |
Electrochemistry revolves around redox reactions that occur in electrochemical systems. Oxidation happens at the anode and reduction at the cathode, leading to a flow of electrons that generates an electric current. This principle is used to build cells and batteries that power our devices.
The oxidation and reduction processes are always coupled. For example, in the Daniel Cell, zinc acts as the anode (oxidized) while copper acts as the cathode (reduced). Understanding electron flow and electrode reactions forms the foundation For All further electrochemical studies.
Electrochemical and Electrolytic Cells
Types, Components, and Reactions
| Cell Type | Energy Conversion | Example |
|---|---|---|
| Electrochemical (Galvanic) Cell | Chemical energy → Electrical energy | Daniel Cell |
| Electrolytic Cell | Electrical energy → Chemical change | Electrolysis of NaCl |
| Concentration Cell | Same electrodes, different ion concentrations | Hydrogen concentration cell |
Electrochemical cells convert chemical energy into electrical energy using spontaneous redox reactions. In contrast, electrolytic cells require external power sources to drive non-spontaneous reactions. Both types consist of electrodes (anode and cathode), electrolytes, and a salt bridge.
In a galvanic cell like the Daniel Cell, the reaction \(Zn + Cu^{2+}
ightarrow Zn^{2+} + Cu\) occurs spontaneously. Electrons move from zinc (anode) to copper (cathode), generating potential difference. In electrolytic cells, electricity decomposes compounds—for example, electrolysis of molten sodium chloride yields sodium and chlorine gas.
Conductance of Electrolytic Solutions
Specific, Molar, and Equivalent Conductance
| Type | Symbol | Definition | Unit |
|---|---|---|---|
| Specific Conductance | \(\kappa\) | Conductance of 1 cm3 of solution | S cm-1 |
| Molar Conductance | \(Lambda_m\) | Conductance of all ions from 1 mole of electrolyte | S cm2 mol-1 |
| Equivalent Conductance | \(Lambda_{eq}\) | Conductance of ions from 1 gram equivalent | S cm2 eq-1 |
Conductance depends on ion concentration and mobility. With dilution, the distance between ions increases, reducing inter-ionic attraction, which increases conductance. Kohlrausch’s Law of Independent Migration of Ions helps in calculating limiting molar conductance of strong electrolytes.
Mathematically, molar conductance is expressed as \(Lambda_m = \kappa imes \frac{1000}{C}\), where \(\kappa\) is specific conductance and \(C\) is concentration. This relationship forms the basis for many numerical questions in board exams and JEE papers.
Nernst Equation and EMF of Cell
Cell Potential and Nernst Equation
| Parameter | Symbol | Expression |
|---|---|---|
| Cell EMF | \(E_{cell}\) | \(E_{cell} = E^0_{cell} – \frac{RT}{nF} \ln Q\) |
| Standard EMF | \(E^0_{cell}\) | Measured under standard conditions |
| Reaction Quotient | \(Q\) | Ratio of activities of products and reactants |
The Nernst Equation relates the EMF of a cell to temperature, concentration, and number of electrons transferred. It helps calculate the potential under non-standard conditions. For example, at 298 K, the simplified form is \(E_{cell} = E^0_{cell} – \frac{0.0591}{n} \log Q\).
This relationship allows prediction of the direction of redox reactions and equilibrium constants. Students should memorize the equation and understand how each variable affects the potential. Practicing related numericals ensures command over electrochemistry calculations.
Applications and Corrosion
Practical Uses of Electrochemistry
| Application | Description | Example |
|---|---|---|
| Batteries | Convert chemical energy to electrical energy | Lead-acid battery, Li-ion battery |
| Electroplating | Coating a metal with another for protection | Silver plating of copper |
| Corrosion Control | Prevention of rusting using sacrificial anode | Galvanization of iron |
| Electrolysis | Decomposition using electrical energy | Electrolysis of water |
Electrochemistry is used extensively in industries and everyday life. Rechargeable batteries rely on reversible redox reactions, while electroplating enhances the durability and aesthetics of metals. Corrosion prevention methods, like galvanization, are based on electrochemical principles.
In biological systems, electrochemical gradients control nerve impulses and metabolic reactions. Understanding electrochemistry thus bridges the gap between physics, chemistry, and biology, making it one of the most interdisciplinary topics in science.