What is Electrochemistry Class 12: Definition & Key Concepts

Understanding Electrochemistry: Definition and Scope

Electrochemistry is the branch of chemistry that deals with the interconversion of chemical energy and electrical energy through redox reactions. In Class 12 NCERT Chemistry, this chapter explains how electrons move in chemical reactions and how this movement can be harnessed to generate electric current.

Key points:

• Electrochemistry studies oxidation-reduction (redox) reactions involving electron transfer.
• It explains how batteries and cells work.
• It covers both spontaneous and non-spontaneous reactions.

This foundational knowledge helps students understand practical applications like batteries, corrosion, and electrolysis.

Galvanic Cells: How Chemical Energy Produces Electricity

A galvanic cell (or voltaic cell) converts chemical energy into electrical energy spontaneously. It consists of two different metals connected by a salt bridge and immersed in electrolyte solutions.

Components of a galvanic cell:

• Anode: Electrode where oxidation occurs (loss of electrons).
• Cathode: Electrode where reduction occurs (gain of electrons).
• Salt bridge: Maintains electrical neutrality by allowing ion flow.

Example: Daniell cell

• Zinc electrode (anode) in ZnSO₄ solution
• Copper electrode (cathode) in CuSO₄ solution

Electrons flow from zinc to copper through an external circuit, producing electric current.

Cell reaction:

$$\text{Zn (s)} \rightarrow \text{Zn}^{2+} + 2e^-$$ $$\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu (s)}$$

Electrolytic Cells: Using Electricity to Drive Chemical Reactions

Electrolytic cells use electrical energy to cause non-spontaneous chemical reactions. This process is called electrolysis.

Key features:

• Requires an external power source.
• Electrodes are connected to a battery or power supply.
• Anode is positive; cathode is negative.

Example: Electrolysis of molten sodium chloride

• At cathode: $$\text{Na}^+ + e^- \rightarrow \text{Na (s)}$$
• At anode: $$2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^-$$

Electrolysis is used in metal extraction, electroplating, and purification.

Electrode Potential and Standard Electrode Potential Explained

Electrode potential is the voltage developed at an electrode when it is in contact with its ions. It indicates the tendency of a species to gain or lose electrons.

• Measured in volts (V).
• Standard electrode potential ($E^\circ$) is measured under standard conditions (1 M concentration, 1 atm pressure, 25 °C).

The standard hydrogen electrode (SHE) is assigned $E^\circ = 0$ V and used as a reference.

Electrode potentials help predict the direction of redox reactions and calculate cell voltage:

$$E_{cell}^\circ = E_{cathode}^\circ - E_{anode}^\circ$$

Higher positive $E^\circ$ means stronger oxidizing agent.

Comparing Galvanic and Electrolytic Cells

Understanding the differences between galvanic and electrolytic cells is important for Class 12 students. Here's a comparison table:

This comparison helps clarify the roles and applications of each cell type.

Faraday’s Laws of Electrolysis and Their Applications

Faraday’s laws quantify the relationship between electric charge and the amount of substance produced at electrodes during electrolysis.

First Law: The mass of substance deposited or liberated at an electrode is proportional to the quantity of electricity passed.

$$m \propto Q$$

Where $m$ = mass, $Q$ = charge in coulombs.

Second Law: The masses of different substances liberated by the same quantity of electricity are proportional to their chemical equivalent weights.

Formula to calculate mass deposited:

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

Where:

• $M$ = molar mass
• $n$ = number of electrons transferred
• $F$ = Faraday constant (96500 C/mol)

Applications include electroplating, refining metals, and quantitative analysis.