Unit 7: Electrochemistry

What is Electrochemistry?

Electrochemistry is the branch of chemistry that studies the relationship between electrical energy and chemical reactions. This unit explores the principles of electrochemical cells, including both galvanic (voltaic) and electrolytic cells. It covers topics such as redox reactions, electrochemical potential, and the application of electrochemistry in various technologies and industries.

Key Topics in Electrochemistry:

• Electrochemical Cells: Understanding the components and operation of galvanic (voltaic) cells and electrolytic cells.
• Redox Reactions: Examining oxidation-reduction reactions and their role in electrochemical processes.
• Electrode Potentials: Studying standard electrode potentials and how they relate to the spontaneity of reactions.
• Nernst Equation: Applying the Nernst equation to calculate cell potentials under non-standard conditions.
• Applications of Electrochemistry: Exploring real-world applications, such as batteries, electroplating, and corrosion prevention.

Benefits of Studying Electrochemistry:

• Understanding Energy Conversion: Provides insights into how chemical energy is converted into electrical energy and vice versa.
• Practical Applications: Enhances knowledge relevant to technologies like batteries, fuel cells, and electroplating.
• Scientific Insight: Builds a foundation for advanced studies in physical chemistry, materials science, and engineering.

This unit is essential for students to grasp the fundamental concepts of electrochemistry, its applications, and its impact on modern technology. Mastery of these concepts is crucial for understanding various electrochemical processes and their practical uses.

1. Spontaneous chemical reactions take place in:

a. Electrolytic cell
b. Galvanic cell
c. Nelson’s cell
d. Downs cell

2. Formation of water from hydrogen and oxygen is:

a. Redox reaction
b. Acid-base reaction
c. Neutralization
d. Decomposition

3. Which one of the following is not an electrolytic cell?

a. Downs cell
b. Galvanic cell
c. Nelson’s cell
d. Both a and c

4. The oxidation number of chromium in K₂Cr₂O₇ is:

a. +2
b. +6
c. +7
d. +14

5. Which one of the following is not an electrolyte?

a. Sugar solution
b. Sulphuric acid solution
c. Lime solution
d. Sodium chloride solution

6. The most common example of corrosion is:

a. Chemical decay
b. Rusting of iron
c. Rusting of aluminium
d. Rusting of tin

7. Nelson’s cell is used to prepare caustic soda along with gases. Which of the
following gas is produced at cathode:

a. Cl₂
b. H₂
c. O₃
d. O₂

8. During the formation of water from hydrogen and oxygen, which of the
following does not occur:

a. Hydrogen has oxidized
b. Oxygen has reduced
c. Oxygen gains electrons
d. Hydrogen behaves as oxidizing agent

9. The formula of rust is:

a. Fe₂O₃.nH₂O
b. Fe₂O₃
c. Fe(OH)₃.nH₂O
d. Fe(OH)₃

10. In the redox reaction between Zn and HC1, the oxidizing agent is:

a. Zn
b. H⁺
c. Cl^–
d. H₂