Chapter 5: Electrochemistry Notes, Polytechnic 1st Semester Chemistry Notes BTER

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Welcome to Rajasthan Polytechnic BTER! This blog post is dedicated to Chapter 5 of the Chemistry syllabus for 1st-semester Polytechnic students. Here, you'll find detailed notes on Chapter 5: Electrochemistry, designed to provide you with all the essential information you need. For more quality content and notes, make sure to visit our website regularly.

This chapter covers the fundamental concepts of electrochemistry, including Faraday’s laws of electrolysis, corrosion, and various methods to prevent corrosion. Understanding these concepts is essential for students to grasp how chemical reactions occur in electrolytic cells and how metals degrade over time.

5.1 Faraday’s Laws of Electrolysis

Faraday's laws of electrolysis describe the relationship between the quantity of electricity passed through an electrolyte and the amount of substance deposited or dissolved at the electrodes. These laws are crucial for understanding how electroplating, battery charging, and other electrochemical processes work.

First Law of Electrolysis:
The amount of chemical substance deposited or dissolved at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte. Mathematically,
$\text{m} \propto \text{Q}$
where,

• $\text{m}$ = Mass of the substance deposited or dissolved (in grams)
• $\text{Q}$ = Total charge passed (in coulombs)

Second Law of Electrolysis:
When the same quantity of electricity is passed through different electrolytes, the mass of substances deposited or dissolved is proportional to their equivalent weights.
$\frac{m_1}{E_1} = \frac{m_2}{E_2} = \frac{m_3}{E_3}$
where,

• $m_1, m_2, m_3$ = Masses of different substances deposited or dissolved
• $E_1, E_2, E_3$ = Their respective equivalent weights

Simple Numerical Problem Example:
Calculate the mass of copper deposited when a current of 2 amperes is passed through a copper sulfate solution for 30 minutes. (Equivalent weight of copper = 31.75 g)

Solution:

1. Convert time into seconds: $30 \times 60 = 1800$ seconds
2. Calculate total charge: $Q = \text{I} \times \text{t} = 2 \times 1800 = 3600$ coulombs
3. Using the formula:
   $m = \frac{Q \times E}{96500} = \frac{3600 \times 31.75}{96500} \approx 1.185 \, g$

5.2 Introduction to Corrosion of Metals

Corrosion is the gradual destruction of metals due to chemical or electrochemical reactions with their environment. It leads to the deterioration of materials and structures, causing economic losses and safety hazards.

5.2.1 Definition

Corrosion is the process by which metals are slowly converted into their oxides, hydroxides, or other compounds by reacting with moisture, air, acids, or other chemicals present in the environment.

5.2.2 Types of Corrosion

1. Chemical Corrosion: This occurs due to direct chemical reactions between the metal surface and the environment, such as oxidation in the presence of dry gases like oxygen or chlorine.
2. Electrochemical Corrosion: This involves electrochemical reactions, where metal acts as an anode, losing electrons and forming ions, leading to degradation. It occurs in the presence of an electrolyte, such as water.

5.3 H₂ Liberation and O₂ Absorption Mechanism of Electrochemical Corrosion

Electrochemical corrosion typically occurs in two main mechanisms:

1. Hydrogen Evolution (H₂ Liberation):

   • Common in acidic environments.
   • Metal acts as the anode and dissolves, releasing electrons.
   • These electrons are used to reduce hydrogen ions (H⁺) to hydrogen gas (H₂) at the cathode. $\text{M} \rightarrow \text{M}^{2+} + 2e^{-}$
     $2H^{+} + 2e^{-} \rightarrow H_{2}$

2. Oxygen Absorption (O₂ Absorption):

   • Common in neutral or alkaline environments.
   • Oxygen is reduced at the cathode, forming hydroxide ions (OH⁻). $\text{M} \rightarrow \text{M}^{2+} + 2e^{-}$
     $O_{2} + 2H_{2}O + 4e^{-} \rightarrow 4OH^{-}$

5.4 Factors Affecting Rate of Corrosion

Several factors can influence the rate of corrosion:

1. Nature of Metal: Metals with higher reactivity corrode faster. For example, iron corrodes faster than copper.
2. Presence of Impurities: Impurities can create galvanic cells on the metal surface, accelerating corrosion.
3. Environment: Humidity, temperature, and presence of corrosive chemicals (like salt) can increase the rate of corrosion.
4. pH of the Environment: Corrosion is more severe in acidic conditions.
5. Oxygen Availability: More oxygen means faster oxidation and corrosion.

5.5 Internal Corrosion Preventive Measures

1. Purification: Removing impurities from metals during production can reduce corrosion.
2. Alloying: Adding other elements like chromium or nickel to metals (e.g., in stainless steel) can improve corrosion resistance.
3. Heat Treatment: Processes like annealing and quenching can enhance the microstructure of metals, making them more resistant to corrosion.

5.6 External Corrosion Preventive Measures

1. Metal Coatings:

   • Anodic Coatings: Applying a metal like zinc on iron. Zinc corrodes preferentially, protecting the iron (Galvanizing).
   • Cathodic Coatings: Applying a less reactive metal like tin on iron. If the coating is broken, iron will corrode.

2. Organic Inhibitors: These are chemical substances added to the environment that slow down the rate of corrosion. They form a protective film on the metal surface.

This chapter on electrochemistry covers essential concepts like Faraday's laws, corrosion, and its prevention. Understanding these principles is crucial for students in the Polytechnic 1st Semester Chemistry course. Keep revisiting Rajasthan Polytechnic for more detailed notes and explanations to ace your exams.

These Notes are according to Syllabus provided by BTER (Board of Technical Education Rajasthan, Jodhpur) for Polytechnic 1st Semester Students.

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Garima Kanwar