Can add Faradays Laws - General Awareness for SSC CGL PDF Download

What is an Electrode?

An electrode is a designated point in an electrochemical setup where electric current enters or leaves the electrolyte or circuit. It acts as the cathode when current exits the electrode and as the anode when current enters it.

Electrodes are essential components of electrochemical cells and must possess strong electrical conductivity. Inert electrodes, which do not take part in chemical reactions, are also common. They can be made from materials like gold, platinum, carbon, graphite, or various metals, providing a surface for oxidation-reduction (redox) reactions.

Types of Electrodes

Electrodes fall into two main categories:

• Reactive electrodes: These participate directly in the cell's reactions and may dissolve into the electrolyte. Common examples include copper, silver, and zinc, often used in potentiometric studies.
• Inert electrodes: These remain chemically unaffected by the reactions. Typical examples are carbon and platinum electrodes.

Faraday's Laws of Electrolysis

First Law

• Discovered by Michael Faraday, this law states that the mass of a substance deposited or liberated at an electrode is directly proportional to the amount of electric charge passed through the circuit. Charge is measured in coulombs (or ampere-seconds).
• From Equation 1: m proportional to Q
• Removing the proportionality constant yields: m = Z * Q
• Here, m is the mass (in grams) of the substance deposited or released, Q is the charge (in coulombs), and Z is the electrochemical equivalent (ECE) of the substance—the mass deposited per unit charge.

Second Law

• This law indicates that when the same quantity of charge passes through different electrolytes, the masses of the substances deposited are proportional to their equivalent weights (or chemical equivalents).
• This is expressed as: W_​1 / E₁ = W₂ / E₂ = ...
• Where W represents the mass of the substance, and E is its equivalent weight.
• The equivalent weight is calculated as: E = Atomic weight / Valency

Illustrative Example

Consider three electrolytic cells in series, each undergoing a different reduction reaction, with y moles of electrons passing through:

1. In the first cell: Na⁺ + e^- -> Na → Deposits 23y grams of sodium.
2. In the second: Cu²⁺ + 2e^- -> Cu → Deposits 31.75y grams of copper (atomic weight 63.5 / valency 2).
3. In the third: Al³⁺ + 3e^- -> Al → Deposits 9y grams of aluminum (atomic weight 27 / valency 3).

One mole of electrons is required to reduce one equivalent of ions. The charge of one electron is 1.602 x 10^-19 C, so one mole of electrons (Avogadro's number: 6.023 x 10²³) carries: (6.023 x 10²³) x (1.602 x 10^-19) = 96,500 C

This value, 96,500 coulombs, is known as 1 Faraday (F). Passing 1 Faraday through a cell deposits 1 gram-equivalent of the substance, leading to the combined form of the laws: W = (Q / 96,500) x E

Or, for the electrochemical equivalent: Z = E / 96,500

Applications of Faraday's Laws

These principles underpin several practical processes, including:

• Extracting metal ions from aqueous solutions.
• Facilitating redox reactions via electrolysis.
• Producing heavy water through electrolysis.
• Manufacturing galvanic and fuel cells.
• Preventing metal corrosion through electroplating.