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A topic from the subject of Electrolysis in Chemistry.

  • 1 year ago
  • 6 min read
Faraday's Laws of Electrolysis
Introduction

Electrolysis is the process of using an electrical current to drive a chemical reaction. Faraday's Laws of Electrolysis, discovered by Michael Faraday in the early 19th century, describe the quantitative relationship between the amount of electricity passed through an electrolytic cell and the amount of chemical change that occurs.

Basic Concepts

Electrolytic Cell: A device consisting of two electrodes (cathode and anode) immersed in an electrolyte solution, through which an electric current is passed.

Electrolyte: A substance that contains ions and allows electricity to flow through it.

Electrodes: Conductors through which electricity enters and leaves the electrolytic cell.

Cathode: Electrode where reduction occurs (negative electrode).

Anode: Electrode where oxidation occurs (positive electrode).

Equipment and Techniques

Power Supply: Provides a constant voltage or current to the cell.

Electrodes: Inert materials (e.g., graphite, platinum) that do not participate in the electrochemical reactions.

Voltmeter: Measures the voltage across the cell.

Ammeter: Measures the current passing through the cell.

Electroanalytical Balance: Used to accurately measure the mass of electrodes before and after electrolysis.

Types of Experiments

Quantitative Analysis: Determining the amount of a substance in a solution by measuring the mass of metal deposited during electrolysis (gravimetric analysis).

Electroplating: Coating an object with a metal by electrodeposition.

Electrosynthesis: Producing chemical compounds by electrolysis.

Data Analysis

Faraday's Laws of Electrolysis relate the amount of electricity (current) passed through an electrolytic cell to the amount of chemical change:

Faraday's First Law: The mass (m) of a substance deposited or dissolved during electrolysis is directly proportional to the amount of charge (Q) passed through the cell: m = ZQ/F

  • Z: Electrochemical equivalent (grams deposited per coulomb)
  • F: Faraday constant (96,485 coulombs per mole of electrons)

Faraday's Second Law: When the same amount of electricity is passed through different electrolytic cells, the masses of different substances deposited or dissolved are directly proportional to their respective electrochemical equivalents: m1/Z1 = m2/Z2 = m3/Z3

Applications

Electrorefining: Purifying metals by electrolysis.

Electroplating: Coating surfaces for protection or decoration.

Electrosynthesis: Producing chemicals such as hydrogen and chlorine.

Quantitative Chemical Analysis: Determining the concentration of ions in solution.

Conclusion

Faraday's Laws of Electrolysis are fundamental principles that govern the quantitative aspects of electrochemical reactions. They provide a basis for understanding and harnessing the power of electrolysis for various industrial and analytical applications.

Faraday's Laws of Electrolysis

Faraday's First Law of Electrolysis (Law of Mass): The mass of a substance deposited or liberated at an electrode during electrolysis is directly proportional to the amount of charge passing through the circuit. Mathematically, this can be represented as: m ∝ Q, where 'm' is the mass of the substance and 'Q' is the charge.

Faraday's Second Law of Electrolysis (Law of Equivalents): When the same amount of charge passes through different electrolytic solutions, the masses of the substances liberated at each electrode are directly proportional to their equivalent weights (or electrochemical equivalents). Mathematically, this is expressed as: m1/E1 = m2/E2 = ... = constant, where 'm' represents the mass and 'E' represents the equivalent weight.

Key Points:
  • Electrolysis: The process of using direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. This involves the decomposition of an electrolyte (either molten or in solution) into its constituent elements.
  • Electrolytes: Substances that conduct electricity when dissolved in water or molten. They contain freely moving ions which carry the electric current.
  • Electrolytic Cell: The apparatus used to carry out electrolysis. It consists of two electrodes (anode and cathode) immersed in an electrolyte, connected to a DC power source.
  • Cations: Positively charged ions that migrate towards the cathode (negative electrode) during electrolysis, undergoing reduction (gain of electrons).
  • Anions: Negatively charged ions that migrate towards the anode (positive electrode) during electrolysis, undergoing oxidation (loss of electrons).
  • Electrochemical Equivalent (E): The mass of a substance deposited or liberated by one coulomb of charge. It is related to the equivalent weight (EW) by the formula: E = EW/F, where F is Faraday's constant (approximately 96485 Coulombs/mol).
  • Faraday's Constant (F): The charge carried by one mole of electrons (approximately 96485 Coulombs/mol).
  • Applications of Faraday's Laws: These laws are crucial in electroplating, electrorefining of metals, and other electrochemical processes.

Faraday's Laws of Electrolysis

Faraday's Laws of Electrolysis describe the quantitative relationship between the amount of electricity passed through an electrolyte solution and the amount of substance deposited or liberated at the electrodes.

Faraday's First Law of Electrolysis:

The mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed through the electrolyte.

Mathematically: m ∝ Q, where:

  • m = mass of substance deposited/liberated
  • Q = quantity of electricity (in Coulombs) = I × t (where I = current in Amperes and t = time in seconds)

Therefore, m = ZQ = ZIt, where Z is the electrochemical equivalent (mass deposited per unit charge).

Faraday's Second Law of Electrolysis:

When the same quantity of electricity is passed through different electrolytes, the masses of the substances deposited or liberated are proportional to their equivalent weights (atomic mass/valency).

Experiment Example 1: Electrolysis of Copper(II) Sulfate Solution

Materials: Copper electrodes, copper(II) sulfate solution (CuSO4), DC power supply, ammeter, stopwatch, weighing balance.

Procedure:

  1. Weigh the cathode (negative electrode) accurately.
  2. Set up the electrolysis apparatus: connect the copper electrodes to the DC power supply via an ammeter.
  3. Immerse the electrodes in the copper(II) sulfate solution.
  4. Switch on the power supply and record the current (I) and the time (t) for which the electrolysis is carried out.
  5. After a specific time, switch off the power supply and carefully remove the cathode.
  6. Wash and dry the cathode thoroughly.
  7. Weigh the cathode again.
  8. Calculate the mass of copper deposited (m = final mass - initial mass).
  9. Calculate the quantity of electricity passed (Q = I × t).
  10. Calculate the electrochemical equivalent of copper (Z = m/Q).

Observations: Record the current, time, initial mass, and final mass of the cathode.

Results: Calculate the mass of copper deposited and the electrochemical equivalent.

Experiment Example 2: Electrolysis of Water (demonstrates the second law)

Materials: Dilute sulfuric acid solution (to increase conductivity), two platinum electrodes, DC power supply, gas collection tubes (inverted over the electrodes).

Procedure:

  1. Fill the gas collection tubes with the dilute sulfuric acid solution and invert them over the platinum electrodes.
  2. Connect the electrodes to a DC power supply.
  3. Switch on the power supply and observe the gas produced at each electrode.
  4. Measure the volume of gas collected at each electrode.

Observations: Note the volume of hydrogen and oxygen gases collected. The ratio should be approximately 2:1 reflecting their equivalent weights and the stoichiometry of the reaction: 2H2O → 2H2 + O2

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