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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 by which an electric current is passed through a substance, causing chemical reactions to occur. Faraday's Laws of Electrolysis describe the relationship between the amount of electric current passed through a substance and the amount of chemical change that occurs.

Basic Concepts

  • Electric current: The flow of electric charge through a conductor.
  • Electrolyte: A substance that contains ions, which are atoms or molecules that have lost or gained electrons.
  • Electrodes: Conductors that are connected to the power source and immersed in the electrolyte.
  • Anode: The electrode where oxidation occurs.
  • Cathode: The electrode where reduction occurs.
  • Faraday's constant: The amount of charge (approximately 96485 Coulombs) required to produce one mole of a substance.

Equipment and Techniques

  • Power source: A device that provides an electric current (e.g., battery, power supply).
  • Electrolytic cell: A container that holds the electrolyte and electrodes.
  • Voltmeter: A device that measures voltage.
  • Ammeter: A device that measures current.

Types of Experiments

  • Quantitative electrolysis: The amount of electric current passed through a substance is measured, and the amount of chemical change that occurs is determined (e.g., mass of metal deposited).
  • Qualitative electrolysis: The products of electrolysis are identified, and the reactions that occur are determined.

Data Analysis

  • Faraday's first law: The mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed through the electrolyte (Q = It, where Q is charge in Coulombs, I is current in Amperes, and t is time in seconds).
  • Faraday's second law: When the same quantity of electricity is passed through solutions of different electrolytes, the masses of the substances deposited or liberated are directly proportional to their equivalent weights (molar mass divided by the number of electrons transferred per ion).

Applications

  • Electroplating: Coating a metal object with a thin layer of another metal (e.g., chrome plating).
  • Electrorefining: Purifying a metal by removing impurities (e.g., copper refining).
  • Production of chemicals: Electrolysis is used to produce a variety of chemicals, such as hydrogen, chlorine, and sodium hydroxide (chlor-alkali process).

Conclusion
Faraday's Laws of Electrolysis are fundamental principles of electrochemistry with a wide range of applications. They provide a quantitative understanding of the relationship between the amount of electric current passed through a substance and the amount of chemical change that occurs.

Faraday's Laws of Electrolysis
Key Points
  • The mass of a substance deposited or liberated at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte.
  • The masses of different substances deposited or liberated by the same quantity of electricity are proportional to their equivalent weights (molar mass divided by the charge on the ion).
  • The electrochemical equivalent of an element is the mass of the element deposited or liberated by one coulomb of electricity. It is equal to its molar mass divided by its valence times the Faraday constant (F).
Faraday's Laws Explained
  • 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. Mathematically, this is expressed as: m = ZQ, where 'm' is the mass deposited, 'Z' is the electrochemical equivalent, and 'Q' is the quantity of electricity (in Coulombs).
  • Faraday's Second Law of Electrolysis: When the same quantity of electricity is passed through solutions of different electrolytes, the masses of the substances deposited or liberated are proportional to their equivalent weights. This implies that the same number of moles of electrons will deposit or liberate one equivalent weight of any substance.

Electrochemical Equivalent (Z): The electrochemical equivalent of a substance is the mass (in grams) deposited or liberated when one coulomb of electricity is passed through its solution. It's calculated as: Z = M / (n * F), where 'M' is the molar mass, 'n' is the valence (charge on the ion), and 'F' is the Faraday constant (approximately 96485 C/mol).

Applications of Faraday's Laws of Electrolysis
  • Electroplating: Coating a metal object with a thin layer of another metal for protection or aesthetics.
  • Electrorefining: Purifying metals by electrolysis.
  • Electrometallurgy: Extraction of metals from their ores using electrolysis.
  • Production of Hydrogen and Oxygen: Electrolysis of water to produce hydrogen and oxygen gas.
Faraday's Laws of Electrolysis Experiment

Materials:

  • 9-volt battery
  • Copper wire
  • Beaker
  • Copper sulfate solution (CuSO₄)
  • Voltmeter
  • Ammeter
  • Stopwatch
  • Balance (for measuring mass change of electrodes - optional but recommended for a complete demonstration of Faraday's Laws)
  • Sandpaper (for cleaning copper electrodes)

Procedure:

  1. Clean the copper wires using sandpaper to remove any oxide layer.
  2. Weigh the copper electrodes using the balance and record their masses. (Optional, but crucial for verifying Faraday's Laws)
  3. Pour the copper sulfate solution into the beaker.
  4. Immerse the copper electrodes into the copper sulfate solution, ensuring they do not touch each other.
  5. Connect the positive terminal of the battery to one copper electrode (anode) via a copper wire.
  6. Connect the negative terminal of the battery to the other copper electrode (cathode) via a copper wire.
  7. Connect the voltmeter in parallel across the electrodes to measure the voltage.
  8. Connect the ammeter in series with the circuit to measure the current.
  9. Start the stopwatch.
  10. Record the voltage, current, and time at regular intervals (e.g., every minute) for a predetermined time (e.g., 10-20 minutes). Note that longer electrolysis times will yield more easily measurable mass changes.
  11. After the predetermined time, stop the stopwatch and disconnect the circuit.
  12. Carefully remove the copper electrodes from the solution, rinse them with distilled water, and dry them thoroughly.
  13. Weigh the copper electrodes again and record their new masses. (Optional, but crucial for verifying Faraday's Laws)

Key Considerations:

  • Ensure the copper electrodes are clean before starting the experiment to obtain accurate results.
  • Use fresh copper sulfate solution to avoid contamination.
  • Monitor the voltage and current during the experiment to ensure they remain relatively constant. Significant variations may affect the accuracy.
  • Record the time accurately. The longer the experiment runs, the easier it will be to observe and measure the effects predicted by Faraday's Laws.

Data Analysis and Significance:

The mass change of the cathode (gain in mass) and the total charge passed (current × time) can be used to verify Faraday's Laws of Electrolysis. By calculating the electrochemical equivalent of copper, you can demonstrate the proportionality between the amount of substance deposited and the quantity of electricity passed. The data collected should show that the mass of copper deposited at the cathode is directly proportional to the charge passed through the electrolyte, confirming Faraday's first law. Additionally, the mass deposited for a given charge should be consistent with the electrochemical equivalent of copper, verifying Faraday's second law (when different metals are used).

This experiment demonstrates Faraday's Laws of Electrolysis, which state that:

  1. The mass of a substance deposited or liberated at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte (charge, Q = It).
  2. The mass of a substance deposited or liberated at an electrode is directly proportional to its equivalent weight.

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