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A topic from the subject of Decomposition in Chemistry.

  • 11 months ago
  • 6 min read

Electrolytic Decomposition

Introduction

Electrolysis is the process of causing a chemical reaction using electrical energy. It's often used to drive non-spontaneous reactions, frequently resulting in the decomposition of molecules in the original compounds. This is known as "Electrolytic Decomposition".

Basic Concepts
Electrolysis

Electrolysis uses an electric current to drive a non-spontaneous chemical reaction. This process takes place in an electrolytic cell, which is typically divided into two compartments (half-cells).

Electrolytic Decomposition

Electrolytic decomposition, or electrolysis, breaks down compound molecules into their constituent atoms using an electric current. This is the reverse of a chemical combination reaction.

Equipment and Techniques
  • Electrolytic Cell: This cell drives a non-spontaneous reaction when connected to a battery or power supply. It contains two electrodes: an anode (positive electrode) and a cathode (negative electrode). The electrodes are often inert materials like platinum or graphite to prevent them from reacting.
  • Power Source: A battery or power supply provides the electrical energy needed for the reaction.
  • Potentiometer (Voltmeter): This measures the potential difference (voltage) between the two electrodes.
  • Ammeter: This measures the current (amperes) flowing through the circuit.
Types of Experiments
Water Electrolysis

This is a common electrolytic decomposition experiment. Water (H₂O) is decomposed into hydrogen (H₂) and oxygen (O₂) gases. An electrolyte, such as sulfuric acid or sodium hydroxide, is typically added to increase conductivity.

Copper Sulphate Electrolysis

Copper sulphate (CuSO₄) electrolysis decomposes copper sulphate into copper (Cu) metal and sulphate ions (SO₄²⁻). The copper metal will be deposited on the cathode.

Data Analysis

Measuring the electric current and the amount of substance decomposed allows calculation of the number of moles of electrons transferred in the reaction, which is related to the electrochemical equivalent of the substance.

Applications
Electroplating

Electroplating coats a metal object with a thin layer of another metal. The object to be coated acts as the cathode in an electrolytic cell.

Electrorefining

Electrorefining purifies metals such as copper, gold, and silver. Impure metal is used as the anode, and pure metal is deposited at the cathode.

Other Applications:

Electrolytic decomposition is also used in the production of many chemicals, such as chlorine and sodium hydroxide (chlor-alkali process), and in the extraction of certain metals from their ores.

Conclusion

Electrolytic decomposition is crucial in various chemical and industrial processes. As technology advances, even more applications are expected to emerge.

Electrolytic Decomposition

Electrolytic Decomposition, also known as electrolysis, is a fundamental concept in chemistry involving the breakdown of a compound into its constituent elements using electric current. It is a redox process that involves the transfer of electrons from one species to another.

Key Points

  • Electric current as a catalyst: In Electrolytic Decomposition, an electric current acts as the catalyst to initiate the breakdown of a compound.
  • Electrodes and electrolyte: The process involves two electrodes (anode and cathode) and an electrolyte. The electrolyte is the compound/substance which is decomposed. The electrolyte must be molten or in solution to allow ion movement.
  • Redox reactions: Decomposition is essentially a series of oxidation-reduction reactions. At the anode, oxidation occurs (loss of electrons) while at the cathode, reduction occurs (gain of electrons).

Main Concepts

  1. Electrolyte: This is the compound that undergoes decomposition during electrolysis. It can be a molten salt or a solution of salts in a suitable solvent (e.g., water).
  2. Electric Current: An imposed direct current (DC) facilitates the movement of ions in the electrolyte, causing the compound to decompose. The voltage must be sufficient to overcome the decomposition potential of the electrolyte.
  3. Electrodes: Two electrodes – anode (positive) and cathode (negative) – are immersed in the electrolyte. The chosen compound decomposes at these electrodes. The choice of electrode material is important and can influence the products formed.
  4. Oxidation and Reduction: These two half-reactions occur simultaneously during electrolysis. Oxidation is the loss of electrons and happens at the anode, while reduction is the gain of electrons and occurs at the cathode. The overall reaction is the sum of the oxidation and reduction half-reactions.
  5. Faraday's Laws of Electrolysis: These laws quantitatively describe the relationship between the amount of substance decomposed and the quantity of electricity passed through the electrolyte. The first law states that the mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed. The second law states that the masses of different substances deposited or liberated by the same quantity of electricity are proportional to their equivalent weights.

Examples: Electrolysis is used in various applications, including the production of metals (e.g., aluminum from alumina), electroplating, and the purification of metals. The electrolysis of water is a common demonstration, producing hydrogen and oxygen gas.

Experiment: Electrolytic Decomposition of Water

Objective: The aim of this experiment is to observe and understand electrolysis, specifically the electrolytic decomposition of water into hydrogen and oxygen gases.

Materials Needed:
  • A 9V battery
  • Connecting wires
  • Two graphite pencils (or carbon electrodes)
  • Beaker (250ml)
  • Distilled water
  • Salt (1 teaspoon) – acts as an electrolyte to increase conductivity
  • Two test tubes
  • A source of power (9V battery is sufficient)
Procedure:
  1. Fill the beaker with distilled water.
  2. Add a teaspoon of salt to the water and stir until it dissolves completely. This increases the conductivity of the water, allowing for a more efficient electrolytic process.
  3. Sharpen both ends of the graphite pencils to expose more graphite surface area for better conductivity. Connect each pencil to a wire. Attach the other ends of the wires to the positive (+) and negative (-) terminals of the battery.
  4. Carefully place the sharpened ends of the pencils (electrodes) into the beaker, ensuring they are submerged in the salt water but not touching each other or the sides of the beaker. Maintain a reasonable distance between the electrodes.
  5. Observe the electrodes. You should see bubbles forming around both pencils. This indicates the start of the electrolysis process. The bubbles are the hydrogen and oxygen gases being produced.
  6. Invert the two test tubes and fill them completely with the salt water solution. Carefully place the inverted test tubes over each pencil to collect the gases produced. Make sure the open end of the test tube is below the water level to prevent air from entering.
  7. (Optional) To improve gas collection, you may add a small amount of detergent to the water. This will help trap the gas bubbles, making the volume of gas easier to measure.
  8. Allow the electrolysis to proceed for a sufficient amount of time (at least 15-20 minutes) to collect a noticeable volume of gases.
Observations:

After some time, you will observe that one test tube fills approximately twice as fast as the other. The test tube filling faster contains hydrogen gas (H2), which is collected at the cathode (negative electrode). The test tube filling slower contains oxygen gas (O2), collected at the anode (positive electrode). This 2:1 ratio demonstrates the stoichiometric relationship in the balanced equation for the electrolysis of water: 2H2O → 2H2 + O2

Significance:

Electrolysis is a fundamental electrochemical process that demonstrates the use of electrical energy to drive a non-spontaneous chemical reaction. This experiment visually demonstrates the decomposition of water into its constituent elements, hydrogen and oxygen, highlighting the stoichiometric ratio between them. It provides a practical illustration of concepts such as oxidation-reduction reactions (H2O is reduced to H2 at the cathode and oxidized to O2 at the anode), electrochemistry, and Faraday's laws of electrolysis.

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