Back to Library

(AI-Powered Suggestions)

Related Topics

A topic from the subject of Electrolysis in Chemistry.

  • 11 months ago
  • 6 min read
Electrolytic Cells vs Galvanic Cells
Introduction

Electrolytic cells and galvanic cells are two types of electrochemical cells. Electrolytic cells use electrical energy to drive non-spontaneous chemical reactions, while galvanic (or voltaic) cells use spontaneous chemical reactions to generate electrical energy. Both types of cells are important in various applications, including battery technology, chemical synthesis, and metal refining.

Basic Concepts
  • Electrochemical Cells: Electrochemical cells are devices that convert chemical energy to electrical energy (galvanic cells) or electrical energy to chemical energy (electrolytic cells). These conversions involve the movement of electrons through an external circuit connected to the cell.
  • Electrodes: Electrodes are conductors where the electrochemical reactions occur. They are often made of metals (e.g., copper, iron, zinc, platinum, graphite).
  • Anodes and Cathodes: The anode is the electrode where oxidation (loss of electrons) occurs. The cathode is the electrode where reduction (gain of electrons) occurs.
  • Anode Half-Reaction: The anode half-reaction describes the oxidation process at the anode. It shows the loss of electrons.
  • Cathode Half-Reaction: The cathode half-reaction describes the reduction process at the cathode. It shows the gain of electrons.
  • Overall Cell Reaction: The overall cell reaction is the sum of the anode and cathode half-reactions. It represents the net chemical change.
  • Electromotive Force (EMF): EMF is the potential difference between the anode and cathode. In galvanic cells, EMF is positive; in electrolytic cells, EMF is negative (requiring an external power source).
Key Differences: Galvanic vs. Electrolytic Cells
Feature Galvanic Cell Electrolytic Cell
Reaction Type Spontaneous Non-spontaneous
Energy Conversion Chemical to Electrical Electrical to Chemical
EMF Positive Negative
Anode Oxidation Oxidation
Cathode Reduction Reduction
External Power Source Not required Required
Equipment and Techniques
  • Cell Chamber: Contains the electrolytes and electrodes.
  • Electrodes: Often made of inert metals (platinum, graphite) to avoid unwanted reactions.
  • Salt Bridge (Galvanic): Connects the half-cells, allowing ion flow to maintain electrical neutrality.
  • Power Supply (Electrolytic): Provides the electrical energy to drive the non-spontaneous reaction.
  • Multimeter: Measures EMF and current.
Types of Experiments
  • Electrolysis of Water: Decomposition of water into hydrogen and oxygen using electricity.
  • Electroplating: Coating a metal surface with another metal via electrodeposition.
  • Battery Construction: Building a simple battery to demonstrate galvanic cell principles.
  • Corrosion Study: Investigating factors affecting the corrosion of metals.
Data Analysis
  • Faraday's Law: Relates the amount of substance produced or consumed to the charge passed.
  • Nernst Equation: Calculates EMF under non-standard conditions.
  • Polarization Curves: Show the relationship between current and voltage.
Applications
  • Batteries: Galvanic cells storing and releasing electrical energy.
  • Electrolysis: Production of chemicals and refining of metals.
  • Corrosion Control: Cathodic protection and sacrificial anodes.
  • Fuel Cells: Generate electricity from fuel oxidation.
Conclusion

Electrolytic and galvanic cells are crucial in many electrochemical processes. Understanding their principles is essential in chemistry, energy, and industrial applications.

Electrolytic Cells vs. Galvanic Cells
Key Points
  • Electrolytic cells use electrical energy to drive a non-spontaneous chemical reaction.
  • Galvanic cells use a spontaneous chemical reaction to generate electrical energy.
  • Both types of cells consist of two electrodes (an anode and a cathode) immersed in an electrolyte solution.
  • In an electrolytic cell, an external power source is used to force electrons to flow from the anode to the cathode (against the natural flow).
  • In a galvanic cell, electrons flow spontaneously from the anode to the cathode, generating an electric current.
Main Concepts
Electrolytic Cells:
  • Consist of an anode and a cathode connected by a wire and an external power source.
  • The anode is the electrode where oxidation occurs (electrons are lost).
  • The cathode is the electrode where reduction occurs (electrons are gained).
  • An external power source is required to drive the non-spontaneous reaction.
  • The direction of electron flow is from the anode to the cathode (forced by the external power source).
  • Examples: Electrolysis of water, electroplating, charging a battery.
Galvanic Cells:
  • Consist of an anode and a cathode connected by a wire (often with a salt bridge).
  • The anode is the electrode where oxidation occurs (electrons are lost).
  • The cathode is the electrode where reduction occurs (electrons are gained).
  • The chemical reaction that occurs is spontaneous.
  • The direction of electron flow is from the anode to the cathode, generating an electric current.
  • Examples: Batteries (alkaline, zinc-carbon, lithium-ion), fuel cells.
Comparison:
Electrolytic Cell Galvanic Cell
Energy Source External power source Spontaneous chemical reaction
Direction of Electron Flow Anode to Cathode Anode to Cathode
Type of Reaction Non-spontaneous (requires energy input) Spontaneous (releases energy)
ΔG (Gibbs Free Energy Change) Positive (+ΔG) Negative (-ΔG)
Cell Potential (Ecell) Negative (-Ecell) Positive (+Ecell)
Examples Electrolysis of water, electroplating, charging a battery Batteries, fuel cells
Experiment: Electrolytic Cells vs. Galvanic Cells

Objective: To demonstrate and understand the differences between electrolytic and galvanic cells, and to observe the processes of electrolysis and electroplating.

Materials:
  • Two beakers
  • Copper sulfate solution (CuSO4)
  • Zinc sulfate solution (ZnSO4)
  • Copper electrodes (Cu)
  • Zinc electrodes (Zn)
  • Voltmeter
  • Ammeter
  • Direct Current (DC) Power Supply (Battery)
  • Connecting wires
  • Salt bridge (optional, for improved galvanic cell performance - e.g., a U-shaped tube filled with agar-agar gel containing a salt like potassium nitrate)
Procedure: Part 1: Electrolytic Cell
  1. Fill two beakers with copper sulfate solution and zinc sulfate solution, respectively. Place a copper electrode in the CuSO4 solution and a zinc electrode in the ZnSO4 solution. These should be separate beakers.
  2. Connect the copper electrode to the positive terminal (anode) of the DC power supply and the zinc electrode to the negative terminal (cathode).
  3. Connect a voltmeter in parallel with the electrodes to measure the voltage across the cell.
  4. Connect an ammeter in series with the circuit to measure the current.
  5. Turn on the DC power supply and observe the readings on the voltmeter and ammeter. Note any changes in the electrodes or solutions.
  6. Record your observations, including any changes observed at the electrodes (e.g., gas evolution, color changes, deposition of metal).
  7. (Optional) If using separate beakers, add a salt bridge to complete the circuit and facilitate ion flow between the two half-cells.
Part 2: Galvanic Cell
  1. Prepare a galvanic cell. One approach is to place a copper electrode in a beaker containing copper sulfate solution and a zinc electrode in a separate beaker containing zinc sulfate solution. Connect the electrodes with a wire.
  2. Connect a voltmeter in parallel with the electrodes to measure the cell potential (voltage).
  3. Connect an ammeter in series with the circuit to measure the current flow.
  4. Observe the readings on the voltmeter and ammeter. Note the polarity of the voltage reading (positive or negative).
  5. Record your observations, including any changes observed at the electrodes. Note the direction of electron flow.
  6. (Necessary) Connect the two solutions with a salt bridge to complete the circuit.
Results:
  • Electrolytic Cell: The voltmeter will show a positive voltage reading (indicating that the power supply is forcing the non-spontaneous reaction), and the ammeter will show a positive current reading. You should observe electrolysis; for example, copper will deposit on the cathode in the copper sulfate cell.
  • Galvanic Cell: The voltmeter will show a positive voltage reading (representing the cell potential, the electromotive force, EMF), and the ammeter will show a positive current reading. The zinc electrode will be the anode (oxidation) and the copper electrode will be the cathode (reduction). The reaction is spontaneous.
Conclusions:
  • Electrolytic cells use electrical energy to drive a non-spontaneous chemical reaction, while galvanic cells use a spontaneous chemical reaction to generate electrical energy.
  • In an electrolytic cell, the anode is positive and oxidation occurs; the cathode is negative and reduction occurs.
  • In a galvanic cell, the anode is negative and oxidation occurs; the cathode is positive and reduction occurs. (Note the reversal of electrode polarities compared to the electrolytic cell).
  • Electrolytic cells are used for processes such as electroplating and the production of certain chemicals, while galvanic cells are used for powering devices such as batteries and fuel cells.

Share on: