Electrochemical Cell

What is Electrochemical Cell

An electrochemical cell is a device that converts chemical energy into electrical energy through electrochemical reactions. These cells consist of two electrodes (anode and cathode) immersed in an electrolyte solution, which allows the flow of ions between the electrodes, and an external circuit that allows the flow of electrons. Electrochemical cells are commonly used in batteries and fuel cells.

Electrochemical Cell

Components of electrochemical cell

An electrochemical cell consists of several essential components that work together to facilitate the conversion of chemical energy into electrical energy (or vice versa). The main components of an electrochemical cell are as follows:

  1. Electrodes: Electrodes are conductive materials, typically made of metals or materials with conductive coatings, that serve as the sites where the redox (reduction-oxidation) reactions take place. There are two types of electrodes in an electrochemical cell:
  • Anode: The electrode where oxidation occurs, leading to the loss of electrons.
  • Cathode: The electrode where reduction occurs, leading to the gain of electrons.
  1. Electrolyte: The electrolyte is a substance, often a solution or a molten salt, that allows the flow of ions between the anode and cathode while preventing the direct flow of electrons. It serves as the medium for ion transport during the redox reactions.
  2. External Circuit: An external circuit provides a conductive pathway that allows electrons to flow from the anode to the cathode. This flow of electrons creates an electric current that can be used to do work, such as powering electrical devices.
  3. Salt Bridge (or Separator): In some electrochemical cells, a salt bridge or a porous separator is used to maintain electrical neutrality within the cell. It allows the flow of ions between the anode and cathode while preventing the mixing of the electrolyte solutions in each half-cell.

These four components—electrodes, electrolyte, external circuit, and salt bridge/separator—work together to enable the redox reactions to proceed, facilitating the conversion of chemical energy into electrical energy in galvanic cells (also known as voltaic cells) and the reverse process in electrolytic cells. The type and arrangement of these components vary depending on the specific type of electrochemical cell and its intended application.

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How electrochemical cell work

Electrochemical cells work by harnessing redox (reduction-oxidation) reactions to convert chemical energy into electrical energy (galvanic cells) or vice versa (electrolytic cells). Here’s a step-by-step explanation of how electrochemical cells function:

1. Redox Reactions:

  • Galvanic Cells (Batteries): In galvanic cells, also known as voltaic cells, spontaneous redox reactions occur. These reactions involve the transfer of electrons from one substance (the reducing agent) to another (the oxidizing agent).
  • Electrolytic Cells (Electrolysis): In electrolytic cells, non-spontaneous redox reactions are driven by an external electrical energy source, typically a battery or power supply. This energy input forces the redox reactions to proceed in a direction that would not occur spontaneously.

2. Anode and Cathode:

  • Galvanic Cells: In a galvanic cell, there are two electrodes—an anode and a cathode. At the anode, oxidation takes place, resulting in the release of electrons into the external circuit. The anode is typically the negative electrode. At the cathode, reduction occurs, with electrons being accepted from the external circuit. The cathode is usually the positive electrode.
  • Electrolytic Cells: In an electrolytic cell, the anode is positively charged, and the cathode is negatively charged. Here, the anode attracts negatively charged ions from the electrolyte, where they undergo oxidation. The cathode attracts positively charged ions from the electrolyte, where they undergo reduction.

3. Electron Flow:

  • Galvanic Cells: Electrons released during the oxidation process at the anode flow through the external circuit toward the cathode, creating an electric current that can be used to perform work, such as powering a device.
  • Electrolytic Cells: An external power source forces electrons to flow from the cathode to the anode through the external circuit. This energy input drives non-spontaneous reduction and oxidation reactions at the cathode and anode, respectively.

4. Ion Flow:

  • In both types of cells, ions flow within the electrolyte to maintain overall electrical neutrality. In galvanic cells, this ion flow occurs through a salt bridge or separator, while in electrolytic cells, it happens naturally within the electrolyte.

5. Cell Potential:

  • The difference in electric potential between the anode and cathode is known as cell potential (Ecell). In galvanic cells, this potential is positive and represents the cell’s ability to do electrical work. In electrolytic cells, the external power source determines the direction and magnitude of Ecell.

6. Energy Conversion:

  • Galvanic Cells: Chemical energy is converted into electrical energy as redox reactions proceed spontaneously. This electrical energy can be used to power devices or store energy in batteries.
  • Electrolytic Cells: Electrical energy from an external source is used to drive non-spontaneous chemical reactions, converting electrical energy into chemical energy or other forms of energy, such as heat (as in the case of water electrolysis).

In summary, electrochemical cells function by facilitating and controlling redox reactions at the anode and cathode, allowing for the conversion of chemical energy into electrical energy (galvanic cells) or electrical energy into chemical energy (electrolytic cells), depending on the specific type of cell and its purpose.

Electrochemical cell working with example

Electrochemical cells work by converting chemical energy into electrical energy (or vice versa) through a series of redox (reduction-oxidation) reactions. These cells consist of two half-cells, each with an electrode immersed in an electrolyte solution. Here’s a simplified explanation of how an electrochemical cell works using a common example: the galvanic cell, represented by a standard alkaline battery.

Components of a Galvanic Cell:

  1. Anode: The anode is one of the two electrodes. In a standard alkaline battery, it is typically made of zinc (Zn). At the anode, the oxidation half-reaction takes place: Anode: Zn(s) → Zn²⁺(aq) + 2e⁻ In this reaction, zinc metal is oxidized, losing two electrons and forming zinc ions (Zn²⁺).
  2. Cathode: The cathode is the second electrode, often composed of manganese dioxide (MnO₂) in an alkaline battery. At the cathode, the reduction half-reaction occurs: Cathode: MnO₂(s) + 2H₂O(l) + 2e⁻ → Mn(OH)₂(s) In this reaction, manganese dioxide accepts two electrons and reacts with water to form manganese hydroxide (Mn(OH)₂).
  3. Electrolyte: The electrolyte is an alkaline solution that allows the flow of ions between the anode and cathode. It contains hydroxide ions (OH⁻) in an alkaline battery, which move to maintain electrical neutrality.
  4. External Circuit: The anode and cathode are connected through an external circuit, typically via a wire. This allows the flow of electrons from the anode to the cathode through the external load (e.g., a flashlight bulb).
  5. Salt Bridge (or Separator, not shown in a basic battery): In more complex cells, like the Daniell cell, a salt bridge or separator allows the flow of ions between the half-cells without mixing their solutions, maintaining electrical neutrality.

How It Works:

  1. At the anode, zinc undergoes oxidation, releasing two electrons into the external circuit. These electrons flow through the wire, creating an electric current that powers an external device, such as a flashlight bulb.
  2. In the cathode, manganese dioxide accepts these electrons and reacts with water, forming manganese hydroxide. This reaction helps maintain electrical neutrality by consuming the electrons coming from the anode.
  3. Simultaneously, hydroxide ions (OH⁻) from the electrolyte solution move toward the anode to balance the charge generated by the formation of zinc ions (Zn²⁺) at the anode.
  4. The overall process continues until the reactants at the anode and cathode are exhausted. As the cell operates, chemical energy in the form of the reactants is converted into electrical energy to power the connected device.

In this way, a galvanic cell, like an alkaline battery, converts the chemical energy stored in its reactants (zinc and manganese dioxide) into electrical energy, providing a continuous source of power for various applications.

Half-Cells and Cell Potential

Certainly, here’s a more concise explanation:


Half-cells are the two components of an electrochemical cell where either oxidation or reduction occurs. Each consists of an electrode and an electrolyte.

  1. Anode Half-Cell: Oxidation occurs here, releasing electrons.
  2. Cathode Half-Cell: Reduction occurs here, accepting electrons.

Cell Potential (EMF):

Cell potential (Ecell) measures the potential difference between half-cells, determining if the reaction is spontaneous (positive Ecell) or requires energy input (negative Ecell). Ecell = E(cathode) – E(anode).

Primary cell and Secondary Cell

Primary cells and secondary cells are two distinct categories of batteries, each with its own characteristics and applications:

Primary Cell:

  • Single Use: Primary cells, also known as non-rechargeable or disposable cells, are designed for single use only. Once the chemical reactions in the cell are exhausted, the battery cannot be recharged, and it must be replaced.
  • Chemical Reactions: They rely on irreversible chemical reactions between the electrodes and electrolyte to generate electrical energy. These reactions deplete the cell’s energy over time.
  • Long Shelf Life: Primary cells generally have a longer shelf life and can be stored for extended periods without significant loss of energy.
  • Examples: Common examples of primary cells include alkaline batteries, zinc-carbon batteries, and lithium primary batteries.

Secondary Cell:

  • Rechargeable: Secondary cells, also known as rechargeable cells or storage batteries, can be recharged and reused multiple times by reversing the chemical reactions that occur during discharge.
  • Reversible Reactions: They employ reversible chemical reactions between the electrodes and electrolyte, allowing energy to be stored and released repeatedly.
  • Limited Lifespan: Secondary cells have a finite number of charge-discharge cycles, and their capacity gradually decreases over time. Eventually, they need to be replaced.
  • Initial Cost: They often have a higher initial cost compared to primary cells, but their cost per cycle is typically lower.
  • Examples: Common examples of secondary cells include lithium-ion batteries, nickel-cadmium (NiCd) batteries, and nickel-metal hydride (NiMH) batteries.

In summary, primary cells are single-use batteries with irreversible chemical reactions, while secondary cells are rechargeable batteries with reversible chemical reactions. The choice between these two types depends on factors such as the intended application, cost considerations, and the need for reusability.

Types of Electrochemical Cells

Certainly, let’s delve into more detail about two specific types of electrochemical cells: galvanic cells and electrolytic cells.

1. Galvanic (Voltaic) Cells:

  • Spontaneous Reactions: Galvanic cells are designed to facilitate spontaneous redox reactions. These reactions release energy, which is harnessed as electrical energy.
  • Two Half-Cells: A galvanic cell consists of two half-cells: an anode and a cathode. Each half-cell contains an electrode immersed in an electrolyte solution.
  • Anode: At the anode, oxidation occurs, leading to the loss of electrons. Electrons are released into the external circuit.
  • Cathode: At the cathode, reduction occurs, where electrons are accepted from the external circuit.
  • Salt Bridge: To maintain electrical neutrality, a salt bridge or a porous separator is used to allow ion flow between the two half-cells without mixing the solutions. This ensures the circuit remains complete.
  • Cell Potential: The difference in electrical potential between the anode and cathode is the cell potential (Ecell), measured in volts (V). It determines the cell’s ability to do electrical work.
  • Examples: Common galvanic cells include alkaline batteries (e.g., AA batteries), lead-acid batteries (e.g., car batteries), and lithium-ion batteries.

2. Electrolytic Cells:

  • Non-Spontaneous Reactions: Electrolytic cells, in contrast to galvanic cells, drive non-spontaneous redox reactions using an external source of electrical energy.
  • Two Electrodes: Electrolytic cells also have two electrodes, but the process is different. The anode is positively charged, and the cathode is negatively charged.
  • Electrolysis: Electrolytic cells are often used for processes like electrolysis, where water or other compounds are broken down into their constituent elements. For example, in water electrolysis, water is split into hydrogen and oxygen gases.
  • Power Source: An external power source (e.g., a battery or power supply) provides the necessary electrical energy to drive the non-spontaneous reaction.
  • Endothermic: Electrolytic processes are typically endothermic, meaning they absorb energy in the form of heat.
  • Applications: Electrolytic cells have a wide range of applications, including electroplating (depositing a layer of metal onto an object’s surface), metal refining, production of chemicals, and water treatment.

In summary, galvanic cells generate electrical energy through spontaneous redox reactions, while electrolytic cells consume electrical energy to drive non-spontaneous redox reactions. Both types of cells play crucial roles in various applications, from powering devices to industrial processes and chemical production.

Certainly, here’s the information about galvanic and electrolytic cells presented in a table format:

PropertyGalvanic CellElectrolytic Cell
TypeSpontaneous redox reactionsNon-spontaneous redox reactions
PurposeConverts chemical energy into electrical energyDrives non-spontaneous chemical reactions
ComponentsTwo half-cells with anode and cathode, separated by a salt bridge or separatorTwo electrodes (anode and cathode) in an electrolyte solution
AnodeOxidation occurs, electrons are released into the external circuitPositively charged electrode where oxidation occurs
CathodeReduction occurs, electrons are accepted from the external circuitNegatively charged electrode where reduction occurs
Energy SourceChemical energy from redox reactionsExternal electrical power source (e.g., battery)
Cell Potential (Ecell)Positive (produces electrical energy)Determined by the external voltage source (negative for driving reactions)
Reaction SpontaneitySpontaneous, releases energyNon-spontaneous, requires input of energy
ApplicationsBatteries (e.g., alkaline, lithium-ion)Electroplating, metal refining, chemical synthesis, water treatment
ExamplesAlkaline batteries, lead-acid batteriesElectrolysis of water to produce hydrogen and oxygen

This table provides a side-by-side comparison of the key characteristics of galvanic cells and electrolytic cells.

which reaction occurs at the anode in an electrochemical cell

In an electrochemical cell, the reaction that occurs at the anode depends on the specific cell chemistry and the overall redox (reduction-oxidation) reaction involved. The anode is where oxidation takes place, and it loses electrons during this process.

Here are a few examples of anode reactions in different types of electrochemical cells:

  1. Galvanic Cell (Battery):
  • In a standard zinc-carbon battery, the anode reaction is:
    • Anode: Zn(s) → Zn²⁺(aq) + 2e⁻
    • In this reaction, solid zinc (Zn) is oxidized to form zinc ions (Zn²⁺) and release two electrons (2e⁻).
  1. Hydrogen Fuel Cell:
  • In a hydrogen fuel cell, the anode reaction involves the oxidation of hydrogen gas (H₂):
    • Anode: H₂(g) → 2H⁺(aq) + 2e⁻
    • Hydrogen gas is oxidized to form two hydrogen ions (H⁺) and release two electrons (2e⁻).
  1. Electrolysis of Water:
  • During the electrolysis of water to produce hydrogen and oxygen gases, the anode reaction is:
    • Anode: 2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e⁻
    • Water molecules are oxidized at the anode to produce oxygen gas (O₂), hydrogen ions (H⁺), and release four electrons (4e⁻).

These are just a few examples, and the specific reaction at the anode can vary depending on the electrochemical cell and the substances involved. In general, the anode is where oxidation occurs, leading to the release of electrons into the external circuit.

where does oxidation occur in an electrochemical cell

Oxidation occurs at the anode in an electrochemical cell. The anode is one of the two electrodes in the cell, the other being the cathode. During the oxidation half-reaction at the anode:

  1. One or more reactants undergo oxidation, typically losing electrons in the process.
  2. Electrons are released into the external circuit and flow through it.
  3. Positively charged ions or cations are generated as a result of the oxidation and migrate into the electrolyte solution to maintain overall electrical neutrality within the cell.

The specific chemical reaction that occurs at the anode can vary widely depending on the type of electrochemical cell and the substances involved. However, in all cases, the anode is the site where oxidation takes place, and it serves as the source of electrons for the external circuit, ultimately leading to the flow of electrical current.

Application of Electrochemical Cell

Electrochemical cells have a wide range of applications across various fields due to their ability to convert chemical energy into electrical energy and vice versa. Here are some common applications of electrochemical cells:

  1. Batteries: Electrochemical cells are widely used in batteries to power portable electronic devices, such as smartphones, laptops, and watches. They are also used in larger-scale applications like electric vehicles (EVs) and grid energy storage systems.
  2. Fuel Cells: Fuel cells are electrochemical cells that generate electricity through the reaction of a fuel (e.g., hydrogen) with an oxidant (e.g., oxygen). They are used to power vehicles, provide backup power, and are being explored for stationary power generation.
  3. Electrolysis: Electrolytic cells are used in the process of electrolysis to split water into hydrogen and oxygen gases. This is a key technology for hydrogen production, which can be used as a clean energy source or for various industrial processes.
  4. Corrosion Protection: Electrochemical cells are employed to protect metal structures from corrosion through methods like sacrificial anodes. In these systems, a less reactive metal corrodes instead of the protected metal.
  5. Electroplating: Electroplating processes use electrochemical cells to deposit a layer of one metal onto the surface of another. This is commonly used for decorative purposes, corrosion resistance, and improving the conductivity of materials.
  6. Sensors: Electrochemical cells are utilized in various types of sensors, such as pH sensors, gas sensors (e.g., oxygen sensors in cars), and biosensors for detecting specific biological molecules.
  7. Medical Devices: Implantable medical devices like pacemakers and defibrillators use electrochemical cells as power sources to deliver electrical impulses to the heart.
  8. Environmental Monitoring: Electrochemical sensors are employed in environmental monitoring to measure parameters like the concentration of pollutants, gases, and dissolved ions in water.
  9. Analytical Chemistry: Electrochemical cells are used in analytical techniques like cyclic voltammetry and amperometry for quantitative analysis and detection of specific substances in solution.
  10. Electrochemical Machining: This process uses electrochemical cells to remove material from a workpiece. It is used in precision machining for applications like aerospace components.
  11. Energy Conversion: Some electrochemical cells, such as photovoltaic cells (solar cells), convert sunlight into electricity. These cells are crucial for solar power generation.
  12. Biomedical Research: Electrochemical cells play a role in various research applications, including studying biological processes, drug discovery, and investigating the behavior of biomolecules.
  13. Wastewater Treatment: Electrochemical cells can help in the removal of contaminants from wastewater through processes like electrocoagulation and electrooxidation.
  14. Hydrogen Storage: Electrochemical cells are studied for their potential to store hydrogen as a clean and efficient energy carrier.
  15. Desalination: Electrodialysis, an electrochemical process, is used in desalination to remove salt and impurities from brackish water.

These applications highlight the versatility of electrochemical cells in meeting the energy and chemical needs of modern society while addressing environmental and technological challenges. Researchers continue to explore and develop new applications for these cells in various fields.

Advantages of electrochemical cell

Electrochemical cells offer several advantages that make them highly valuable for various applications in science, technology, and industry. Here are some of the key advantages of electrochemical cells:

  1. Efficiency: Electrochemical cells are highly efficient at converting chemical energy into electrical energy and vice versa. This efficiency is crucial in applications where energy conversion is essential, such as batteries and fuel cells.
  2. Portable Power: Electrochemical cells, especially batteries, provide a portable and reliable source of electrical energy. They are commonly used in portable electronic devices, ensuring they can operate without a constant external power source.
  3. Clean Energy: Fuel cells, which are a type of electrochemical cell, produce electricity with minimal pollution. Hydrogen fuel cells, for example, emit only water vapor as a byproduct, making them a clean energy option for vehicles and power generation.
  4. Longevity: Many electrochemical cells, such as lithium-ion batteries, have a long cycle life, allowing them to be recharged and used for extended periods before needing replacement.
  5. Scalability: Electrochemical cells can be designed and scaled to meet specific energy and power requirements, making them suitable for both small-scale and large-scale applications.
  6. Fast Response: Electrochemical sensors and supercapacitors have fast response times, making them suitable for applications that require rapid measurement or energy discharge.
  7. Versatility: Electrochemical cells can be tailored to work with a wide range of materials and electrolytes, allowing for flexibility in design and application.
  8. Environmentally Friendly: Some electrochemical processes, like electrolysis, can be used to generate green hydrogen or other valuable products, contributing to environmental sustainability.
  9. Low Self-Discharge: Many modern batteries and supercapacitors have low self-discharge rates, allowing them to retain stored energy for extended periods without significant loss.
  10. Reliability: Electrochemical cells are known for their reliability and stability, making them suitable for critical applications such as medical devices, emergency backup power systems, and aerospace technology.
  11. High Energy Density: Certain types of electrochemical cells, like lithium-ion batteries, offer high energy density, allowing them to store a significant amount of energy in a relatively compact space.
  12. Safety Features: Advances in electrochemical cell design have led to the incorporation of safety features, such as overcharge protection and thermal management systems, making them safer for everyday use.
  13. Redox Flexibility: The redox reactions in electrochemical cells can involve a wide range of chemicals, allowing for the development of various cell chemistries tailored to specific applications.

These advantages make electrochemical cells indispensable in a wide array of fields, from powering everyday consumer electronics to enabling clean energy solutions and facilitating critical scientific research. Ongoing research and development continue to improve the efficiency, reliability, and environmental sustainability of electrochemical cells, further expanding their potential applications.

disadvantages of Electrochemical cell

While electrochemical cells offer many advantages, they also have certain disadvantages and limitations that need to be considered in various applications. Here are some of the disadvantages of electrochemical cells:

  1. Limited Energy Density: Electrochemical cells may have lower energy density compared to other energy storage technologies, such as fossil fuels. This can limit the amount of energy that can be stored in a given volume or weight, especially in some battery types.
  2. Finite Lifespan: All electrochemical cells have a finite lifespan, and their performance can degrade over time due to chemical reactions and wear and tear. This degradation can lead to reduced capacity and efficiency, eventually requiring replacement.
  3. Slow Charging: Some electrochemical cells, like traditional lead-acid batteries, can have slow charging times, which may not be suitable for applications requiring rapid replenishment of energy.
  4. Environmental Impact: The production and disposal of certain electrochemical cell components, such as lithium-ion batteries, can have environmental impacts, including the mining and extraction of raw materials and the recycling or disposal of used cells.
  5. Chemical Compatibility: Electrochemical cells require specific materials for electrodes and electrolytes that are chemically compatible with the intended reactions. This limits the range of materials that can be used and may pose challenges in certain applications.
  6. Temperature Sensitivity: Electrochemical cells can be sensitive to temperature variations. Extreme heat or cold can affect their performance and, in some cases, reduce their lifespan.
  7. Safety Concerns: Electrochemical cells can pose safety risks, especially if damaged or abused. Overcharging, short circuits, or physical damage can lead to thermal runaway, fires, or explosions in certain battery chemistries.
  8. Cost: Some electrochemical cell technologies, particularly advanced ones like solid-state batteries or certain fuel cell systems, can be expensive to manufacture and may not be cost-effective for all applications.
  9. Limited Energy Output: Electrochemical cells are generally suited for providing low to moderate continuous power. High-power demands may require additional systems, such as multiple cells connected in parallel or energy storage solutions like supercapacitors.
  10. Chemical Waste: In some cases, electrochemical cells can produce chemical waste or byproducts that need to be managed or disposed of properly. For example, certain battery chemistries can contain toxic materials.
  11. Complex Maintenance: Some electrochemical systems, like fuel cells, may require specialized maintenance and periodic replacement of components, adding to their operational costs and complexity.
  12. Infrastructure Challenges: Widespread adoption of certain electrochemical technologies, such as hydrogen fuel cells or fast-charging electric vehicle infrastructure, may require significant investments in new infrastructure and distribution networks.

It’s important to note that the disadvantages of electrochemical cells can vary depending on the specific technology, application, and use case. Engineers and researchers continually work on addressing these limitations and improving the performance, safety, and sustainability of electrochemical cell technologies.

Reference : https://www.sciencedirect.com/topics/materials-science/electrochemical-cell

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