Electrochemical Cell:
Definition, Types and Examples from Batteries to Electrolysis
Electrochemical cells are systems that convert energy of chemical reactions into electrical energy. Conversely, electrical energy can also be used for electrochemical reactions.
What is an Electrochemical Cell?
Typical electrochemical cells are batteries, accumulators, fuel cells or electrolytic cells. An electrochemical cell consists of two so-called half-cells, each half-cell consisting of an electrode immersed in an electrolyte. Both half-cells can be immersed in the same electrolyte or in different electrolytes.
Two half-cells in different electrolytes must be ionically connected via an electrolyte bridge, also known as a salt bridge. The easiest variant is a filter paper soaked in KNO3 or KCl. Others are made of glass. A special form of salt bridge is the Haber-Luggin capillary. Every reference electrode with an interior electrolyte operates with a salt bridge. And by the way: every reference electrode is a half-cell.
The chemical reactions in a half-cell can affect the electrolyte, the electrode, or components supplied from outside such as hydrogen gas in fuel cells. Two half-cells that are connected ionically and electrically form a full cell. One of the half-cells emits electrons (oxidation). These electrons flow through the outer circuit into the second half-cell, where corresponding reductions processes take place. The salt bridge allows a flow of anions and cations of the electrolyte so that an equilibrium between the half-cells can occur. Without this ionic connection, the load difference caused by the electron flow would bring the reactions to a standstill.
A typical half-cell which is known in the electrochemistry is the so-called Daniell cell. It consists of a zinc rod in a zinc sulphate solution and a copper rod in a copper sulphate solution. For the ion exchange, both half-cells are connected via a salt bridge. If a multimeter is interposed between the half-cells, a cell voltage of approximately 1.1 Volt can be read off. On each half-cell (often referred to as a half-element), a potential sets, which can be measured by means of a reference electrode. The Daniel element or Daniel cell is a galvanic cell. It is a battery - a primary cell - of historical and educational signficance.
A half-cell consists of an electrode which immerses in an electrolyte.
What are Batteries?
A battery is an electrochemical cell that converts chemical energy into electrical energy through redox reactions (reduction and oxidation).
Oxidation (loss of electrons) happens at the anode. Reduction (gain of electrons) happens at the cathode. Electrons flow through an external circuit, while ions flow inside the battery through the electrolyte.
Basic Components:
Anode (- terminal): The electrode where oxidation occurs, releasing electrons.
Cathode (+ terminal): The electrode where reduction occurs, accepting electrons.
Electrolyte: The medium (liquid, gel, or solid) that allows ions to move but not electrons.
How do Batteries work?
Non-rechargeable batteries are so-called primary batteries.
A chemical reaction at the anode releases electrons and ions. The electrons flow through the external circuit. The ions move through the electrolyte to the cathode. At the cathode, they recombine with electrons. This continues until the chemical reactants are used up. The battery is discharged.
In rechargeable batteries, the discharging-process can be reversed using external power (charging).
Rechargeable batteries are so-called secondary batteries or accumulators.
Types of Batteries
Batteries are chemical systems designed to control redox reactions to ensure a steady, usable flow of electrons.
There are many different types of batteries.
Primary Cells (non-rechargeable): Alkaline, zinc-carbon, lithium primary
Secondary Cells (rechargeable): Nickel-metal hydride, lead-acid, lithium-ion
Batteries are storing chemical energy inside and release electricity until the stored reactants are used up. After that, they must be replaced or recharged.
Fuel Cells are generating electricity continuously as long as they are supplied with external fuel (like hydrogen) and an oxidant (like oxygen).
In short, batteries are energy storage devices, while fuel cells are energy conversion devices.
What are Fuel Cells?
Fuel cells are devices that convert chemical energy into electrical energy through an electrochemical reaction, typically between hydrogen and oxygen. The reaction occurs in the presence of an electrolyte, and the byproducts are usually water and heat, making it a very clean energy source compared to conventional combustion-based methods.
Fuel cells are considered a promising technology for clean energy, especially in sectors like transportation (hydrogen cars, buses) and stationary power generation. The main hurdles currently are cost reduction, hydrogen infrastructure, and improving long-term durability.
Basic Components:
- Anode: The negative side of the fuel cell, where the fuel (often hydrogen) enters.
- Cathode: The positive side of the fuel cell, where oxygen (or air) enters.
- Electrolyte: A substance that allows ions to move between the anode and cathode. In many fuel cells, this is a proton-conducting material like the proton exchange membrane (PEM).
- External Circuit: Where electrons flow from the anode to the cathode, generating electricity.
How do Fuel Cells work?
- Hydrogen is supplied to the anode: The hydrogen molecules are split into protons (H⁺) and electrons (e⁻) by a catalyst (often platinum).
- Electrons flow through the external circuit: The electrons move through an external circuit, creating an electric current that can be used to power devices.
- Protons pass through the electrolyte: The protons (hydrogen ions) travel through the electrolyte towards the cathode.
- Oxygen is supplied to the cathode: Oxygen molecules (from air) combine with the electrons and protons at the cathode to form water (H₂O), which is the main byproduct.
The overall reaction looks like this: 2 H2 + O2 → 2 H2O + Electricity + Heat.
Fuel Cells with Alkaline Electrolytes - AFC (Alkaline Fuel Cell)
The overall reaction of an Alkaline Fuel Cell (AFC) involves the electrochemical conversion of hydrogen and oxygen into water, producing electricity and heat in the process.
Hydrogen is supplied to the anode. Oxygen (usually from air) is supplied to the cathode. The Hydroxide ions (OH-) can pass pass through the Membrane (Alkaline Exchange Membrane AEM). Electrons (e⁻) flow through the external circuit, generating electric current. Water and heat are produced at the cathode.
Anode: Hydrogen Oxidation (Hydrogen Oxidation Reaction HOR)
Hydrogen reacts with hydroxide ions (OH⁻) from the electrolyte to form water and release electrons.
2 H2 + 4 OH- → 4 H2O + 4 e-
Cathode: Oxygen Reduction (Oxygen Reduction Reaction ORR)
Oxygen reacts with water and electrons (from the external circuit) to form hydroxide ions (OH⁻), which migrate back to the anode.
O2 + 2 H2O + 4 e- → 4 OH-
Overall Reaction:
2 H2 + O2 → 2 H2O + Electricity + Heat
Fuel Cells with Acidic Electrolytes - PEMFC (Proton Exchange Membrane Fuel Cell)
The overall reaction of a Proton Exchange Membrane (PEM) Fuel Cell involves the electrochemical conversion of hydrogen and oxygen into water, producing electricity and heat in the process.
Hydrogen is supplied to the anode. Oxygen (usually from air) is supplied to the cathode. The Protons (H⁺) can pass pass through the Membrane (Proton Exchange Membrane PEM). Electrons (e⁻) flow through the external circuit, generating electric current. Water and heat are produced at the cathode.
Anode: Hydrogen Oxidation (Hydrogen Oxidation Reaction HOR)
Hydrogen gas is split into protons (H+) and electrons (e-).
2 H2 → 4 H+ + 4 e-
Cathode: Oxygen Reduction (Oxygen Reduction Reaction ORR)
Oxygen gas reacts with protons (from the anode via the membrane) and electrons (from the external circuit) to form water.
O2 + 4 H+ + 4 e- → 2 H2O
Overall Reaction:
2 H2 + O2 → 2 H2O + Electricity + Heat
Key Differences between AFC and PEMFC
Feature
|
Alkaline Fuel Cell (AFC)
|
PEM Fuel Cell
|
|---|---|---|
Electrolyte | Aqueous KOH (alkaline) | Proton Exchange Membrane |
Ion Conducted | OH⁻ (hydroxide ions) | H⁺ (protons) |
Water generated on | Anode | Cathode |
Operating Conditions | ~60–100°C | ~60–80°C |
Oxygen Source | Needs pure O₂ (CO₂-sensitive) | Pure O₂ or air |
CO₂ Sensitivity | High (CO₂ poisons electrolyte) | Low |
Types of Fuel Cells
There are several types of fuel cells, but the most common ones are:
Proton Exchange Membrane Fuel Cell (PEMFC):
PEMFCs have a quick start-up and relatively low operating temperatures. They are often used in transportation (like hydrogen-powered cars).
Phosphoric Acid Fuel Cell (PAFC):
PAFCs operate at moderate temperatures and are known for their reliability. They are used in stationary applications.
Solid Oxide Fuel Cell (SOFC):
SOFCs operate at higher temperatures and can use a variety of fuels, including natural gas. They are used for stationary power generation.
Molten Carbonate Fuel Cell (MCFC):
MCFCs use molten carbonate salts as electrolytes. They are suitable for large-scale power generation.
Advantages of Fuel Cells
- Clean Energy: The primary byproducts are water and heat, making them environmentally friendly.
- High Efficiency: More efficient than combustion engines since they do not rely on burning fuel.
- Scalable: Can be used in a wide range of applications, from small portable devices to large power plants.
Challenges
- Cost: Fuel cells, particularly PEM cells, require expensive catalysts (like platinum), and the technology is still costly.
- Hydrogen Storage and Infrastructure: Storing and transporting hydrogen safely is a challenge, and the fueling infrastructure is limited.
- Durability: Fuel cells can degrade over time, especially if they are exposed to constant heat and humidity.
Fuel cells are combining hydrogen and oxygen to produce water.
An electrolysis cell does the opposite. Electrolysis is the process of using electricity to drive a chemical reaction, usually splitting water into hydrogen and oxygen.
What is Electrolysis?
Electrolysis is a chemical process that uses electricity to drive a non-spontaneous chemical reaction. It typically involves passing an electric current through an electrolyte, causing chemical changes at the electrodes.
How does Electrolysis work?
Electrolyte: A liquid that conducts electricity, usually a solution of salts, acids, or bases.
Electrodes: Two conductors placed in the electrolyte:
Anode (+): Where oxidation occurs (loss of electrons).
Cathode (−): Where reduction occurs (gain of electrons).
Power Source: Provides the energy to push electrons and force the reaction.
Typical Applications of Electrolysis
- Electroplating (coating objects with a metal)
- Producing gases (like hydrogen, oxygen, chlorine)
- Extraction of metals (e.g., aluminium from bauxite)
- Purifying metals (like copper)
Water Electrolysis
Water electrolysis is the process of splitting water (H₂O) into its basic elements — hydrogen (H2) and oxygen (O2) — using an electric current. It is an important method for producing clean hydrogen fuel, especially when powered by renewable energy.
When you pass electricity through water containing a small amount of an electrolyte (like sulfuric acid for better electrical conductivity), it decomposes into:
- Hydrogen gas (H2) at the cathode (Hydrogen Evolution Reaction HER)
- Oxygen gas (O2) at the anode (Oxygen Evolution Reaction OER)
Reaction: 2 H2O (l) → 2 H2 (g) + O2 (g)
Basic Setup
Electrolyte Water: Water by itself does not conduct electricity well, so an electrolyte is needed to provide ions for the circuit. Often sulfuric acid is used to improve conductivity.
Electrodes: Two solid conductors.
Cathode (−): Reduction (hydrogen is produced)
Anode (+): Oxidation (oxygen is produced)
Diaphragm/Separator: The diaphragm or separator separates the two reaction zones to avoid the mixing of both gases.
Power Supply: Provides the voltage to drive the reaction.
Electrochemical Reaction of Water Splitting in Detail
The reactions in acidic solutions differ from those in alkaline solutions.
Electrolysis in Acidic Solutions
Cathode (−): reduction
Electrons are gained by hydrogen ions (H⁺) to form hydrogen gas.
2 H+ (aq) + 2 e- → H2 (g)
Anode (+): oxidation
Water molecules lose electrons to form oxygen gas and hydrogen ions.
2 H2O (l) → O2 (g) + 4 H+ (aq) + 4 e-
Overall Reaction:
2 H2O (l) → 2 H2 (g) + O2 (g)
Hydrogen and oxygen are produced in a 2:1 volume ratio, because each water molecule has 2 hydrogen atoms and 1 oxygen atom.
Electrolysis in Alkaline Solutions
Cathode (−): reduction
Electrons are gained by water to form hydrogen gas and hydroxide ions (OH-).
4 H2O (l) + 4 e- (aq) → 2 H2 (g) + 4 OH- (aq)
Anode (+): oxidation
Hydroxide ions (OH-) lose electrons to form oxygen gas and water molecules.
4 OH- (aq) → O2 (g) + 2 H2O (l) + 4 e-
Overall Reaction:
2 H2O (l) → 2 H2 (g) + O2 (g)
Hydrogen and oxygen are produced in a 2:1 volume ratio, because each water molecule has 2 hydrogen atoms and 1 oxygen atom.
Batteries, fuel cells, and electrolyzers are called full cells, which always consist of of two half-cells connected through an ionic conductor. These full cells can be analyzed as a whole by measuring current and cell voltage.
It is also very useful to study the behavior of the individual parts separately. Our ElyFlow and FlexCell systems are designed for examining specific reactions that take place at either the anode or the cathode.
In half-cell experiments, the focus is on the potential of the working electrode (the sample being tested). To measure this potential, a reference electrode—such as the HydroFlex hydrogen reference electrode—is needed. This setup uses two electrodes.
If the current also needs to be measured, a third electrode is added, called the counter electrode, resulting in a three-electrode setup.
Two-electrode or Three-electrode set-up?
To measure electrochemical potentials, you need the object to be studied as well as a second electrode with a known and constant potential, a so-called reference electrode. In this case, you have a simple two electrodes set-up, and the measurement happens quasi currentless.
If you want to measure the potential and the current, you have to choose a three-electrode set-up. The third electrode is the so-called counter electrode which is used to detect the electrical current. In this set-up, you must insert the reference electrode with a Haber-Luggin capillary to minimize the voltage drop between the working electrode and the counter electrode and to avoid disruptions which are caused by the current flowing between the working and the counter electrode.
In both cases, it is important to choose the right reference electrode – a large number of reference electrodes is available. Most of them work with an interior electrolyte which contains chlorides, such as potassium chloride. In that case, you will contaminate your measuring solution with potassium chloride. This can cause major problems, especially if you want to study corrosion processes, because chloride can strengthen the corrosion of your material. It is best to work with an electrolyte-free reference electrode (indicator electrode). We recommend the hydrogen reference electrode HydroFlex since it works without an interior electrolyte and every electrochemist should work with it. There are no mistakes caused by diffusion voltages.
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How do electrochemical potentials arise?
You want to measure electrochemical potentials – but how do they arise? Also learn more about reference electrodes of the second kind and salt bridges. Or do you need help with basic terms like activity, concentration, potential, and voltage?
Electrochemical test cell ElyFlow
ElyFlow is a test cell optimized for electrolyte circuits for electrochemistry with a three-electrode set-up. Here you can find some general information.
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The Importance of Field Lines
The potentials which are to be measured are only identical on the so-called equipotential surfaces.
Homogeneous electric field lines are required for correct measurements. Here you can learn more about parallel field lines, equipotential surfaces and IR-drop.
The Importance of Haber-Luggin Capillaries
The electrolytical contact between the working and the reference electrode is made by the Haber-Luggin capillary. Haber-Luggin capillaries are reducing the IR-Drop between reference electrode and working electrode.
Electrolyte Key – Salt Bridge
An electrolyte bridge, also known as a salt bridge, is used for the contacting of various electrolytes. Commercial reference electrodes rely on salt bridges.
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Electrochemical test cell FlexCell-PTFE
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Electrochemical test cell FlexCell-PP
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Electrochemical test cell ElyFlow PTFE
ElyFlow with its three-electrode set-up is optimized for electrolyte circuits.