ELECTROCHEMICAL CELLS: COMPONENTS, HOW THEY WORK, TYPES, EXAMPLE - CHEMISTRY - 2025

The electrochemical cells are devices in which chemical reactions pass where the chemical energy is converted into electrical energy or vice versa. These cells make up the heart of electrochemistry, the soul being the potential exchange of electrons that can occur, spontaneously or not, between two chemical species.

One of the two species oxidizes, loses electrons, while the other is reduced, gaining the transferred electrons. Commonly, the species that is reduced is a metallic cation in solution, which by gaining electrons ends up being electrically deposited on an electrode made of the same metal. On the other hand, the species that oxidizes is a metal, turning into metal cations.

Diagram for an electrochemical cell from Daniel. Source: Rehua

For example, the image above represents Daniel's cell: the simplest of all electrochemical cells. The metallic zinc electrode oxidizes, releasing Zn ²⁺ cations into the aqueous medium. This occurs in the ZnSO ₄ container on the left.

On the right, the solution containing CuSO ₄ is reduced, transforming the Cu ²⁺ cations into metallic copper that is deposited on the copper electrode. During the development of this reaction, the electrons travel through an external circuit activating its mechanisms; and therefore, providing electrical energy for the operation of a team.

Components of electrochemical cells

Electrodes

Electrical currents are generated or consumed in electrochemical cells. To ensure an adequate flow of electrons there must be materials that are good conductors of electricity. This is where the electrodes and the external circuit come in, provided with copper, silver or gold wiring.

The electrodes are the materials that provide the surface where the reactions will take place in the electrochemical cells. There are two types depending on the reaction that occurs in them:

-Anode, electrode where oxidation occurs

-Cathode, electrode where reduction occurs

The electrodes can be made of a reacting material, as in the case of Daniel's cell (zinc and copper); or, of an inert material, as it happens when they are made of platinum or graphite.

The electrons released by the anode must reach the cathode; but not through a solution, but through a metallic cable that joins both electrodes to an external circuit.

Electrolyte dissolution

The solution that surrounds the electrodes also plays an important role, as it is enriched with strong electrolytes; such as: KCl, KNO ₃, NaCl, etc. These ions favor to a certain extent the migration of electrons from the anode to the cathode, as well as their conduction through the vicinity of the electrodes to interact with the species to be reduced.

Sea water, for example, conducts electricity much better than distilled water, with a lower concentration of ions. That is why electrochemical cells have strong electrolyte dissolution among their components.

Saline bridge

The ions of the solution begin to surround the electrodes causing a polarization of the charges. The solution around the cathode begins to become negatively charged, as the cations are being reduced; in the case of the Daniel cell, the Cu ²⁺ cations when deposited as metallic copper on the cathode. Thus, there begins to be a deficit of positive charges.

This is where the salt bridge intervenes to balance the charges and prevent the electrodes from polarizing. Towards the side or compartment of the cathode, cations of the salt bridge will migrate, either K ⁺ or Zn ²⁺, to supplant the Cu ²⁺ consumed. Meanwhile, NO ₃ ^- anions will migrate from the salt bridge towards the anode compartment, to neutralize the increasing concentration of Zn ²⁺ cations.

The salt bridge is composed of a saturated solution of salts, with its ends covered by a gel that is permeable to ions, but impermeable to water.

Types of electrochemical cells and how they work

How an electrochemical cell works depends on what type it is. There are basically two types: galvanic (or voltaic) and electrolytic.

Galvanic

Daniel's cell is an example of a galvanic electrochemical cell. In them the reactions occur spontaneously and the potential of the battery is positive; the greater the potential, the more electricity the cell will supply.

Cells or batteries are precisely galvanic cells: the chemical potential between the two electrodes is transformed into electrical energy when an external circuit intervenes that connects them. Thus, the electrons migrate from the anode, ignite the equipment to which the battery is connected, and are returned directly to the cathode.

Electrolytic

Electrolytic cells are those whose reactions do not occur spontaneously, unless they are supplied with electrical energy from an external source. Here the opposite phenomenon occurs: electricity allows non-spontaneous chemical reactions to develop.

One of the best known and most valuable reactions that takes place within this type of cell is electrolysis.

Rechargeable batteries are examples of electrolytic and galvanic cells: they are recharged to reverse their chemical reactions and reestablish the initial conditions for reuse.

Examples

Daniel's cell

The following chemical equation corresponds to the reaction in Daniel's cell where zinc and copper participate:

Zn (s) + Cu ²⁺ (aq) → Zn ²⁺ (aq) + Cu (s)

But the cations Cu ²⁺ and Zn ²⁺ are not alone but accompanied by the anions SO ₄ ^2-. This cell can be represented as follows:

Zn - ZnSO ₄ - - CuSO ₄ - Cu

Daniel's cell can be built in any laboratory, being very recurrent as a practice in the introduction of electrochemistry. As Cu ²⁺ is deposited as Cu, the blue color of the CuSO ₄ solution will gradually fade.

Platinum hydrogen cell

Imagine a cell that consumes hydrogen gas, produces metallic silver, and at the same time supplies electricity. This is the platinum and hydrogen cell, and its general reaction is as follows:

2AgCl (s) + H ₂ (g) → 2Ag (s) + 2H ⁺ + 2Cl ^-

Here in the anode compartment we have an inert platinum electrode, submerged in water and pumped into gaseous hydrogen. H ₂ is oxidized to H ⁺ and gives up its electrons to the milky AgCl precipitate in the cathode compartment with a metallic silver electrode. On this silver the AgCl will be reduced and the mass of the electrode will increase.

This cell can be represented as:

Pt, H ₂ - H ⁺ - - Cl ^-, AgCl - Ag

Downs cell

And finally, among the electrolytic cells we have that of fused sodium chloride, better known as the Downs cell. Here, electricity is used to move a volume of molten NaCl through the electrodes, thus causing the following reactions:

2Na ⁺ (l) + 2e ^- → 2Na (s) (cathode)

2Cl ^- (l) → Cl ₂ (g) + 2e ^- (anode)

2NaCl (l) → 2Na (s) + Cl ₂ (g) (global reaction)

Thus, thanks to electricity and sodium chloride, metallic sodium and chlorine gas can be prepared.

References

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