Electrochemical cell

An electrochemical battery is a device capable of obtaining electrical energy from chemical reactions, or producing chemical reactions through the introduction of electrical energy. A common example of an electrochemical cell is the standard 1.5 volt “battery”. In reality, a “cell” is a single galvanic cell, while a battery consists of several cells connected in series.


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  • 1 File
    • 1 Notation of a stack
  • 2 types of cells
  • 3 Half cells in a voltaic cell
  • 4 Equilibrium reaction
  • 5 Electrode potentials
  • 6 Types of galvanic cells
    • 1 Primary galvanic cells
    • 2 Secondary galvanic cells
  • 7 Fuente
  • 8 external links


Batteries are irreversible primary elements, such as flashlight or radio batteries, while cells, batteries or accumulators are reversible secondary elements, such as in cars, forklifts, telephone plants or photovoltaic solar installations in doctor’s offices. family doctor, social circles or rural schools.

Irreversible means that the chemical-electrical reactions that occur in the discharge are in one direction until the electrodes are completely exhausted.

In reversible cells, the process occurs in two directions, that is, when discharging, the cell delivers current and when charging, the cell receives current that is chemically stored.

Notation of a stack

The material of the negative electrode is first written using its chemical symbol, then a vertical line is placed to separate the electrode from the ion solution of given concentration.

Immediately the symbol II is used to indicate the salt bridge, then the second ionic solution of known concentration is placed, the sign II and finally the sign of the metal:

Zn (.) I Zn 2+ (1mol/L) II Cu 2+ (1mol/L) I Cu (.)

These are all conventional cell notations

The cells that produce electrical energy are called galvanic cells, which are caused by the consumption of chemical energy.

The cells that consume energy are the electrochemical cells; they consume current from an external current source, thereby storing electrolytic chemical energy.

cell types

There are two fundamental types of cells and in both a redox reaction takes place, and the conversion or transformation of one type of energy into another:

  • The voltaic cell transforms a spontaneous chemical reaction into an electrical current, like batteries. They also receive the names of galvanic cell, galvanic battery or voltaic battery. They are widely used, so most of the examples and images in this article refer to them.
  • The electrolytic cell transforms an electrical current into a chemical oxidation-reduction reaction that does not take place spontaneously. In many of these reactions, a chemical substance is broken down, which is why this process is called electrolysis. They also receive the names of electrolytic cell or electrolytic tank. Unlike the voltaic cell, in the electrolytic cell, the two electrodes do not need to be separated, so there is only one vessel in which the two half-reactions take place.

Half cells in a voltaic cell

A galvanic cell or voltaic cell consists of two half-cells electrically connected by a metallic conductor, and also by a salt bridge. Each half cell consists of an electrode and an electrolyte. The two half cells may use the same electrolyte, or they may use different electrolytes. Chemical reactions in the cell may involve the electrolyte, the electrodes, or an external substance (as in fuel cells that may use hydrogengas as reactant). In a complete voltaic cell, chemical species in one half-cell lose electrons (oxidation) to their electrode while species in the other half-cell gain electrons (reduction) from their electrode. A salt bridge is often used to provide an ionic contact between the two half-cells with different electrolytes, to prevent the solutions from mixing and causing unwanted side reactions. This salt bridge may simply be a strip of filter paper soaked in saturated potassium nitrate solution. Other devices to achieve solution separation are porous vessels and gelled solutions. A porous container is used in the Bunsen stack (right).

They are also called semi-reactions because in each of them a part of the redox reaction takes place:

  • The loss of electrons (oxidation) takes place at the anode.
  • The gain of electrons (reduction) at the cathode.

equilibrium reaction

Each half-cell has a characteristic voltage called the half-cell potential or reduction potential. The different substances that can be chosen for each half-cell give rise to different potential differences of the whole cell, which is the parameter that can be measured. You can not measure the potential of each half cell, but the difference between the potentials of both. Each reaction is undergoing an equilibrium reaction between the different oxidation states of the ions; when equilibrium is reached, the cell cannot provide any more voltage. In the half-cell that is undergoing oxidation, the closer to equilibrium the ion/atom with the most positive oxidation state is, the more potential this reaction will give. Similarly, in the reduction reaction,

electrode potentials

The potential or electromotive force of a cell can be predicted through the use of the electrode potentials, the voltages of each half cell. The voltage difference between the reduction potentials of each electrode gives a prediction for the measured cell potential.

Cell potentials have a possible range from 0 to 6 volts. Batteries using electrolytes dissolved in water generally have cell potentials less than 2.5 volts, since the very strong oxidizing and reducing agents, which would be required to produce a higher potential, tend to react with water.

Types of galvanic cells

Galvanic cells or cells are classified into two broad categories:

  • Primary cells transform chemical energy into electrical energy, irreversibly (within the limits of practice). When the initial amount of reactants present in the cell is depleted, energy cannot be easily restored or returned to the electrochemical cell by electrical means.
  • The secondary cells can be recharged, that is, they can reverse their chemical reactions by supplying electrical energy to the cell, until their original composition is restored.

Primary galvanic cells

Primary galvanic cells can produce current immediately after connection. Disposable batteries are intended to be used only once and are disposed of later. Disposable batteries cannot be reliably recharged as the chemical reactions are not easily reversible and the active materials cannot return to their original form.

They generally have higher energy densities than rechargeable cells,4 but disposable cells do not do well in high-drain applications with loads less than 75 ohms (75 Ω).

Secondary galvanic cells

Secondary galvanic cells must be charged before use; they are usually assembled with active materials and objects in the low-energy (discharged) state. Rechargeable galvanic cells or secondary galvanic batteries can be regenerated (colloquially, recharged) by applying an electrical current, which reverses the chemical reactions that occur during use. Devices for the adequate supply of such currents that regenerate the active substances contained in the cell or battery are inappropriately called chargers or rechargers.

The oldest form of rechargeable battery is the lead-acid battery. This electrochemical cell is remarkable as it contains an acidic liquidin a sealed container, which requires the cell to be kept upright and the area to be well ventilated to ensure safe dispersal of the hydrogen gas produced by these cells during overload. The lead-acid cell is also very heavy for the amount of electrical power it can supply. Despite this, its low manufacturing cost and high surge current levels make its use common when a large capacity (greater than 10A h) is required or when weight and poor handling are not important. Lead-acid battery with absorbent glass felt cells, showing apart the two electrodes and, in the middle, the absorbent glass material that prevents acid spills.automobile as a substitute for the lead-acid wet cell, because it is maintenance-free. The VRLA cell uses immobilized sulfuric acid as the electrolyte, reducing the possibility of leaks and extending service life.7 Immobilization of the electrolyte has been achieved, generally in one of two ways:


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