Electrochemical cell

From Wikipedia, the free encyclopedia
Jump to navigation Jump to search
A demonstration electrochemical cell setup resembling the Daniell cell. The two half-cells are linked by a salt bridge carrying ions between them. Electrons flow in the external circuit.

An electrochemical cell (EC) is a device capable of either generating electrical energy from chemical reactions or facilitating it otherwise. The electrochemical cells which generate an electric current are called voltaic cells or galvanic cells and the other ones are called electrolytic cells which are used to drive chemical reactions like electrolysis.[1][2] A common example of an galvanic cells is a standard 1.5 - volt[3] cell meant for consumer use. A battery consists of one or more cells, connected in either parallel or series pattern.[4]

Galvanic cell

Electrolytic cell


The Bunsen cell, invented by Robert Bunsen.

An electrochemical cell consists of two half-cells. 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. The chemical reactions in the cell may involve the electrolyte, the electrodes, or an external substance (as in fuel cells that may use hydrogen gas as a reactant). In a full electrochemical cell, species from one half-cell lose electrons (oxidation) to their electrode while species from the other half-cell gain electrons (reduction) from their electrode.

A salt bridge (e.g., filter paper soaked in KNO3, NaCl, or some other electrolyte) is often employed to provide ionic contact between two half-cells with different electrolytes, yet prevent the solutions from mixing and causing unwanted side reactions. An alternative to a salt bridge is to allow direct contact (and mixing) between the two half-cells, for example in simple electrolysis of water.

As electrons flow from one half-cell to the other through an external circuit, a difference in charge is established. If no ionic contact were provided, this charge difference would quickly prevent the further flow of electrons. A salt bridge allows the flow of negative or positive ions to maintain a steady-state charge distribution between the oxidation and reduction vessels, while keeping the contents otherwise separate. Other devices for achieving separation of solutions are porous pots and gelled solutions. A porous pot is used in the Bunsen cell (right).

Equilibrium reaction

Each half-cell has a characteristic voltage. Various choices of substances for each half-cell give different potential differences. Each reaction is undergoing an equilibrium reaction between different oxidation states of the ions: When equilibrium is reached, the cell cannot provide further voltage. In the half-cell that is undergoing oxidation, the closer the equilibrium lies to the ion/atom with the more positive oxidation state the more potential this reaction will provide. Likewise, in the reduction reaction, the closer the equilibrium lies to the ion/atom with the more negative oxidation state the higher the potential.

Cell potential

The cell potential can be predicted through the use of electrode potentials (the voltages of each half-cell). These half-cell potentials are defined relative to the assignment of 0 volts to the standard hydrogen electrode (SHE). (See table of standard electrode potentials). The difference in voltage between electrode potentials gives a prediction for the potential measured. When calculating the difference in voltage, one must first rewrite the half-cell reaction equations to obtain a balanced oxidation-reduction equation.

  1. Reverse the reduction reaction with the smallest potential ( to create an oxidation reaction/ overall positive cell potential)
  2. Half-reactions must be multiplied by integers to achieve electron balance.

Cell potentials have a possible range of roughly zero to 6 volts. Cells using water-based electrolytes are usually limited to cell potentials less than about 2.5 volts due to high reactivity of the powerful oxidizing and reducing agents with water that needed to produce a higher voltage. Higher cell potentials are possible with cells using other solvents instead of water. For instance, lithium cells with a voltage of 3 volts are commonly available.

The cell potential depends on the concentration of the reactants, as well as their type. As the cell is discharged, the concentration of the reactants decreases and the cell potential also decreases.

See also


  1. ^ "Electrolytic Cells". hyperphysics.phy-astr.gsu.edu. Retrieved 2018-05-17. 
  2. ^ "Electrochemical Cells". hyperphysics.phy-astr.gsu.edu. Retrieved 2018-05-17. Electrochemical cells which generate an electric current are called voltaic cells or galvanic cells, and common batteries consist of one or more such cells. In other electrochemical cells an externally supplied electric current is used to drive a chemical reaction which would not occur spontaneously. Such cells are called electrolytic cells. 
  3. ^ "Byjus ( a tutorial site)". 
  4. ^ "Electrochemical Cells". hyperphysics.phy-astr.gsu.edu. 
Retrieved from "https://en.wikipedia.org/w/index.php?title=Electrochemical_cell&oldid=846166080"
This content was retrieved from Wikipedia : http://en.wikipedia.org/wiki/Electrochemical_cell
This page is based on the copyrighted Wikipedia article "Electrochemical cell"; it is used under the Creative Commons Attribution-ShareAlike 3.0 Unported License (CC-BY-SA). You may redistribute it, verbatim or modified, providing that you comply with the terms of the CC-BY-SA