Electrochemical cells are devices based on the principle that when a chemical oxidation-reduction reaction takes place, electrons are transferred from one chemical species to another. In one type of electrochemical cell—called a voltaic or galvanic cell—these electrons are deliberately taken outside the cell and made to flow through an electric circuit to operate some kind of electrical device. A flashlight battery is an example of a voltaic electrochemical cell.
In the other type of electrochemical cell, called an electrolytic cell, the reverse process is takes place: Electrons in the form of an electric current are pumped through the chemicals in the cell to force an oxidation-reduction reaction to take place. An example of an electrolytic cell is the setup used to decompose water into hydrogen and oxygen by electrolysis.
Thus, a voltaic cell produces electricity from a chemical reaction, while an electrolytic cell produces a chemical reaction from electricity. Voltaic and electrolytic cells are considered separately below, following a general discussion of the relationship between chemistry and electricity.
In order to understand the intimate relationship between chemical reactions and electricity, we can consider a very simple oxidation-reduction reaction: the spontaneous reaction between a sodium atom and a chlorine atom to form sodium chloride:
What happens in this reaction is that an electron is passed from the sodium atom to the chlorine atom, leaving the sodium atom positively charged and the chlorine atom negatively charged. (Under normal conditions, the chlorine atoms are paired up into diatomic chlorine molecules, Cl2 ; but that does not change the present argument.) When a large number of sodium atoms and chlorine atoms are mixed together and react, a large number of electrons move from sodium atoms to chlorine atoms. These moving electrons constitute a flow of electricity. The “push” or potential for this electron flow comes from the sodium atoms’ eagerness to get rid of electrons and the chlorine atoms’ relative eagerness to grab them.
The practical problem when large numbers of sodium and chlorine atoms react is that the electrons are flowing in every direction—wherever a sodium atom can find a chlorine atom. We therefore cannot harness the electron flow to do useful electrical work. To use the electricity to light up a bulb, for example, we must make the electrons flow in a single direction through a wire; then we can put a bulb in their path and they will have to push through the filament to get from the sodium atoms to the chlorine atoms, lighting the filament up in the process. In other words, we must separate the sodium atoms from the chlorine atoms, so that they can only transfer their electrons on our terms: through the wire that we provide. Such an arrangement constitutes a voltaic or galvanic cell. It has the effect of converting chemical potential energy—a chemical push—into electrical potential energy—an electrical push: in other words, a voltage.
The sodium-plus-chlorine reaction is difficult to use in practice, because chlorine is a gas and sodium is a highly reactive metal that is nasty to handle. But many other chemical reactions can be used to make voltaic cells for generating electricity. All that is needed is a reaction between a substance (atoms, molecules, or ions) that wants to give up electrons and a substance that wants to grab onto electrons: in other words, an oxidation-reduction reaction. Then it is just a matter of arranging the substances so that the passing of electrons from one to the other must take place through an external wire. Strictly speaking, the resulting devices are voltaic cells, but people generally call them batteries.
As an illustration of how a voltaic cell works, we can choose the metallic elements silver (Ag) and copper (Cu) with their respective ions in solution, Ag+ and Cu++. Because copper atoms are more eager to give up electrons than silver atoms are, the copper atoms will tend to force the Ag+ions to take them. Or to say it the other way, Ag+ions are more eager to grab electrons than Cu++ions are, so they will take them away from copper atoms to become neutral silver atoms. Thus, the spontaneous reaction that will take place when all four species are mixed together is
This equation says that a piece of copper metal dipped into a solution containing silver ions will dissolve and become copper ions, while at the same time silver ions “plate out” as metallic silver. (This isn’t how silverplating is done, however, because the silver comes out as a rough and nonadhering coating on the copper. The silverplating of dinnerware and jewelry is done in an electrolytic cell.) To make a useful voltaic cell out of the copper-silver system, we must put the Cu and Cu++ in one container, the Ag and Ag+in a separate container, and then connect them with a wire. Bars of copper and silver metal should be dipped into solutions of copper nitrate, Cu(NO3)2, and silver nitrate, AgNO3, respectively. A salt bridge should be added between the two containers. This is a tube filled with an electrolyte—a solution of an ionic salt such as potassium nitrate KNO3, which allows ions to flow through it. Without the salt bridge, electrons would tend to build up in the silver container and the reaction would stop because the negative charge has no place to go. The salt bridge allows the negative charge, this time in the form of NO3 <-b1.0001>- ions, to complete the circuit by crossing the bridge. Now the circuit is complete and the reaction can proceed, producing a steady flow of electrons through the wire and keeping
Anode— A positively charged electrode.
Cathode— A negatively charged electrode.
Electrode— A conductor, usually a piece of metal, used to lead electricity (electrons) into or out of a region.
Electrolysis— The process by which an electrical current is used to break a compound apart into its components.
Electrolyte— The chemical solution in which an electric current is carried by the movement and discharge of ions.
Oxidation— The process in which an atom’s oxidation state is increased, by its losing one or more electrons.
Oxidation state or oxidation number— A positive or negative whole number that expresses how many units of combining power an atom is exhibiting toward other atoms. For example, sodium in NaCl has an oxidation number of→+1, while sulfur in Na2 S has an oxidation number of 2.
Oxidation-reduction reaction— A chemical reaction in which one or more atoms are oxidized, while one or more other atoms are reduced.
Reduction— The process by which an atom’s oxidation state is decreased, by its gaining one or more electrons.
the bulb lit until something runs out—either the copper bar is all dissolved or the silver ions are all depleted: our “battery” isdead.
In principle, a voltaic cell can be made from the four constituents of any oxidation-reduction reaction: any two pairs of oxidizable and reducible atoms, ions, or molecules. For example, any two elements and their respective ions can be made into a voltaic cell. Examples: Ag/Ag+ with Cu/Cu++ (as above), or Cu/Cu++ with Zn/Zn++(zinc), or H2/H+ (hydrogen) with Fe/Fe+++ (iron), or Ni/Ni++ (nickel) with Cd/Cd++(cadmium). The last cell is the basis for the rechargeable nickel-cadmium (nicad) batteries that power many electrical devices from razors to computers. When voltaic cells are used for portable purposes, they are “dry cells”: instead of a liquid solution, they contain a nonspillable paste. The lead storage battery in automobiles, however, does contain a liquid: a sulfuric acid solution.
There are many chemical reactions that, unlike the sodium-chlorine and copper-silver reactions above, simply will not occur spontaneously. One example is the breakup of water into hydrogen and oxygen:
This will not happen all by itself (that is, without the added energy) because water is an extremely stable compound. We can force this reaction, however, by pumping energy into the water in the form of an electric current. When we do this—pass an electric current through a chemical system to drive chemical reactionsmdash;it creates an electrolytic cell.
Electrolytic cells are used for a variety of purposes other than electrolyzing water. They are used to obtain metals such as sodium, magnesium, and aluminum from their compounds; to refine copper; to produce important industrial chemicals such as sodium hydroxide, chlorine, and hydrogen, and to electroplate metals such as silver, gold, nickel, and chromium onto jewelry, tableware, and industrial machine parts.
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Oxtoby, David W., et al. The Principles of Modern Chemistry. 5th ed. Pacific Grove, CA: Brooks/Cole, 2002. Umland, Jean B. General Chemistry. St. Paul: West, 1993.
Fun Science Gallery. “Experiments in Eletrochemistry”
<http://www.funsci.com/fun3_en/electro/electro.htm> (accessed November 10, 2006).
Georgia State University. “Electrochemical Cells” <http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/electrochem.html> (accessed November 10, 2006).
Robert L. Wolke
"Cell, electrochemical." The Gale Encyclopedia of Science. . Encyclopedia.com. (January 17, 2019). https://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/cell-electrochemical
"Cell, electrochemical." The Gale Encyclopedia of Science. . Retrieved January 17, 2019 from Encyclopedia.com: https://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/cell-electrochemical