Batteries (physics)

Type of physical science: Chemistry

Field of study: Chemical reactions

Batteries are devices that both store energy as chemical energy and convert this energy to electrical energy, which can be used to do work. The energy conversion occurs through chemical reactions.

Overview

Batteries are devices employed to convert chemical energy, which is stored within the battery, into electrical energy, which can then be used to perform work. The energy conversion occurs through a particular class of chemical reactions. The focus of this discussion is on the properties of these reactions and their production of electrical power.

Energy is stored in chemical bonds that link atoms together to form molecules, and also in the forces of attraction between electrons and the nucleus of atoms. Electrons are elementary particles that possess a unit negative electrical charge. Some of the energy in chemical species can be converted to a flow of electrons by way of chemical reactions. The source of the electrons is the chemical species. Such chemical reactions take place in batteries at interfaces, each of which separates a solution from what is called an electrode (reaction can also take place within porous electrodes). An electrode is made of a substance that is a good conductor of electricity, usually a metal. The electrodes of a battery connect to wires at terminals.

Both the chemical reactions and the electrodes associated with batteries can be divided into two types. One type of reaction causes electrons to be transferred from the interface into the material of the electrode. Electrodes at which these reactions occur are called anodes. The electrons travel from the anode to a wire, connected at a terminal. The wire, which is usually regarded as separate from the battery, is in turn usually connected to an electrical circuit.

Electrons move throughout the circuit and then return to the battery through a wire, which is connected, via another terminal, to a different electrode; this electrode is called a cathode. At the interface of a solution and a cathode, the second type of reaction occurs, in which electrons are transferred from the cathode to chemical species. The solution in the battery contains ions, which are molecules or atoms that have either gained or lost electrons. Ions are thus electrically charged particles, and the charge they possess corresponds to the number of electrons lost (positively charged) or gained (negatively charged). Positively charged ions are called cations, and negatively charged ions are called anions. A solution containing ions is called an electrolyte, and is necessary in a battery in order that electrical charge be transferred from one type of electrode to the other through the movement of the ions. Thus, the movement of ions completes the circular flow of charge.

One example of a battery is the lead storage battery used in automobiles. The electrodes in the lead storage battery are immersed in a solution containing water and sulfuric acid. In an aqueous solution, each sulfuric acid molecule breaks up into three ions: two hydrogen cations (which are protons), each of which possesses one positive charge, and a sulfate ion, which consists of one sulfur and four oxygen atoms. Each sulfate ion has two negative charges.

The anode is made of solid lead. In the chemical reaction that occurs at the surface of the anode, sulfate ions from the solution combine with the lead atoms to produce a compound called lead sulfate, a neutral molecule that consists of one lead atom, one sulfur atom, and four oxygen atoms. The lead sulfate is a solid that remains stuck to the anode surface. During this reaction, two electrons are transferred from each sulfate ion that reacts at the anode. The electrons then travel through an external circuit and return to the battery at the cathode. The cathode is made of lead (IV) oxide, a material in which each molecule consists of one lead atom and two oxygen atoms. At the surface of the cathode, the lead oxide reacts with ions in the solution to produce solid lead sulfate and water. Electrons are transferred from the cathode to the chemical species during the reaction.

In order to discuss the driving force for the production of electricity in chemical reactions, one must first examine some reactions of batteries in situations decoupled from the battery: Consider a jar partially filled with a solution containing water and copper sulfate, a neutral molecule consisting of one copper atom, four oxygen atoms, and one sulfur atom. If the amount of copper sulfate is small enough, all of it will dissociate into ions; each molecule breaks up into one divalent copper cation and one sulfate ion ("divalent" is used to denote that the copper ion has two positive charges). The solution will appear blue, characteristic of solutions containing divalent copper ions. If a piece of zinc metal is placed in the solution, the zinc dissolves. During the dissolution process, zinc atoms are converted to zinc ions in solution. Each zinc ion has two positive charges. But what happened to the two electrons that were associated with a given zinc atom before it dissolved? At the same time that zinc is dissolving, copper atoms form on the surface of the zinc metal. During the process, electron pairs are transferred from zinc atoms to the copper ions, transforming the copper ions to copper atoms. The atomic structure of copper is such that, given a choice, a pair of electrons prefer to be located in a copper atom rather than a zinc atom. This process is evidenced by the fading of the blue color that denotes the presence of divalent copper ions.

The same reactions can be made to occur in a system in which copper and copper ions are physically separated from the jar containing zinc and zinc ions. The bar of zinc metal is placed in one jar, all the copper ions in a second jar. Both jars will have a common electrolyte that contains sulfate ions. The second jar also needs a metal, but instead of zinc a bar of copper metal is used. The same net process will occur as the one that occurs using the one-jar system if the zinc and copper bars are connected with a wire. The zinc atoms are transformed to zinc ions, and two electrons are produced at the interface for each atom transformed. At the bar of copper metal, copper ions are transformed to copper atoms. Without the wire, these two processes would stop after a relatively brief period of time. Since electrons are produced at the zinc bar and consumed at the copper bar, the copper bar will appear to the electrons at the zinc bar as more positive than the zinc. Electrons always try to move toward the greatest positive charge, thus they move from the zinc bar, through the wire, to the copper bar. The electrons that are accepted by the copper ions at the copper bar can be thought of as those produced at the zinc bar.

Both the movement of electrons and the chemical reactions will come to a halt unless a charge balance is maintained. The requirement can be maintained throughout the whole system by connecting the jars with what is called a salt bridge: a bent tube filled with a salt. Anions, in this case sulfate ions, travel from the copper jar to the zinc jar through the salt bridge. The negative charge they carry keeps the whole system's charge balanced. Otherwise, the jar with the zinc bar would become positively charged as a result of the production of the positively charged zinc ions, and the jar containing the copper bar would become negatively charged as a result of the loss of the positively charged copper ions.

The original jar containing both zinc and copper and their ions supported the reaction, which is represented by the equation: one zinc atom + one divalent copper cation – one copper atom + one divalent zinc cation which represents the fact that one zinc atom and one copper ion are transformed by the reaction to one copper atom and one zinc ion.

In the two-jar arrangement, the overall reaction is split into two reactions. Each of these reactions is called a half-cell reaction. In the jar containing the zinc bar, the following reaction occurs: one zinc atom + one divalent zinc cation + two electrons that is, one zinc atom is transformed to a zinc ion and two electrons are produced. The reaction that occurs in the jar containing the copper bar is one divalent copper cation + two electrons + one copper atom

That is, two electrons are transferred to a copper ion, and the transfer produces a copper atom.

In the two-jar arrangement, the electrons produced in the jar containing zinc are consumed in the jar containing copper. Thus, the overall chemical process is the same as in the one-jar arrangement; in other words, adding equations 2 and 3 gives equation 1. In the two-jar arrangement, however, electrons flow through a wire, thus producing a current, and this wire can be connected to a circuit in which the flow of electrons can be utilized to do work.

The two-jar arrangement, including the coupling with the wire and salt bridge, is called by several names: electrochemical cell, galvanic cell, voltaic cell, or battery. Batteries may contain several of the two-jar-type arrangement in one unit. The basic principles involved in producing a current in the two-jar arrangement are the same as those behind the construction of commercial batteries; commercial batteries may include several cells linked in a series.

Chemists have learned to decouple chemically reactions like those in equation 1 and couple them electrically to make batteries. The reaction in equation 1 is a specific example of a class of reactions called oxidation-reduction reactions. Oxidation-reduction reactions are reactions in which electrons are transferred between different molecules, and/or atoms, and/or ions. In these types of reactions, a species that receives electrons is said to be reduced; this is easy to remember because the charge of the species changes in the negative direction. In the case of the zinc-copper reaction, the copper ion is reduced; it receives two electrons and is transformed into a copper atom. A species that loses electrons during the reaction is said to be oxidized. In the zinc-copper reaction, a zinc atom is oxidized; it loses two electrons and is transformed into a divalent cation. (As a historical note, the term "oxidation" derives from the word "oxygen," and was first applied when it was realized that metals that reacted with oxygen lost electrons.) Also in the above example, the copper bar is the cathode, and the cathode is always the site of the reduction reaction, usually marked positive at the associated terminal in a battery. The zinc bar is the anode, and the anode is always the site of the oxidation reaction, marked negative at the terminal.

In principle, any oxidation-reduction reaction can be coupled electrically and decoupled chemically so that the reduction occurs in one jar and the oxidation occurs in a second jar. As a second example, solid cobalt will react with divalent nickel cations to produce divalent cobalt cations and solid nickel; cobalt is oxidized and nickel is reduced. The reaction may be divided into two half-cell reactions by placing a bar of cobalt metal in one jar containing an electrolyte, and a bar of nickel metal in a second jar containing the same electrolyte and, in addition, divalent nickel cations. Coupled electrically, electrons will be produced at the cobalt metal, travel through the wire, and then be consumed at the nickel metal bar.

If in one oxidation-reduction reaction the ion from an ion-atom couple is reduced, this does not imply that the ion will be reduced in all oxidation-reduction reactions. For example, if solid copper is placed in a solution of silver nitrate, copper atoms react with monovalent silver cations to produce divalent copper ions and silver atoms. Here, copper is oxidized, whereas in the zinc-copper reaction, copper was formed by the reduction of copper ions.

When oxidation-reduction reactions are chemically decoupled into two half-cell reactions, one can measure what is called the potential energy difference for the electrons between the two chemical environments, that is, between the two jars. This potential is also called voltage and can be measured by hooking up a voltmeter to the wire that connects the two metal bars. The higher the voltage (measured in volts), the greater the driving force that pushes the electrons from one jar to the other. The voltage that is obtained depends on the chemical composition in the two jars and the concentration of the chemicals. In making batteries, one usually desires to have large voltage.

Applications

Electrolytes of batteries can be a solid, liquid, or even a gas. The fact that the electrolyte does not have to be a runny liquid is demonstrated by the dry cell battery which has commonly been used in flashlights. Here, the electrolyte is a moist paste that contains ammonium ions. (An ammonium ion consists of one nitrogen and four hydrogen atoms, and possesses one positive charge.) The paste fills a zinc container, which is lined with a porous separator; the separator allows some ions to pass through it and plays the same role as the salt bridge in the two-jar system. The zinc container is the anode. A carbon rod is inserted into the paste and serves as the cathode. At the cathode, ammonium ions are reduced, that is, they accept electrons, and are converted into ammonia gas and hydrogen gas. At the anode, zinc ions and electrons are produced. The dry cell also provides an example of a case in which reactions at the electrodes must be coupled to other chemical reactions. Since the products of the cathode reaction are gases, they would cause a sealed dry cell to explode; however, the paste electrolyte contains quantities that react with these gases. The voltage produced by a typical dry cell battery is 1.5 volts.

A battery that is more expensive than the dry cell but more efficient is the alkaline cell.

Here, the anode is made of a paste composed of zinc particles, potassium hydroxide, and water.

The potassium hydroxide dissociates into monovalent potassium cations and hydroxyl anions.

Each hydroxyl anion consists of a hydrogen atom and an oxygen atom; if a hydrogen cation is chemically combined with the hydroxyl anion, a water molecule is formed. It is because of the presence of hydroxyl ions that the battery is called an alkaline battery; mixtures that contain more hydroxyl ions than hydrogen cations are called alkaline. The paste is enclosed in a plastic sleeve and surrounds a brass cylinder, which collects electrons produced by the reactions in the anode. The sleeve separates the anode from the cathode, which consists of a paste containing graphite particles (a form of carbon that is used in pencils), water, and magnesium (IV) oxide.

The zinc in the anode of the alkaline battery reacts with the hydroxyl anions to produce a zinc oxide, water, and electrons. The electrons are consumed at the cathode by the manganese oxide during a reaction with water. An alkaline battery typically produces a voltage of 1.54 volts.

The dry cell and alkaline batteries are examples of primary batteries. In this type of battery, a chemical equilibrium is eventually reached, at which point the battery fails to produce a current. The dead battery is thrown away. On the other hand, when equilibrium is approached in the lead storage battery, the battery can be recharged.

The fact that the lead battery is rechargeable is a consequence of an important principle of oxidation-reduction reactions. Consider again reactions in a two-jar arrangement. The reactions, such as those in the zinc-copper system, are spontaneous: They are naturally driven in the direction that they occur. A measure of the driving force is the measured potential. If, instead of measuring the potential, a potential is applied, the reaction can be reversed. For example, in the zinc-copper system, it is possible in the two-jar arrangement to make the copper dissolve and zinc ions react to form zinc atoms. To do this, one must apply a potential with a sign opposite to and a magnitude greater than the potential measured for the spontaneous reactions. A lead battery is recharged in the same way. The fact that the product of the spontaneous reactions in the lead storage battery, lead sulfate, forms on the electrode is also important. If the lead sulfate dissolved, then it would have to diffuse back to the electrodes during recharging. This would take a very long time, and it would make the battery useless for operation in an automobile.

Context

The first battery was constructed at the end of the eighteenth century by Alessandro Volta, an Italian physicist. It consisted of alternating silver and zinc disks, separated by cloths saturated with salt solutions, and it was, and still is, commonly called a pile. This and several other piles invented at the beginning of the nineteenth century were examples of primary batteries. During this time and for the remainder of the nineteenth century, much research effort was put into the invention and development of storage batteries. That the lead storage battery emerged out of nineteenth-century science was largely attributable to telegrams. The field of telegraphy was a driving force for the invention of the lead storage battery; although primary batteries served as adequate power sources up until about 1840, they proved to be insufficient for the lengthening telegraph lines and sophisticated instruments. Research on batteries with rechargeable cells laid the foundation for the successful storage battery industry of the 1880's, even though the primary goal of fulfilling the needs of telegraphy was not achieved. Several discoveries led up to the development of storage batteries. One of them was made by a German chemist by the name of Johann Ritter. Ritter seems to have been the first, in 1802, to have obtained solids on electrodes during the production of current in piles. Although this may seem to be an event of small proportions, previous to Ritter's discovery only the formation of gases had been observed in the operation of piles. As was noted, it is the formation of lead sulfate on the electrodes of the lead storage battery that allows it to be recharged. Technically, Ritter was the first to construct a storage battery, but it was not until approximately twenty years later that further developments occurred. One of the difficulties at the beginning of the eighteenth century was the lack of electrical measuring devices; Ritter used the response of frog's muscle to estimate the current produced.

One question that had to be answered before progress could be made on storage batteries was, what was causing the electrical current to flow in piles? Many theories existed, among them the idea that it was caused by chemical reactions. By connecting different combinations of batteries, Sir Humphry Davy discharged and recharged cells so that he kept changing the direction of the current between electrodes of these cells. He then was able to relate the change in current with the chemical changes on the electrodes. These experiments provided strong evidence in favor of the chemical reaction theory.

After it was finally accepted that chemical reactions produced electricity in batteries, chemists realized that it was an advantage to place strongly electron-accepting materials (species that are easily reduced) on the cathode. These substances made the current of primary batteries large and constant. It was a German chemist, P. S. Muncke, who, in 1835, discovered that lead (IV) oxide (Pb02) was an excellent material for the cathode; it was easily reduced and was a good conductor. It was Auguste de la Rive, however, who first designed a battery using Pb02.

The chemist Gaston Plante spent thirty years, from 1859 to 1889, developing the design of the lead storage battery; he was not, as is suggested in some historical accounts, the original inventor. One of Plante's important contributions was to show that lead and lead oxide will react spontaneously with sulfuric acid to form lead sulfate. Then, when the reaction was reversed, one would obtain more lead oxide and less lead. Until this was known, lead batteries possessed cathodes composed of lead covered with a thin layer of lead oxide. Lead is nonporous, whereas lead oxide is porous. Because of Plante's discovery, pure lead oxide cathodes were developed in which the sulfuric acid could penetrate and allow more molecules to react, and which increased the energy storage capacity of the battery.

The early years of the twentieth century marked the final stages of the development of the nickel-cadmium storage battery by Waldemar Jungner, a Swedish chemist, and the development of the nickel-iron battery by Thomas Edison. These are examples of storage batteries that have alkaline electrolytes. Both batteries are used in a variety of appliances.

Much research continues to be devoted to improving electrodes, electroytes, and the understanding of the fundamentals of electrochemical reactions. Theory predicts that large improvements can still be made in the efficiency of the electrochemical conversion of chemical energy into electrical energy. The electrochemical conversion should be more efficient than other methods, as it does not use an intermediate step in which heat is evolved, as is the case in combustion.

Principal terms

ANODE: an electrode at which an oxidation reaction occurs

CATHODE: an electrode at which reduction occurs

CHARGE: a measure of electricity; an electron has a charge of 1.6019-19 coulombs

ELECTRODE: a material used for conducting electrons into and out of solutions

ELECTROLYTE: a solution that contains ions

IONS: atoms or molecules that have lost or gained electrons so that they are no longer electrically neutral

LEAD (IV) OXIDE: a neutral molecule consisting of one lead atom and two oxygen atoms; the roman numeral IV is used to denote the fact that the electron density around the lead atom has, on average, four fewer electrons than the lead atom does in the state of pure lead

OXIDATION: a loss of electrons by an atom, molecule, or ion

REDUCTION: a gain of electrons by an atom, molecule, or ion

Bibliography

Graham, Robert W. PRIMARY BATTERIES: RECENT ADVANCES. Park Ridge, N.J.: Noyes Data Corporation, 1978. Reviews the patent literature on primary batteries between 1975 and 1978. The legal jargon as well as much of the highly formal technical language has been removed, leaving an informative description of the U.S. patent literature. The topics include zinc dry cell batteries, batteries for pacemakers, watch battery design, and seawater batteries.

Jasinski, Raymond. HIGH-ENERGY BATTERIES. New York: Plenum Press, 1967. A summary of the basic principles and technology of battery performance. Although somewhat technical, the book contains many descriptive parts describing theory and applications. This book also puts forth a good argument for continued research in the area of batteries, based mostly on the prediction of theory that there remains unrealized potential.

Kotz, John C., and Keith F. Purcell. CHEMISTRY AND CHEMICAL REACTIVITY. Philadelphia: Saunders College Publishing, 1987. One of a large number of freshman chemistry texts that present a chapter (chapter 21) on oxidation-reduction reactions including sections on batteries. This particular text is loaded with large, colorful photographs and illustrations that, in themselves, provide interesting descriptions of several different batteries as well as the related chemical reactions.

Ostwald, Friedrich Wilhelm. ELECTROCHEMISTRY: HISTORY AND THEORY. 2 vols. Translated by N. P. Date. New Delhi: Amerind, 1980. Translation of a classic book, originally published in 1896. The author, who won a Nobel Prize in 1909, possessed a remarkable personality, which is captured by his writings. The book gives a detailed history of electrochemistry, including advances in battery technology. Ostwald describes the controversies over scientific views that now are considered pillars of scientific theory.

Schallenberg, Richard H. BOTTLED ENERGY. Philadelphia: American Philosophical Society, 1982. Traces the history of both the research and development of storage batteries from 1800 to around 1950. The book describes both the scientific developments and economic driving forces for these developments at a nontechnical level. Much can be learned from this book regarding the personalities of the scientists and inventors, as well as of the competition among them.

Volta's pile