Ox-redox Reactions

Type of physical science: Chemistry

Field of study: Chemical reactions

Chemical reactions in which one or more electrons are transferred from one substance to another are termed oxidation-reduction, or redox, reactions. Redox reactions are utilized for the commercial production of many chemical compounds and in batteries.

Overview

Atoms are composed of electrically positive, negative, and neutral particles, called protons, neutrons, and electrons, respectively. Chemical processes in which electrons are transferred from one chemical substance to another are representative of a major chemical reaction class called oxidation-reduction reactions, or, more commonly, redox reactions.

Oxidation is the loss of electrons in a chemical process, while reduction is a gain of electrons. In a redox reaction, electrons are transferred from the substance being oxidized to the substance being reduced. The material being reduced pulls electrons from another material, and is thus termed an oxidizing agent. A reducing agent causes another material to gain electrons and is itself oxidized. Any redox reaction is characterized by the presence of both an oxidizing and a reducing agent.

A common example of a spontaneous redox reaction is the corrosion of iron. When elemental iron is exposed to oxygen in the presence of water, the iron is oxidized, losing two electrons, resulting in the formation of common rust. Corrosion can also occur when two dissimilar metals are placed in contact with each other.

In order to understand any chemical reaction, it is necessary to determine the ratio in which the starting materials, or reactants, combine, as well as the ratio of reactants to products.

This information is conveyed using a shorthand notation for the chemical process called a chemical equation. To be useful, the equation must be balanced. In other words, if there are five atoms of element A present in the reactants, there must be five atoms of element A in the products. With redox processes, there is another factor to consider: The total number of electrons lost in oxidation processes in the reaction must be the same as the number of electrons gained in reduction processes. This simply means that the electron transfer must be balanced.

For some reactions, the number of electrons transferred may be easily determined, as in the case of rusting iron mentioned earlier. Rust is an example of an ionic compound, in which positively charged atoms, or cations, are combined with negatively charged atoms, or anions, to form a crystalline solid, which is held together by the electrostatic interactions of the opposing charges. Cations and anions in ionic compounds have characteristic charges. For example, sodium and chlorine form charged atoms, or ions, with charges of positive one and negative one, respectively. Formation of sodium chloride, or table salt, involves the transfer of one electron from a sodium atom to a chlorine atom.

In general, for reactions involving the formation or decomposition of ionic compounds, the number of electrons transferred is easy to determine, as change in the electrical charges of the atoms in the compound indicates the number of electrons lost or gained by the atoms. Many compounds utilize another type of bonding, covalent bonding, in which electrons are not transferred but shared between two atoms in the bond. Since atoms in this type of compound do not have a charge, it is difficult to view clearly redox reactions of covalently bound compounds as electron transfer processes.

In such cases, oxidation numbers must be utilized. Oxidation numbers are determined according to a specific set of rules and assigned to the atoms in a compound. They represent an assigned charge and are used specifically to balance redox reactions. In the case of ionic compounds, the oxidation number is simply the charge on the cation or anion. For free elements, such as metallic iron, the oxidation state is assumed to be zero. Many elements in compounds have characteristic oxidation numbers that they are assigned in virtually all compounds. These characteristic numbers typically can be determined from the location of an element in the periodic table. All assigned oxidation states must add up to the net charge of the species. For example, consider chlorine trifluoride, a neutral molecule in which three fluorine atoms are attached to a central chlorine atom. Fluorine is assigned an oxidation number of negative one in all its compounds. In order for the oxidation numbers to add up to zero overall for the neutral molecule, the chlorine atom must have an oxidation number of positive three.

In redox processes, the oxidation states of elements are changed in the conversion of reactants to products. Oxidation and reduction are thus more generally defined in terms of oxidation number changes, which take place most commonly via electron, oxygen atom, or hydrogen atom transfer. Increasing the oxidation number of an element is oxidation, while in a reduction process, the oxidation number decreases. A properly balanced redox reaction is one in which the total increase in oxidation number of species oxidized is equal in magnitude to the total decrease in oxidation number of species being reduced.

There are two techniques for balancing electron transfer reactions, both of which require assignment of oxidation numbers to all atoms. This allows determination of the elements being oxidized or reduced and the number of electrons involved in each process. In the oxidation number method, first the oxidation number changes are balanced by using simple ratios. Then the remaining components of the reaction are balanced based on the redox ratio. To utilize the method of half-reactions, the reaction is divided into two parts, an oxidation half and a reduction half. Each half-reaction is balanced independently, with appropriate numbers of electrons added to each half-reaction to balance the electrical charges. Each half-reaction is then multiplied by integers, so that the number of electrons lost in the oxidation half-reaction equals the number of electrons gained in the reduction half-reaction. These half-reactions are then recombined to yield the balanced net reaction. The half-reaction method is particularly useful when one is considering the reactions that take place at the electrodes in batteries, which utilize redox reactions to produce electricity.

In order to write a balanced chemical equation, one needs to have knowledge of the reaction products. While these can be determined by experimentation, it is desirable to predict the products of a particular reaction before it is run. In the case of redox reactions, this prediction can be made on the basis of relative tendencies of substances in the reaction to gain or lose electrons. As an example, zinc metal will displace dissolved copper from water. In this process, the zinc loses two electrons, while the dissolved copper gains two electrons and comes out of the solution in its metallic form. A substance that displaces zinc from solution will also displace copper. A series of metals listed in order of increasing reducing strength constitutes an activity series. Such a series gives an idea of the relative strengths of various oxidizing and reducing agents and allows one to predict in which direction a reaction will probably proceed.

Applications

Redox reactions have been utilized by humankind since prehistoric times. Combustion, the reaction of oxygen with another chemical substance, is probably the single most important and widely utilized example of a redox process. A combustion reaction is characterized by a flame, which liberates both heat and light. Oxygen acts as the oxidizing agent; historically, this reactivity is the derivation of the term "oxidation." Man utilizes combustion reactions to heat homes, power automobiles and other types of machinery, generate electrical power, and for a bewildering variety of other applications.

Naturally occurring redox processes are important in nature, in both living and nonliving systems. Many biological processes in plants and animals involve redox reactions.

Plants convert sunlight into chemical energy in photosynthesis, in which the carbon in carbon dioxide is reduced to a lower oxidation state to form carbohydrates, and the oxygen is oxidized to its electrically neutral, elemental state. This stored energy is released in the respiration process, in which the organic fuel is oxidized by molecular oxygen. Several steps, which involve a series of enzymes, or biological catalysts, are involved in the transfer of electrons from the fuel to, ultimately, oxygen. This sequence of steps is termed oxidative phosphorylation. An interesting aspect of this sequence of electron transfer reactions is that each successive step involves a slightly stronger oxidizing agent, culminating with oxygen, a very strong oxidizer.

This great tendency of elemental oxygen to cause oxidation is manifested in many ways in nature. Most naturally occurring materials tend to exist in their oxidized forms as a result of prolonged contact with the highly oxidizing atmosphere of the earth. In particular, most metallic elements are found in nature not in the elemental, metallic state, but rather as oxidized ions in minerals. The minerals, or ores, containing metallic elements are actually ionic compounds in which the metal exists as a cation. In order to isolate the metal from the ore in its pure, metallic state, the metal ion in the ore must be chemically reduced somehow. Various techniques are used to reduce metallic ores, with the technique and conditions utilized dependent on the metal to be isolated. These techniques are collectively a part of modern technology known as extractive metallurgy, which is simply the process of extracting a metal from its natural ore.

The name is somewhat misleading, as the metal is not simply extracted unchanged from the ore but is chemically altered by the addition of electrons in the reduction process.

Probably the most widely known method of extractive metallurgy is the reduction of the ore by carbon or carbon monoxide. Iron, which is the fourth most common element in the earth's crust, must be extracted from its ores by such a reductive process. To produce iron, the ore, which is typically an oxide such as hematite, is placed in a blast furnace along with limestone and coke, which both serve as a fuel to generate very high temperatures in the furnace and produce carbon to act as a reducing agent. The reduction of iron to its elemental state takes place in the furnace at a temperature of about 2,000 degrees Celsius. Many other common metals, such as tin and lead, may also be obtained by reduction of oxide ores by carbon. These other metals typically do not require the extremely high temperature for reduction that is needed for iron.

While this carbon reduction technique works well for a variety of metals, there are others that cannot readily be reduced in this manner. In such cases, it is possible to utilize electricity to reduce the metallic ion to its neutral state. Such a method was developed by Charles Martin Hall in 1889 for the production of aluminum from cryolite. The ore is placed in a large vessel lined with iron electrodes. Graphite electrodes are immersed in the cryolite/alumina melt to reduce the aluminum. Application of electricity is also utilized to produce metallic sodium and chlorine gas from molten table salt, sodium chloride. In this process, the sodium ions are reduced and the chloride ions are oxidized, resulting in formation of the free elements. When sodium chloride is dissolved in water and an electric potential applied, different products, including sodium hydroxide, hydrogen gas, and chlorine gas are produced. This process can even be used to produce the hypochlorite ion, the active ingredient in chlorine bleach, when the mixture is stirred. These reactions are examples of electrolysis processes, or cleavage by electricity.

Electrolysis reactions cause electron transfer reactions to take place that do not occur naturally.

A related use of redox reactions induced by the application of a potential is in electrorefining of metals such as gold or copper. Under suitable conditions, these metals can be plated onto an electrode in an extremely pure state. Electroplating, which is used to place a thin layer of pure metal on the surface of an object, is another example of a redox process induced by application of an electrical potential.

Most of the applications mentioned previously utilize extreme conditions such as high temperature, or the application of an electrical potential, to induce a redox process. Spontaneous redox reactions, reactions that occur naturally without the application of any external driving force, are also extremely important industrially. Probably the most common and important use of spontaneous redox chemistry in the commercial sector is in batteries. Electrical current is caused by the movement of electrons through a wire. Redox processes involve transfer of electrons and can thus be utilized to generate an electric current. This is accomplished by forcing the oxidation and reduction half-reactions to take place in separate compartments, or cells. These cells are connected in such a way that the transferred electrons are constrained to flow through a wire connecting them to generate a current. From tiny watch batteries to fuel cells utilized in modern spacecraft, batteries are easily transported and versatile sources of chemical energy.

Most of the redox reactions that take place in batteries may be reversed to regenerate starting materials by application of an electrical potential to the battery. Secondary, or rechargeable, batteries, such as the lead storage battery used in automobiles, are designed to facilitate this reversal so that the battery may be recharged and reused.

Corrosion of metals, especially iron, is a naturally occurring process. Redox chemistry may be utilized to prevent this process and protect metallic structures in the environment. This is done by attaching the metal that is to be protected to another metal that is more easily oxidized, called a sacrificial anode. The sacrificial anode is oxidized preferentially, thus protecting the other metal from corrosion. This technique is used to protect buried pipelines, as well as to preserve the metal in offshore oil platforms from the effects of the highly corrosive environment at sea.

Context

Modern concepts of redox chemistry have ultimately been derived from the first reaction of this type to be studied, combustion. So fundamentally important was this reaction to man in recorded history that fire was classified along with earth, air, and water as an element by Empedocles.

What is now known as redox chemistry was utilized prehistorically the first time a metal, probably copper, was obtained by reduction from its ore. Extraction of tin led to the Bronze Age, and development by the Hittites in about 1200 B.C. of a method for extracting iron from its ores opened the door to the iron age and modern society. By the 1800's, these techniques of extractive metallurgy had reached a high degree of sophistication, and many cities grew from this industry. Much of the technique in metallurgy was developed by trial and error, with little real understanding of the chemical processes involved.

Discovery of oxygen in the eighteenth century and subsequent study of its chemical reactions provided greater understanding of this type of reactivity. Perhaps the most important work contributing to modern understanding of redox reactions as electron transfer began in the late eighteenth century, when Luigi Galvani discovered animal electricity. To study this phenomenon, Alessandro Volta placed two dissimilar metals in a circuit to cause animal convulsions. This apparatus was the first battery, or voltaic cell, which utilized a spontaneous chemical reaction to generate electricity.

These simple beginnings sparked a dramatic burst of discovery and technological development in the following years that included the invention of telegraph, current for which was provided by a spontaneous redox reaction. With the discovery that electrochemical processes could cause chemical changes that were very similar to those caused by oxygen, oxidation took on a broader meaning. With the discovery of the electron in 1911, redox processes became linked with the concept of electron transfer. This basic idea is still predominant in modern chemistry, although refined somewhat to explain redox reactions of covalently bound molecules, in which the electrical charges of chemical elements involved do not change substantially.

Many modern devices that are taken for granted rely on redox reactions to function, including any battery-operated device. Oxidation-reduction chemistry is very much a part of basic research in modern chemistry. Electron transfer chemistry is an integral part of virtually all chemistry textbooks and a major topic in many chemistry courses. The importance of redox chemistry is perhaps best illustrated by the awarding of the 1983 Nobel Prize in Chemistry to Henry Taube for his work on electron transfer mechanisms. As more and more specialty chemicals are developed, manipulation of oxidation states through redox processes may lead to significant advances in chemistry, biology, and material science.

Principal terms

ELECTRODE: the surface in an electrochemical cell at which oxidation or reduction takes place

ELECTRON: a subatomic particle that is the fundamental unit of negative charge

OXIDATION: loss of electron(s) in a chemical process, resulting in an increase in oxidation number

OXIDATION NUMBER: a number assigned to an atom according to a set of rules used to determine the number of electrons transferred in a chemical reaction

REDUCTION: gain of electron(s) in a chemical process, resulting in a decrease in oxidation number

Bibliography

Asimov, Isaac. ASIMOV'S CHRONOLOGY OF SCIENCE AND DISCOVERY. New York: Harper & Row, 1989. In this very readable and well organized work designed for the general reader, Asimov places scientific and technological accomplishments in their historical perspective. Included are several aspects of redox chemistry, for example, various types of batteries.

Basolo, Fred. "High Oxidation State Metals in Organometallic Chemistry." In 1990 YEARBOOK OF SCIENCE AND THE FUTURE, edited by David Calhoun. Chicago: Encyclopaedia Britannica, 1989. This narrative outlines the development of a relatively new class of compounds in which metal atoms are in uncharacteristically high oxidation states.

Davis, Edith. "A Revised Approach to Solving Redox Equations." JOURNAL OF CHEMICAL EDUCATION 67 (August, 1990): 671. This brief article presents a fresh approach to balancing redox equations.

Hall, Stephen S. "The Age of Electricity." In INVENTORS AND DISCOVERERS: CHANGING OUR WORLD. Washington, D.C.: National Geographic Society, 1988. This chapter describes the development of electricity, beginning with the early experiments of Volta and Galvani. Very well written and illustrated, with both photographs and paintings.

Labianca, Dominick A. "The Chemical Basis of the Breathalyzer: A Critical Analysis." JOURNAL OF CHEMICAL EDUCATION 67 (March, 1990): 259. This article deals with one of the most common and important everyday uses of a redox reaction, in the breath test for alcohol. Discusses the strengths and weaknesses of the method.

McQuarrie, Donald A., and Peter A. Rock. "Oxidation-Reduction Reactions." In GENERAL CHEMISTRY. 3d ed. New York: W. H. Freeman, 1991. In this chapter, the principles of redox chemistry are introduced, along with rules for assigning oxidation numbers and detailed methods of balancing redox equations. Other chapters in this readable, well-illustrated text discuss electrochemical applications and describe extractive metallurgy.

Electrolysis

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