Oxidation-Reduction Reaction

History

Oxidation numbers

Examples of oxidation-reduction reactions Combustion

Corrosion

Biological processes

Current and future uses

Resources

Oxidation-reduction reactions, also known as redox reactions, are chemical processes in which electrons are transferred from one atom, ion, or molecule to another. Explosions, fires, batteries, and even our own bodies are powered by oxidation-reduction reactions. When iron rusts or colored paper bleaches in sunlight, oxidation-reduction has taken place.

Oxidation-reduction reactions are a combination of two processes: oxidation, in which electrons are lost, and reduction, in which electrons are gained. The two processes cannot occur independently of each other. A mnemonic device used by chemists to help keep things straight is “LEO says ‘GER,”’ which stands for Loss of Electrons, Oxidation. Gain of Electrons, Reduction.

The driving force of oxidation-reduction reactions is the transfer of electrons. Although it is sometimes difficult to remember what happens to electrons during oxidation and what happens during reduction, a look at familiar processes can help keep this straight. Some of the first oxidation-reduction reactions understood by chemists were those that involved oxygen, the most plentiful element on Earth. When combined with other elements in a compound or molecule, oxygen frequently is an electron “hog,” taking electrons away from many other elements, which oxidizes them. The oxygen then takes the negatively charged electrons and becomes a negatively charged ion. The oxygen has been “reduced,” somewhat like taking in negative thoughts will reduce a person’s positive attitude. An example of this is the reaction between oxygen in the air and iron. The metal iron becomes positively charged and the oxygen becomes negatively charged. The two charged ions now attract each other and hang around together in the form of iron oxide, or rust.

History

Probably the earliest human use of oxidation-reduction reactions occurred 4,500-7,500 years ago in the Copper/Bronze Age. Copper ores were heated in the presence of carbon to produce copper metal. In this process, the copper in the ore was reduced to copper metal and the carbon was oxidized to carbon dioxide. This same process was applied to iron ores during the Iron Age, which occurred 3,500-4,500 years ago.

Oxidation-reduction reactions have long been a part of pottery making as well. Color differences in the clay or glaze can be produced when firing pottery under oxidizing conditions, when lots of oxygen is present, or under low-oxygen conditions, such as with a partially closed kiln or a fire with green leaves on it. Clay containing iron will be orange-red if fired under oxidizing conditions (due to the presence of red iron oxide) and black in reducing conditions, when black iron oxide—in which the iron has a lower oxidation number—forms. Among the people who have historically used oxidizing and reducing fire conditions are the Native Americans in the southwestern United States and the Greeks in the early Bronze Age.

Oxidation-reduction reactions are also used in explosives, substances that burn (oxidize) so rapidly that they cause huge amounts of pressure. Gunpowder, thought to be the first explosive used, was used in China as early as the sixth century to make fireworks, and by 960 AD for military applications; it had migrated to Europe around the thirteenth century. Nitrocellulose and nitroglycerin were developed in 1846 and 1847, respectively. TNT (trinitrotoluene), first developed in 1863, saw widespread use in World War I. Since 1955, a commonly used cheap and powerful explosive has been a mixture of ammonium nitrate and fuel oil. This was used in the in 1995 Oklahoma City bombing.

An important step in the understanding of oxidation-reduction reactions was the discovery of oxygen. Joseph Priestley (1733-1804) was the first scientist on record to prepare oxygen in the laboratory. This historic reaction was also an oxidation-reduction reaction. Priestley heated mercury oxide and formed elemental mercury and oxygen. In this reaction, mercury was reduced and the oxide ion was oxidized. Antoine Lavoisier (1743-1794) recognized that when substances are burned, they combine with oxygen. He even figured out that our bodies burn food and give off carbon dioxide as we produce energy. Tragically, the life of this great chemist was ended prematurely when he was beheaded during the French Revolution.

Oxidation numbers

Oxidation numbers, sometimes called oxidation states, help chemists keep track of the numbers of electrons that surround each atom in a chemical reaction, and how they change in oxidation-reduction reactions. When an atom gains an electron (is reduced), its oxidation number is increased by one. There are some simple rules for assigning oxidation numbers to elements in chemical compounds:

1. The oxidation number of an element, having neither gained nor lost any of its electrons, is zero. For example, the oxidation number of pure copper, Cu, is zero, as is the oxidation number of each oxygen atom in a molecule of oxygen, O₂.
2. The oxidation number of an elemental ion is the same as its charge. An ion of copper with a +2 charge, Cu²⁺, has an oxidation number of +2. A fluoride ion, F^—, has an oxidation number of —1.
3. Some elements almost always form compounds in which they have a particular oxidation number. Aluminum always forms a+3 ion and therefore exists in the+3 oxidation state in compounds. Sodium and other alkali metals almost always form a+1 ion; its oxidation state is+1. Hydrogen can form compounds in which the hydrogen atom has an oxidation number of either+ or— 1. When hydrogen has an oxidation number of+1, it is written on the left-hand side of the chemical formula. If its oxidation number is – 1, it is written on the right-hand side. Oxygen usually has a – 2 oxidation number. Chlorine and other halogens usually take on a–1 charge. Other elements are not so predictable. Nitrogen can have oxidation numbers of +5, +4, +3, +2, +1, and–3.
4. The sum of oxidation numbers in a neutral molecule or compound is zero. Table salt, with the chemical formula of sodium chloride, NaCl, is made up of two ions, a positively charged sodium ion and a negatively charged chloride ion. A water molecule consists of two hydrogen atoms, each with an oxidation number of+1, and an oxygen atom with an oxidation number of—2.

It is often easier to follow oxidation-reduction reactions if they are split into two half-reactions. One half reaction indicates what is happening to the chemical substances and electrons in the oxidation portion of the reaction. The other half-reaction does the same for the reduction portion. The complete reaction is the sum of the two half-reactions.

A useful tool for chemists is a table of standard reduction potentials. This table lists common half-reactions, and assigns each a numerical value that indicates how easily the reduction reaction proceeds—that is, how eagerly electrons are accepted. A high standard reduction potential value indicates that the substance is easily reduced. A low standard reduction potential indicates that the substance is easily oxidized—it prefers to lose electrons. In general, a substance will oxidize something that has a lower reduction potential than it has. The halogens, chemical elements found in group 17 of the periodic table, are strong oxidizing agents because their atoms readily accept negative ions. The alkali metals such as sodium, found on the left side of the periodic table in group 1, are strong reducing agents because their atoms readily give up an electron, becoming positive ions. The arbitrary zero point for standard reduction potentials has been designated as this reaction:

2H⁺ + 2 electrons→ H₂

This reaction has been assigned a potential of 0.000 volts under standard conditions. The standard reduction potential for fluorine gas is 2,890 volts, while that for sodium metal is—2.714 volts.

Examples of oxidation-reduction reactions Combustion

Let us look at an oxidation-reduction chemically as we examine what happened to the dirigible Hindenburg in 1937. The Hindenburg was a dirigible filled with hydrogen, which gave it the lift it needed to keep afloat. The Hindenburg was a luxurious mode of transportation complete with a dining room and 25 private rooms. However, its voyage from Germany to the United States ended tragically on May 6, 1937, with the destruction of the airship and the loss of 36 lives because of the explosive combination of hydrogen and oxygen illustrated by the equation below. The oxidation numbers of each element are indicated below the chemical formulas:

Hydrogen underwent a loss of electrons; it was oxidized. Oxygen underwent a gain of electrons; it was reduced. In terms of half-reactions, the oxidation half reaction shows what happens to the hydrogen:

H₂ → 2H⁺ 2 electrons

while the reduction half-reaction illustrates what happens to the oxygen:

O₂ + 2 electrons – O^{2 –}

Hydrogen and oxygen combined once again to produce a fireball in the sky in 1986. This time, the space shuttle Challenger was destroyed by an explosion and all seven crew members were killed. Cold temperatures before the launch fatigued the O-rings that sealed Challenger ’s booster tanks containing 500,000 gal (1.9 million l) of liquid hydrogen and oxygen. The controlled combination of hydrogen and oxygen was intended to provide power needed to launch Challenger just as the combustion of gasoline powers a car. A spark ignited the two liquids and set off a massive, uncontrolled oxidation-reduction reaction.

Oxidation-reduction reactions are often accompanied by release of heat and sometimes flame. Combustion reactions are oxidation-reduction reactions that occur when oxygen oxidizes another material. For example, burning carbon in a lump of coal produces carbon dioxide. The reaction can be illustrated as:

In this reaction, carbon is oxidized, going from an oxidation number of 0 to+4. The oxygen is reduced from an oxidation number of 0 to—2. A similar reaction occurs when hydrocarbon fuel is burned.

Corrosion

Corrosion reactions also involve oxidation. However, these reactions are limited to the oxidation of metals, do not give off the light associated with combustion, and usually occur when moisture is present. Corrosion occurs most rapidly when metals are strained and bent; the metals rapidly oxidize in the strained regions. Corrosion can be inhibited by covering metal surfaces with paint or metals that are less easily oxidized. An example is the plating of iron with chromium on nickel. In some cases, more easily oxidized metals are used to coat or come in contact with the metal that is being protected. Then these will react more readily with the oxygen. An example is galvanizing: coating iron with zinc. Some substances such as aluminum quickly form an oxide coating in areas that are exposed, but this coating is inert to oxygen and this prevents further corrosion. That is why aluminum does not rust.

Biological processes

Photosynthesis consists of a series of oxidation-reduction reactions that begin when the carbon in carbon dioxide is reduced, and electrons are passed to molecules in the plant. When living things break down molecules of food to produce energy, carbon dioxide, and water, oxidation-reduction has taken place in the form of cellular respiration. As in photo-synthesis, a series of chemical reactions are necessary to complete cellular respiration.

Another important biological process, the nitrogen cycle, is composed of a series of oxidation and reduction reactions. Bacteria take nitrogen from the air and reduce it to ammonia and nitrates, nutrients that plants use to make proteins, nucleic acids, and other nitrogen-containing molecules needed for their metabolism. Other bacteria in soil convert nitrates back into nitrogen gas. Many of the oxidation-reduction reactions that occur in living organisms are regulated by enzymes.

Current and future uses

Dangerous as they may be, oxidation-reductions are used all the time. Burning, bleaching, metallurgy, and photography all rely on oxidation-reduction reactions. An important application of oxidation-reduction reactions is in electrochemical cells. (These types of cells should not be confused with biological cells. The word cell comes from cella, Latin for chamber or small room.) In an electrochemical cell, the oxidation reaction is physically separated from the reduction reaction, and the electrons pass between the two reactions through a conductor. Oxidation occurs at the anode and reduction occurs at the cathode. Electrochemical cells can produce electricity or consume it. Batteries and dry cells are commonly used electrochemical cells that produce electricity. A car battery is usually a 12-volt battery, made from a combination of six cells that produce two volts each.

Cells that use electricity can be used to deposit metals onto surfaces in a process known as electro-plating, which is can be used to make jewelry, mirrors, and shiny surfaces that resist abrasion, tarnish, and corrosion. Metal salts in a solution called the plating bath are reduced to metal at the cathode of the electro-chemical cell.

Oxidation-reduction reactions are widely used to produce chemicals used in manufacturing. The chemical that is produced in the most volume in the United States is sulfuric acid, which is made by oxidizing sulfur with oxygen to produce sulfur trioxide (SO₃), then dissolved in water to produce sulfuric acid, H₂SO₄.

Not all important oxidation-reduction reactions involve oxygen. A commonly produced chemical that does not contain oxygen is ammonia. To produce ammonia, NH₃, by an oxidation-reduction reaction, nitrogen and hydrogen are combined with a catalyst under pressure at 932^°F (500^°C). The nitrogen is oxidized and the hydrogen is reduced. The resulting ammonia can then be used to make fertilizers, dyes, explosives, cleaning solutions, and polymers.

Hydrogen acts as a reducing agent in many manufacturing processes. It can be used to make shortening from vegetable oils in a process known as hydrogenation. It can even reduce ions of metals such as silver and tungsten to pure metals.

Oxidation-reduction reactions are an important component of chemical analysis. Potassium permangante and cerium (IV) solutions can be used as strong oxidizing agents in the analysis of iron, tin, peroxide, vanadium, molybdenum, titanium, and uranium. Potassium dichromate is an oxidizing agent used in the analysis of organic materials in water and wastewater.

Oxidation-reduction reactions can be used to bleach materials and sanitize water. Sodium hypochlorite is used as a liquid laundry bleach and as a solid component of dishwasher powders and cleansers. Calcium hypochlorite is often used to sanitize swimming pools, killing bacteria in water by oxidizing them. Ozone is a powerful oxidizing agent that can also be used to purify water by destroying bacteria and organic pollutants. Water that has been sanitized by ozone is free of the unpleasant taste, smell, and byproducts associated with chlorinated water.

Metals are rarely found free in nature, but occur in ores, where metals are in their oxidized form. They must be reduced to the metals (oxidation number zero) in order to be used. Some metals are easily reduced. For example, mercury can be produced from a mercury sulfide ore simply by heating it in air. Iron is produced from ore by heating with coke (impure carbon) and oxygen. The coke reduces the iron in the ore. Other metals are more difficult to reduce and are obtained only after electrons are pumped into their ores using electricity. Aluminum is such a metal. As long as oxygen is around, corrosion will act to reverse the reduction of the metals achieved in metallurgy. Metals that are most resistant to corrosion are those with high standard reduction potentials such as gold and platinum.

Oxidation-reduction reactions are also responsible for food spoilage. The main source of oxidation is oxygen from the air. Preservatives that are added to foods are often reducing agents.

Oxidation reactions are important in many reactions that keep our bodies going. But oxidation has also been blamed for aging, cancer, hardening of the arteries, and rheumatoid arthritis. Research is being done to evaluate the benefits of antioxidants in foods and dietary supplements. These are natural reducing agents such as fat soluble vitamin E and vitamin C (ascorbic acid).

These substances might inhibit the damaging byproducts of oxidation reactions that can occur in the human body after exposure to some toxic chemicals. One concern, however, is that substances do not always act the same way in the human body that they do in nature. For example, vitamin C is a reducing agent. If lemon juice is brushed onto a cut apple, the vitamin C in the lemon juice will prevent the browning of the apple caused by oxidation of the apple by the air. However, vitamin C might act as an oxidizing agent in the body.

The reaction can be harnessed as a source of energy. When hydrogen and oxygen are carefully fed into a fuel cell, the oxidation-reduction reaction can be used to provide electrical power, for example, for spacecraft. The only byproduct of the reaction between hydrogen and oxygen is nonpolluting water. Another application of the hydrogen/oxygen reaction is to use hydrogen combustion to power vehicles. Currently, hydrogen is produced from water using electricity and it takes more energy to make the hydrogen than is obtained from its combustion. In the future, hydrogen might be made using solar energy and would provide a nonpolluting fuel.

The natural ability of algae and other water plants to oxidize harmful materials in sewage has been used in sewage lagoons, also known as oxidation pond systems. Small volumes of raw sewage can be treated simply by directing the sewage into shallow ponds containing algae and other water vegetation. In Belgium, nitrates are removed from wastewater by bacteria that reduce the nitrates to nitrogen which can be safely released into the atmosphere.

See also Cell, electrochemical.

Resources

BOOKS

Atkins, P.W., and J. A. Beran. General Chemistry. 2nd ed. New York: Scientific American Books, 1992, pps. 102-117.

Kostiner, Edward. Study Keys to Chemistry. Theme 17 “Oxidation-Reduction Reactions and Electrochemistry.” Barron’s Educational Series, Inc., 1992.

KEY TERMS

Combustion —An oxidation-reduction reaction that occurs so rapidly that noticeable heat and light are produced.

Corrosion —A reaction in which a metal is oxidized and oxygen is reduced, usually in the presence of moisture.

Disproportionation —An oxidation-reduction reaction in which the same chemical species is oxidized and reduced.

Electrochemical cell —A device in which an oxidation reaction is physically separated from a reduction reaction in a way that allows electrons to flow between them.

Half-reaction —The isolated oxidation or reduction reaction that is a part of a complete oxidation-reduction reaction.

Oxidant (oxidizing agent) —A chemical substance that oxidizes materials by removing electrons from them.

Oxidation —A process in which a chemical substances loses electrons and undergoes an increase in oxidation number.

Reductant (reducing agent) —A chemical substance that reduces materials by donating electrons to them.

Reduction —The process by which an atom’s oxidation state is decreased by a gain of one or more electrons.

Lide, D. R., ed. CRC Handbook of Chemistry and Physics. Boca Raton, FL: CRC Press, 2001.

Raven, Peter H. and George B. Johnson. “Oxidation-Reduction: The Flow of Energy in Living Things” and “The Nitrogen Cycle,” Biology. 3rd ed. Dubuque, IA: Wm. C. Brown Publishers, 1992, pps. 138-139, 490-491.

PERIODICALS

Halliwell, Barry. “Antioxidants: Sense or Speculation?” Nutrition Today, Vol. 29, no. 6 (November/December 1994): 15-19.

OTHER

New Mexico University. College of Arts and Sciences, Department of Chemistry and Biochemistry. “Redox Reactions” <http://www.chemistry.nmsu.edu/studntres/chem112.spring.1997/Redox.html> (accessed October 8, 2006).

Catherine Hinga Haustein

OXIDATION-REDUCTION REACTIONS

CONCEPT

Most people have heard the term "oxidation" at some point or another, and, from the sound of the word, may have developed the impression that it has something to do with oxygen. Indeed it does, because oxygen has a tendency to draw electrons to itself. This tendency, rather than the presence of oxygen itself, is actually what identifies oxidation, defined as a process in which a substance loses electrons. The oxidation of one substance is always accompanied by reduction, or the gaining of electrons, on the part of another substance—hence the term "oxidation-reduction reaction," sometimes called a redox reaction. The world is full of examples of this highly significant form of chemical reaction. One such example is combustion, or an even more rapid form of combustion, explosion. Likewise the metabolism of food, as well as other biological processes, involves oxidation and reduction reactions. So, too, do a number of processes that take place on the surfaces of metals: when iron rusts; when copper turns green; or when aluminum forms a coating of aluminum oxide that prevents it from rusting. Oxidation-reduction reactions also play a major role in electrochemistry, which has a highly useful application to daily life in the form of batteries.

HOW IT WORKS

Chemical Reactions

A chemical reaction is a process whereby the chemical properties of a substance are changed by a rearrangement its atoms. The change produced by a chemical reaction is quite different from a purely physical change, which does not affect the fundamental properties of the substance itself. A piece of copper can be heated, melted, beaten into different shapes, and so forth, yet throughout all those changes, it remains pure copper, an element of the transition metals family.

But suppose a copper roof is exposed to the elements for many years. Copper is famous for its highly noncorrosive quality, and this, combined with its beauty, has made it a favored material for use in the roofs of imposing buildings. (Because it is relatively expensive, few middle-class people today can afford a roof entirely made of copper, but sometimes it is used as a decorative touch—for instance, over the entryway of a house.) Eventually, however, copper does begin to corrode when exposed to air for long periods of time.

Over the years, exposed copper develops a thin layer of black copper oxide, and as time passes, traces of carbon dioxide in the air contribute to the formation of greenish copper carbonate. This explains why the Statue of Liberty, covered in sheets of copper, is green, rather than having the reddish-golden hue of new, uncorroded copper.

EXTERNAL VS. INTERNAL CHANGE.

The preceding paragraphs describe two very different phenomena. The first was a physical change in which the chemical properties of a substance—copper—remained unaltered. The second, on the other hand, involved a chemical change on the surface of the copper, as copper atoms bonded with carbon and oxygen atoms in the air to form something different from copper. The difference between these two types of changes can be likened to varieties of changes in a person's life—an external change on the one hand, and a deeply rooted change on the other.

A person may move to another house, job, school, or town, yet the person remains the same. Many sayings in the English language express this fact: for instance, "Wherever you go, there you are," or "You can take the boy out of the country, but you can't take the country out of the boy." Moving is simply a physical change. On the other hand, if a person changes belief systems, overcomes old feelings (or succumbs to new ones), changes lifestyles in a profound manner, or in any other way changes his or her mind about something important—this is analogous to a chemical change. In these instances, the person, like the surface of the copper described above, has changed not merely in external properties, but in inner composition.

"LEOthe Lion Says 'GER'"

Chemical reactions are addressed in depth within the essay devoted to that subject, which discusses—among other subjects—many ways of classifying chemical reactions. These varieties of chemical reaction are not all mutually exclusive, as they relate to different aspects of the reaction. As noted in the review of various reaction types, one of the most significant is an oxidation-reduction reaction (sometimes called a redox reaction) involving the transfer of electrons.

As its name implies, an oxidation-reduction reaction is really two processes: oxidation, in which electrons are lost, and reduction, in which electrons are gained. Though these are defined separately here, they do not occur independently; hence the larger reaction of which each is a part is called an oxidation-reduction reaction. In order to keep the two straight, chemistry teachers long ago developed a useful, if nonsensical, mnemonic device: "LEO the lion says 'GER'." LEO stands for "Loss of Electrons, Oxidation," and "GER" means "Gain of Electrons, Reduction."

Many, though not all, oxidation-reduction reactions involve oxygen. Oxygen combines readily with other elements, and in so doing, it tends to grab electrons from those other elements' atoms. As a result, the oxygen atom becomes an ion (an atom with an electric charge)—specifically, an anion, or negatively charged ion.

In interacting with another element, oxygen becomes reduced, while the other element is oxidized to become a cation, or a positively charged ion. This, too, is easy to remember: oxygen itself, obviously, cannot be oxidized, so it must be the one being reduced. But since not all oxidation-reduction reactions involve oxygen, perhaps the following is a better way to remember it. Electrons are negatively charged, and the element that takes them on in an oxidation-reduction reaction is reduced—just as a person who thinks negative thoughts are "reduced" if those negative thoughts overcome positive ones.

Oxidation Numbers

An oxidation number (sometimes called an oxidation state) is a whole-number integer assigned to each atom in an oxidation-reduction reaction. This makes it easier to keep track of the electrons involved, and to observe the ways in which they change positions. Here are some rules for determining oxidation number.

• 1. The oxidation number for an atom of an element not combined with other elements in a compound is always zero.
• 2. For an ion of any element, the oxidation number is the same as its charge. Thus a sodium ion, which has a charge of +1 and is designated symbolically as Na⁺, has an oxidation number of +1.
• 3. Certain elements or families form ions in predictable ways:
• a. Alkali metals, such as sodium, always form a +1 ion; oxidation number = +1.
• b. Alkaline earth metals, such as magnesium, always form a +2 ion; oxidation number = +2.
• c. Halogens, such as fluorine, form −1 ions; oxidation number = −1.
• d. Other elements have predictable ways to form ions; but some, such as nitrogen, can have numerous oxidation numbers.
• 4. The oxidation number for oxygen is −2 for most compounds involving covalent bonds.
• 5. When hydrogen is involved in covalent bonds with nonmetals, its oxidation number is +1.
• 6. In binary compounds (compounds with two elements), the element having greater electronegativity is assigned a negative oxidation number that is the same as its chargewhen it appears as an anion in ionic compounds.
• 7. When a compound is electrically neutral, the sum of its elements' oxidation states is zero.
• 8. In an ionic chemical species, the sum of the oxidation states for its constituent elements must equal the overall charge.

These rules will not be discussed here; rather, they are presented to show some of the complexities involved in analyzing an oxidation-reduction reaction from a structural stand-pont—that is, in terms of the atomic or molecular reactions. For the most part, we will be observing oxidation-reductions phenomenologically, or in terms of their outward effects. A good chemistry textbook should provide a more detailed review of these rules, along with a table showing oxidation numbers of elements and binary compounds.

Oxidation-reduction reactions are easier to understand if they are studied as though they were two half-reactions. Half the reaction involves what happens to the substances and electrons in the oxidizing portion, while the other half-reaction indicates the activities of substances and electrons in the reduction portion.

REAL-LIFE APPLICATIONS

Combustion and Explosions

As with any type of chemical reaction, combustion takes place when chemical bonds are broken and new bonds are formed. It so happens that combustion is a particularly dramatic type of oxidation-reduction reaction: whereas we cannot watch iron rust, combustion is a noticeable event. Even more dramatic is combustion that takes place at a rate so rapid that it results in an explosion.

Coal is almost pure carbon, and its combustion in air is a textbook example of oxidation-reduction. Although there is far more nitrogen than oxygen in air (which is a mixture rather than a compound), nitrogen is very unreactive at low temperatures. For this reason, it can be used to clean empty fuel tanks, a situation in which the presence of pure oxygen is extremely dangerous. In any case, when a substance burns, it is reacting with the oxygen in air.

As one might expect from what has already been said about oxidation-reduction, the oxygen is reduced while the carbon is oxidized. In terms of oxidation numbers, the oxidation number of carbon jumps from 0 to 4, while that of oxygen is reduced to −2. As they burn, these two form carbon dioxide or CO₂, in which the two −2 charges of the oxygen atoms cancel out the +4 charge of the carbon atom to yield a compound that is electrically neutral.

COMBUSTION IN HUMAN EXPERIENCE.

Combustion has been a significant part of human life ever since our prehistoric ancestors learned how to harness the power of fire to cook food and light their caves. We tend to think of premodern times—to use the memorable title of a book by American historian William Manchester, about the Middle Ages—as A World Lit Only By Fire. In fact, our modern age is even more combustion-driven than that of our forebears.

For centuries, burning animal fat—in torches, lamps, and eventually in candles—provided light for humans. Wood fires supplied warmth, as well as a means to cook meals. These were the main uses of combustion, aside from the occasional use of fire in warfare or for other purposes (including that ghastly medieval form of execution, burning at the stake). One notable military application, incidentally, was "Greek fire," created by the Byzantines in the seventh century a.d. A mixture of petroleum, potassium nitrate, and possibly quicklime, Greek fire could burn on water, and was used in naval battles to destroy enemy ships.

For the most part, however, the range of activities to which combustion could be applied was fairly narrow until the development of the steam engine in the period from the late seventeenth century to the early nineteenth century. The steam engine applied the combustion of coal to the production of heat for boiling water, which in turn provided the power to run machinery. By the beginning of the twentieth century, combustion had found a new application in the internal combustion engine, used to power automobiles.

EXPLOSIONS AND EXPLOSIVES.

An internal combustion engine does not simply burn fuel; rather, by the combined action of the fuel injectors (in a modern vehicle), in concert with the pistons, cylinders, and spark plugs, it actually produces small explosions in the molecules of gasoline. These produce the output of power necessary to turn the crankshaft, and ultimately the wheels.

An explosion, in simple terms, is a sped-up form of combustion. The first explosives were invented by the Chinese during the Middle Ages, and these included not only fireworks and explosive rockets, but gunpowder. Ironically, however, China rejected the use of gunpowder in warfare for many centuries, while Europeans took to it with enthusiasm. Needless to say, Europeans' possession of firearms aided their conquest of the Americas, as well as much of Africa, Asia, and the Pacific, during the period from about 1500 to 1900.

The late nineteenth and early twentieth centuries saw the development of new explosives, such as TNT or trinitrotoluene, a hydrocarbon. Then in the mid-twentieth century came the most fearsome explosive of all: the nuclear bomb. A nuclear explosion is not itself the result of an oxidation-reduction reaction, but of something much more complex—either the splitting of atoms (fission) or the forcing together of atomic nuclei (fusion).

Nuclear bombs release far more energy than any ordinary explosive, but the resulting blast also causes plenty of ordinary combustion. When the United States dropped atomic bombs on the Japanese cities of Hiroshima and Nagasaki in August 1945, those cities suffered not only the effects of the immediate blast, but also massive fires resulting from the explosion itself.

FUELING THE SPACE SHUTTLE.

Oxidation-reduction reactions also fuel the most advanced form of transportationknown today, the space shuttle. The actual orbiter vehicle is relatively small compared to its external power apparatus, which consists of two solid rocket boosters on either side, along with an external fuel tank.

Inside the solid rocket boosters are ammonium perchlorate (NH₄ClO₄) and powdered aluminum, which undergo an oxidation-reduction reaction that gives the shuttle enormous amounts of extra thrust. As for the larger single external fuel tank, this contains the gases that power the rocket: hydrogen and oxygen.

Because these two are extremely explosive, they must be kept in separate compartments. When they react, they form water, of course, but in doing so, they also release vast quantities of energy. The chemical equation for this is: 2H₂ + O₂ →2H₂O + energy.

On January 28, 1986, something went terribly wrong with this arrangement on the space shuttle Challenger. Cold weather had fatigued the O-rings that sealed the hydrogen and oxygen compartments, and the gases fed straight into the flames behind the shuttle itself. This produced a powerful and uncontrolled oxidation-reduction reaction, an explosion that took the lives of all seven astronauts aboard the shuttle.

The Environment and Human Health

Combustion, though it can do much good, can also do much harm. This goes beyond the obvious: by burning fossil fuels or hydrocarbons, excess carbon (in the form of carbon dioxide and carbon monoxide) is released to the atmosphere, with a damaging effect on the environment.

In fact, oxidation-reduction reactions are intimately connected with the functioning of the natural environment. For example, photosynthesis, the conversion of light to chemical energy by plants, is a form of oxidation-reduction reaction that produces two essentials of human life: oxygen and carbohydrates. Likewise cellular respiration, which along with photosynthesis is discussed in the Carbon essay, is an oxidation-reduction reaction in which living things break down molecules of food to produce energy, carbon dioxide, and water.

Enzymes in the human body regulate oxidation-reduction reactions. These complex proteins, of which several hundred are known, act as catalysts, speeding up chemical processes in the body. Oxidation-reduction reactions also take place in the metabolism of food for energy, with substances in the food broken down into components the body can use.

OXIDATION: SPOILING AND AGING.

At the same time, oxidation-reduction reactions are responsible for the spoiling of food, the culprit here being the oxidation portion of the reaction. To prevent spoilage, manufacturers of food items often add preservatives, which act as reducing agents.

Oxidation may also be linked with the effects of aging in humans, as well as with other conditions such as cancer, hardening of the arteries, and rheumatoid arthritis. It appears that oxygen molecules and other oxidizing agents, always hungry for electrons, extract these from the membranes in human cells. Over time, this can cause a gradual breakdown in the body's immune system.

To forestall the effects of oxidation, some doctors and scientists recommend antioxidants—natural reducing agents such as vitamin C and vitamin E. The vitamin C in lemon juice can be used to prevent oxidizing on the cut surface of an apple, to keep it from turning brown. Perhaps, some experts maintain, natural reducing agents can also slow the pace of oxidation in the human body.

Forming a New Surface on Metal

Clearly, oxidization can have a corrosive effect, and nowhere is this more obvious than in the corrosion of metals by exposure to oxidizing agents—primarily oxygen itself. Most metals react with O₂, and might corrode so quickly that they become useless, were it not for the formation of a protective coating—an oxide.

Iron forms an oxide, commonly known as rust, but this in fact does little to protect it from corrosion, because the oxide tends to flake off, exposing fresh surfaces to further oxidation. Every year, businesses and governments devote millions of dollars to protecting iron and steel from oxidation by means of painting and other measures, such as galvanizing with zinc. In fact, oxidation-reduction reactions virtually define the world of iron. Found naturally only in ores, the element is purified by heating the ore with coke (impure carbon) in the presence of oxygen, such that the coke reduces the iron.

COINAGE METALS.

Copper, as we have seen, responds to oxidation by corroding in a different way: not by rusting, but by changing color. A similar effect occurs in silver, which tarnishes, forming a surface of silver sulfide, or Ag₂S. Copper and silver are two of the "coinage metals," so named because they have often been used to mint coins. They have been used for this purpose not only because of their beauty, but also due to their relative resistance to corrosion. This resistance has, in fact, earned them the nickname "noble metals."

The third member of this mini-family is gold, which is virtually noncorrosive. Wonderful as gold is in this respect, however, no one is likely to use it as a roofing material, or for any such large-scale application involving its resistance to oxidation. Aside from the obvious expense, gold is soft, and not very good for structural uses, even if it were much cheaper. Yet there is such a "wonder metal": one that experiences virtually no corrosion, is cheap, and strong enough in alloys to be used for structural purposes. Its name is aluminum.

ALUMINUM.

There was a time, in fact, when aluminum was even more expensive than gold. When the French emperor Napoleon III wanted to impress a dinner guest, he arranged for the person to be served with aluminum utensils, while less distinguished personages had to settle for "ordinary" gold and silver.

In 1855, aluminum sold for $100,000 a pound, whereas in 1990, the going rate was about $0.74. Demand did not go down—in fact, it increased exponentially—but rather, supply increased, thanks to the development of an inexpensive aluminum-reduction process. Two men, one American and one French, discovered this process at the same time: interestingly, their years of birth and death were the same.

Aluminum was once a precious metal because it proved extremely difficult to separate from oxygen. The Hall-Heroult process overcame the problem by applying electrolysis—the use of an electric current to produce a chemical change—as a way of reducing Al³⁺ ions (which have a high affinity for oxygen) to neutral aluminum atoms. In the United States today, 4.5% of the total electricity output is used for the production of aluminum through electrolysis.

The foregoing statistic is staggering, considering just how much electricity Americans use, and it indicates the importance of this once-precious metal. Actually, aluminum oxidizes just like any other metal—and does so quite quickly, as a matter of fact, by forming a coating of aluminum oxide (Al₂O₃). But unlike rust, the aluminum oxide is invisible, and acts as a protective coating. Chromium, nickel, and tin react to oxygen in a similar way, but these are not as inexpensive as aluminum.

Electrochemistry and Batteries

Electrochemistry is the study of the relationship between chemical and electrical energy. Among its applications is the creation of batteries, which use oxidation-reduction reactions to produce an electric current.

A basic battery can be pictured schematically as two beakers of solution connected by a wire. In one solution is the oxidizing agent; in the other, a reducing agent. The wire allows electrons to pass back and forth between the two solutions, but to ensure that the flow goes both ways, the two solutions are also connected by a "salt bridge." The salt bridge contains a gel or solution that permits ions to pass back and forth, but a porous membrane prevents the solutions from actually mixing.

In the lead storage battery of an automobile, lead itself is the reducing agent, while lead (IV) oxide (PbO₂) acts as the oxidizing agent. A highly efficient type of battery, able to with stand wide extremes in temperature, the lead storage battery has been in use since 1915. Along the way, features have been altered, but the basic principles have remained—a testament to the soundness of its original design.

The batteries people use for powering all kinds of portable appliances, from flashlights to boom boxes, are called dry cell batteries. In contrast to the model described above, using solutions, a dry cell (as its name implies) involves no liquid components. Instead, it utilizes various elements in a range of combinations, including zinc, magnesium, mercury, silver, nickel, and cadmium. The last two are applied in the nickel-cadmium battery, which is particularly useful because it can be recharged over and over again by an external current. The current turns the products of the chemical reactions in the battery back into reactants.

WHERE TO LEARN MORE

"Batteries." Oregon State University Department of Chemistry (Web site). <http://www.chem.orst.edu/ch411/scbatt.htm> (June 4, 2001).

"Batteries and Fuel Cells" (Web site). <http://vectorsite.com/ttfuelc.html> (June 4, 2001).

Borton, Paula and Vicky Cave. The Usborne Book of Batteries and Magnets. Tulsa, OK: EDC Publishing, 1995.

Craats, Rennay. The Science of Fire. Milwaukee, WI: Gareth Stevens Publishing, 2000.

Knapp, Brian J. Oxidation and Reduction. Danbury, CT: Grolier Educational, 1998.

"Oxidation Reduction Links" (Web site). <http://users.erols.com/merosen/redox.htm> (June 4, 2001).

"Oxidation-Reduction Reactions." General Chemistry (Web site). <http://ull.chemistry.uakron.edu/genchem/11/> (June 4, 2001).

"Oxidation-Reduction Reactions: Redox." UNC-Chapel Hill Chemistry Fundamentals (Web site). <http://www.shodor.org/unchem/advanced/redox/> (June 4, 2001).

Yount, Lisa. Antoine Lavoisier: Founder of Modern Chemistry. Springfield, NJ: Enslow Publishers, 1997.

Zumdahl, Steven S. Introductory Chemistry: A Foundation, fourth edition. Boston: Houghton Mifflin, 2000.

KEY TERMS

ANION:

An ion with a negative charge; pronounced "AN-ie-un."

BINARY COMPOUND:

A compound involving two elements.

CHEMICAL REACTION:

A process whereby the chemical properties of a substance are changed by a rearrangement of the atoms in the substance.

CHEMICAL SPECIES:

A generic term used for any substance studied in chemistry—whether it be an element, compound, mixture, atom, molecule, ion, and so forth.

COVALENT BONDING:

A type of chemical bonding in which two atoms share valence electrons.

ELECTROCHEMISTRY:

The study of the relationship between chemical and electrical energy.

ELECTROLYSIS:

The use of an electric current to produce a chemical change.

ELECTRONEGATIVITY:

The relative ability of an atom to attract valence electrons.

CATION:

An ion with a positive charge; pronounced "KAT-ie-un."

ION:

An atom or group of atoms that has acquired an electric charge due to the loss or gain of electrons.

OXIDATION:

A chemical reaction in which a substance loses electrons. It is always accompanied by reduction; hence the term oxidation-reduction reaction.

OXIDATION NUMBER:

A number as signed to each atom in an oxidation-reduction reaction as a means of keeping track of the electrons involved. Another term for oxidation number is "oxidation state."

OXIDATION-REDUCTION REACTION:

A chemical reaction involving the transfer of electrons.

REDOX REACTION:

Another name for an oxidation-reduction reaction.

REDUCTION:

A chemical reaction in which a substance gains electrons. It is always accompanied by oxidation; hence the term oxidation-reduction reaction.

VALENCE ELECTRONS:

Electrons that occupy the highest energy levels in anatom. These are the only electrons involved in chemical bonding.

Oxidation-reduction reaction

Oxidation-reduction reactions, also known as redox reactions, are chemical processes in which electrons are transferred from one atom, ion, or molecule to another. Explosions, fires, batteries, and even our own bodies are powered by oxidation-reduction reactions. When iron rusts or colored paper bleaches in the sun , oxidation-reduction has taken place.

Oxidation-reduction reactions can be thought of as a combination of two processes: oxidation, in which electrons are lost, and reduction, in which electrons are gained. The two processes cannot occur independently of each other. A mnemonic device used by chemists to help keep things straight is "LEO says 'GER,'" which stands for Loss of Electrons, Oxidation. Gain of Electrons, Reduction.

The driving force of oxidation-reduction reactions is the transfer of electrons. Although it is sometimes difficult to remember what happens to electrons during oxidation and what happens during reduction, a look at familiar processes can help keep this straight. Some of the first oxidation-reduction reactions understood by chemists were those involving oxygen . Oxygen, the most plentiful element on Earth , combines readily with numerous other elements. When combined with other elements in a compound or molecule, oxygen frequently is an electron "hog." It takes electrons away from many other elements and this oxidizes them. The oxygen takes the negatively charged electrons and becomes a negatively charged ion. The oxygen has been reduced, somewhat like taking in negative thoughts will reduce a person's positive attitude. An example of this is the reaction between oxygen in the air and iron. The iron metal becomes positively charged and the oxygen becomes negatively charged. The two charged ions now attract each other and hang around together in the form of iron oxide, or rust.

History

Probably the earliest human use of oxidation-reduction reactions occurred 7,500-4,500 years ago in the Copper/Bronze Age. Copper ores were heated in the presence of carbon to produce copper metal. In this process, the copper in the ore was reduced to copper metal and the carbon was oxidized to carbon dioxide . This same process was applied to iron ores during the Iron Age, which occurred 4,500-3,500 years ago.

The use of oxidation-reduction reactions has long been a part of pottery making. Differences in color in the clay or glaze can be produced when firing pottery under oxidizing conditions when lots of oxygen is present or under low oxygen reducing conditions, such as with a partially closed kiln or a fire with green leaves on it. Clay containing iron will be orange-red if fired under oxidizing conditions due to the presence of red iron oxide and black in reducing conditions when black iron oxide-in which the iron has a lower oxidation number-forms. Among the people who have historically used oxidizing and reducing fire conditions are the Native Americans in the southwestern United States and the Greeks in the Early Bronze Age.

Another historic use of oxidation-reduction reactions is in explosives , substances that burn-are oxidized-so rapidly that they cause huge amounts of pressure . Gunpowder, thought to be the first explosive used, originated in China as early as a.d. 960. It appeared in Europe around the thirteenth century. Eighteen forty-six was a banner year for explosives-nitrocellulose and nitroglycerin were both developed that year. TNT (trinitrotolmene) saw widespread use in World War I. Since 1955, a commonly used cheap and powerful explosive has been a mixture of ammonium nitrate and fuel oil. This was used to bomb the Federal Building in Oklahoma City in 1995, and its use may be curtailed in the future.

Fireworks, a colorful and bright form of oxidationreduction reactions, are believed to have first been used in China in the sixth century.

An important step in the understanding of oxidationreduction reactions was the discovery of oxygen. Joseph Priestley (1733-1804) was the first scientist on record to prepare oxygen in the laboratory. This historic reaction was also an oxidation-reduction reaction. Priestley heated mercury oxide and formed elemental mercury and oxygen. In this reaction, mercury was reduced and the oxide ion was oxidized. Antoine Lavoisier (1743-1794) recognized that when substances are burned, they combine with oxygen. He even figured out that our bodies burn food and give off carbon dioxide as we produce energy . Tragically, the life of this great chemist was ended prematurely when he was beheaded during the French Revolution.

Oxidation numbers

Oxidation numbers, sometimes called oxidation states, help chemists keep track of the numbers of electrons that surround each atom in a chemical reaction, and how they change in oxidation-reduction reactions. When an atom gains an electron (is reduced), its oxidation number is increased by one. There are some simple rules for assigning oxidation numbers to elements in chemical compounds. These rules are:

1. The oxidation number of an element, having neither gained nor lost any of its electrons, is zero . For example, the oxidation number of pure copper, Cu, is zero, as is the oxidation number of each oxygen atom in a molecule of oxygen, O2.
2. The oxidation number of an elemental ion is the same as its charge. An ion of copper with a plus two charge, Cu2+, has an oxidation number of plus two. A fluoride ion, F–, has an oxidation number of minus one.
3. Some elements almost always form compounds in which they have a particular oxidation number. Aluminum always forms a plus three ion and therefore exists in the plus three oxidation state in compounds. Sodium and other alkali metals almost always form a plus one ion; its oxidation state is plus one. Hydrogen can form compounds in which the hydrogen atom has an oxidation number of either plus one or minus one. When the hydrogen has an oxidation number of plus one, it is written on the left hand side of the chemical formula. If its oxidation number is minus one, it is written on the right hand side. Oxygen usually has a minus two oxidation number. Chlorine and other halogens usually take on a minus one charge. Other elements are not so predictable. Nitrogen can have oxidation numbers of +5,+4,+3, +2,+1, and -3.
4. The sum of the oxidation numbers in a neutral molecule or compound is zero. Table salt , with the chemical formula of sodium chloride , is made up of two ions, a positively charged sodium ion and a negatively charged chloride ion. A water molecule consists of two hydrogen atoms , each having an oxidation number of plus one, and an oxygen atom with an oxidation number of minus two.

It is often easier to follow oxidation-reduction reactions if they are split into two half reactions. One-half reaction indicates what is happening to the chemical substances and electrons in the oxidation portion of the reaction. The other half-reaction does the same for the reduction portion. The complete reaction is the sum of the two half reactions.

A useful tool for chemists is a table of standard reduction potentials. This table lists common half reactions, and assigns each a numerical value that indicates how easily the reduction reaction proceeds-that is, how eagerly electrons are accepted. A high standard reduction potential value indicates that the substance is easily reduced. A low standard reduction potential indicates that the substance is easily oxidized-it prefers to lose electrons. In general, a substance will oxidize something that has a lower reduction potential than it has. The halogens, chemical elements found in group 17 of the periodic table , are strong oxidizing agents because their atoms readily accept negative ions. The alkali metals such as sodium, found on the left side of the periodic table in group 1, are strong reducing agents because their atoms readily give up an electron, becoming positive ions. The arbitrary zero point for standard reduction potentials has been designated as this reaction:

This reaction has been assigned a potential of 0.000 volts under standard conditions. The standard reduction potential for fluorine gas is 2.890 volts while that for sodium metal is -2.714 volts.

Examples of oxidation-reduction reactions

Combustion

Let us look at an oxidation-reduction more chemically as we examine what happened to the Hindenburg in 1937. The Hindenburg was a dirigible filled with hydrogen, which gave it the lift it needed to keep afloat. The Hindenburg was a luxurious mode of transportation complete with a dining room and 25 private rooms. However, its first voyage, from Germany to the United States, ended tragically with the destruction of the airship and the loss of 36 lives because of the explosive combination of hydrogen and oxygen illustrated by this equation. The oxidation numbers of each element are indicated below the chemical formulas:

Hydrogen underwent a loss of electrons; it was oxidized. Oxygen underwent a gain of electrons; it was reduced. In terms of half reactions, the oxidation half reaction shows what happens to the hydrogen:

while the reduction half reaction illustrates what happens to the oxygen:

Hydrogen and oxygen combined once again to produce a fireball in the sky in 1986. This time, the space shuttle Challenger was destroyed by explosion and all seven crew members aboard were killed. Cold temperatures before the launch fatigued the O-rings that sealed the Challenger's booster tanks containing 500,000 gal (1.9 million l) of liquide hydrogen and oxygen. The controlled combination of hydrogen and oxygen was intended to provide power needed to launch the Challenger just as the combustion of gasoline provides power to a car. A spark ignited the two liquids and set off a massive uncontrolled oxidation-reduction reaction

Oxidation-reduction reactions are ofter accompanied by release of heat and sometimes, flame. Combustion reactions are oxidation-reduction reactions that occur when oxygen oxidizes another material. For example, burning carbon in a lump of coal produces carbon dioxide. The reaction can be illustrated as:

In this reaction, carbon is oxidized, going from an oxidation number of 0 to +4. The oxygen is reduced from an oxidation number of 0 to -2. A similar reaction occurs when hydrocarbon fuel is burned.

Corrosion

Corrosion reactions also involve oxidation of substances by oxygen. However, these reactions are limited to the oxidation of metals, do not give off the light associated with combustion, and usually occur when moisture is present. Corrosion occurs most rapidly when metals are strained and bent; the metals rapidly oxidize in the strained regions. Corrosion can be inhibited by covering metal surfaces with paint or metals which are less easily oxidized. An example is the plating of iron with chromium on nickel. In some cases, more easily oxidized metals are used to coat or come in contact with the metal that is being protected. Then these will react more readily with the oxygen. An example is galvanizing: coating iron with zinc. Some substances such as aluminum quickly form an oxide coating in areas that are exposed, but this coating is inert to oxygen and this prevents further corrosion. That is why aluminum does not "rust."

Biological processes

Photosynthesis consists of a series of oxidation-reduction reactions which begin when carbon in carbon dioxide is reduced and electrons are passed to molecules in the plant . When living things break down molecules of food to produce energy, carbon dioxide, and water, oxidation-reduction has taken place in the form of cellular respiration. As in photosynthesis, a series of chemical reactions are necessary to complete cellular respiration. Another important biological process, the nitrogen cycle , is composed of a series of oxidation and reduction reactions. Bacteria take nitrogen from the air and reduce it to ammonia and nitrates, nutrients that plants use to make proteins , nucleic acids, and other nitrogen-containing molecules needed for their metabolism . Other bacteria in soil convert nitrates back into nitrogen gas. Many of the oxidation-reduction reactions that occur in living organisms are regulated by enzymes.

Current and future uses

Dangerous as they may be, oxidation-reductions are used all the time. Burning, bleaching, batteries, metallurgy , and photography all rely on oxidation-reduction reactions. An important application of oxidation-reduction reactions is in electrochemical cells. (These types of cells should not be confused with biological cells. The word cell comes from cella, Latin for chamber or small room.) In an electrochemical cell, the oxidation reaction is physically separated from the reduction reaction, and the electrons pass between the two reactions through a conductor. Oxidation occurs at the anode and reduction occurs at the cathode . Electrochemical cells can produce electricity or consume it. Batteries and dry cells are commonly used electrochemical cells that produce electricity. The battery for your car is probably a 12 volt battery, made from a combination of six cells producing two volts each.

Cells that use electricity can be used to deposit metals onto surfaces in a process known as electroplating. Electroplating can be used to make jewelry, mirrors , and shiny surfaces resistant to abrasion, tarnishing and corrosion. Metal salts in a solution called the plating bath are reduced to metal at the cathode of the electro-chemical cell.

Oxidation-reduction reactions are widely used to produce chemicals that are used in manufacturing. The chemical that is produced in the most volume in the United States is sulfuric acid . It is made by oxidizing sulfur with oxygen to produce sulfur trioxide (SO3). This is dissolved in water to give sulfuric acid, H2SO4.

Not all important oxidation-reduction reactions involve oxygen. A commonly produced chemical that does not contain oxygen is ammonia. To produce ammonia, NH3, by an oxidation-reduction reaction, nitrogen and hydrogen are combined together under pressure at 932°F (500°C) with a catalyst. The nitrogen is oxidized and the hydrogen is reduced. The resulting ammonia can then be used to make fertilizers , dyes, explosives, cleaning solutions, and polymers.

Hydrogen acts as a reducing agent in many manufacturing processes. It can be used to make shortening from vegetable oils in a process known as hydrogenation . It can even reduce ions of metals such as silver and tungsten to pure metals.

Oxidation-reduction reactions are an important component of chemical analysis. Potassium permangante and cerium (IV) solutions can be used as strong oxidizing agents in the analysis of iron, tin, peroxide, vanadium, molybdenum, titanium , and uranium . Potassium dichromate is an oxidizing agent used in the analysis of organic materials in water and wastewater.

Oxidation-reduction reactions can be used for bleaching materials and sanitizing water. Sodium hypochlorite is used in solution as a liquid laundry bleach and as a solid component of dishwasher powders and cleansers. Calcium hypochlorite is often used for swimming pool sanitation. The hypochlorites kill bacteria in water by oxidizing them. Ozone is a powerful oxidizing agent that can also be used to purify water. The ozone destroys bacteria and organic pollutants. Water that has been sanitized by ozone is free of the unpleasant taste, smell, and byproducts associated with chlorinated water.

Metals are rarely found free in nature, but occur in ores. The metals are in their oxidized form in the ores and must be reduced to the metals (oxidation number zero) in order to be used. Some metals are easily reduced. For example, mercury can be produced from a mercury sulfide ore simply by heating it in air. Iron is produced from ore by heating with coke (impure carbon) and oxygen. The coke reduces the iron in the ore. Other metals are more difficult to reduce and are only obtained after electrons are pumped into their ores using electricity. Aluminum is such a metal. As long as oxygen is around, corrosion will act to reverse the reduction of the metals achieved in metallurgy. Metals that are most resistant to corrosion are those with high standard reduction potentials such as gold and platinum.

Oxidation-reduction reactions are responsible for food spoilage. The main source of oxidation is oxygen from the air. Preservatives that are added to foods are often reducing agents.

Oxidation reactions are important in many reactions that keep our bodies going. But oxidation has also been blamed for aging, cancer , hardening of the arteries , and rheumatoid arthritis . Research is being done to evaluate the benefits of antioxidants in foods and dietary supplements. Antioxidants are natural reducing agents such asfat soluble vitamin E and vitamin C (ascorbic acid). These substances might inhibit damaging byproducts of oxidation reactions that can occur in the human body after exposure to some toxic chemicals. One concern that scientists studying antioxidants have is that substances do not always act the same way in the human body that they do outside of it. For example, vitamin C is a reducing agent. If lemon juice is squirted on a cut apple, the vitamin C in the juice will prevent the browning of the apple that is caused by oxidation of the apple by the air. However, vitamin C might act as an oxidizing agent in the body.

The reaction can be harnessed as a source of energy. When hydrogen and oxygen are carefully fed into an electrochemical cell called a fuel cell, the oxidation-reduction reaction can be used to provide electrical power, for example, for space craft. The only byproduct of the reaction between hydrogen and oxygen is non-polluting water. Another application of the hydrogen/oxygen reaction is to use hydrogen combustion to power vehicles. Currently, hydrogen is produced from water using electricity and it takes more energy to make the hydrogen than is obtained from its combustion. In the future, hydrogen might be made using solar energy and would provide a non-polluting fuel.

The natural ability of algae and other water plants to oxidize harmful materials in sewage has been used in sewage lagoons, also known as oxidation pond systems. Small volumes of raw sewage can be treated simply by directing the sewage into shallow ponds containing algae and other water vegetation. In Belgium, nitrates are removed from wastewater by bacteria that reduce the nitrates to nitrogen which can be safely released into the atmosphere.

See also Cell, electrochemical.

Resources

books

Atkins, P. W., and J. A. Beran. General Chemistry. 2nd ed. New York: Scientific American Books, 1992.

Kostiner, Edward. "Oxidation-Reduction Reactions and Electrochemistry." Study Keys to Chemistry. Barron's Educational Series, Inc, 1992.

Lide, D. R., ed. CRC Handbook of Chemistry and Physics. Boca Raton: CRC Press, 2001.

Raven, Peter H., and George B. Johnson. "Oxidation-Reduction: The Flow of Energy in Living Things" and "The Nitrogen Cycle." Biology. 3rd ed. Dubuque: Wm. C. Brown Publishers, 1992.

periodicals

Halliwell, Barry. "Antioxidants: Sense or Speculation?" Nutrition Today, 29, no. 6, (November/December 1994): 15-19.

other

"Explosion on the Lady Delta." Video. Films for the Humanities and Sciences. P.O. Box 2053, Princeton, NJ 08453-2053.

Catherine Hinga Haustein

KEY TERMS

Combustion

—An oxidation-reduction reaction that occurs so rapidly that noticeable heat and light are produced.

Corrosion

—A reaction in which a metal is oxidized and oxygen is reduced, usually in the presence of moisture.

Disproportionation

—An oxidation-reduction reaction in which the same chemical species is oxidized and reduced.

Electrochemical cell

—A device in which an oxidation reaction is physically separated from a reduction reaction in a way that allows electrons to flow between them.

Half reaction

—The isolated oxidation or reduction reaction that is a part of a complete oxidation-reduction reaction.

Oxidant (oxidizing agent)

—A chemical substance which oxidizes materials by removing electrons from them.

Oxidation

—A process in which a chemical substances loses electrons and undergoes an increase in oxidation number.

Reductant (reducing agent)

—A chemical substance which reduces materials by donating electrons to them.

Reduction

—The process by which an atom's oxidation state is decreased, by its gaining one or more electrons.

Oxidation-reduction reaction

The term oxidation-reduction reaction actually refers to two chemical reactions that always occur at the same time: oxidation and reduction. Oxidation-reduction reactions are also referred to more simply as redox reactions. Oxidation, reduction, and redox reactions can all be defined in two ways.

The simpler definitions refer to reactions involving some form of oxygen. As an example, pure iron can be produced from iron oxide in a blast furnace by the following reaction:

3 C + 2 Fe₂O₃ → 4 Fe + 3 CO₂

In this reaction, iron oxide (Fe₂O₃) gives away its oxygen to carbon (C). In chemical terms, the carbon is said to be oxidized because it has gained oxygen. At the same time, the iron oxide is said to be reduced because it has lost oxygen.

Because of its ability to give away oxygen, iron oxide is called an oxidizing agent. Similarly, because of its ability to take on oxygen, carbon is said to be a reducing agent. Oxidation and reduction always occur together. If one substance gives away oxygen (oxidation), a second substance must be present to take on that oxygen (reduction).

By looking at the above example, you can see that the following statements must always be true:

An oxidizing agent (in this case, iron oxide) is always reduced.

Words to Know

Combustion: An oxidation-reduction reaction that occurs so rapidly that noticeable heat and light are produced.

Corrosion: An oxidation-reduction reaction in which a metal is oxidized and oxygen is reduced, usually in the presence of moisture.

Oxidation: A process in which a chemical substance takes on oxygen or loses electrons.

Oxidizing agent: A chemical substance that gives up oxygen or takes on electrons from another substance.

Reducing agent: A chemical substance that takes on oxygen or gives up electrons to another substance.

Reduction: A process in which a chemical substance gives off oxygen or takes on electrons.

A reducing agent (in this case, carbon) is always oxidized.

Redox and electron exchanges

For many years, chemists thought of oxidation and reduction as involving the element oxygen in some way or another. That's where the name oxidation came from. But they eventually learned that other elements behave chemically in much the same way as oxygen. They decided to revise their definition of oxidation and reduction to make it more general—to apply to elements other than oxygen.

The second definition for oxidation and reduction is not as easy to see. It is based on the fact that when two elements react with each other, they do so by exchanging electrons. In an oxidation-reduction reaction like the one above, the element that is oxidized always loses electrons. The element that is reduced always gains electrons. The more general definition of redox reactions, then, involves the gain and loss of electrons rather than the gain and loss of oxygen.

In the reaction below, for example, sodium metal (Na) reacts with chlorine gas (Cl₂) in such a way that sodium atoms lose one electron each to chlorine atoms:

2 Na + Cl₂ → 2 NaCl

Because sodium loses electrons in this reaction, it is said to be oxidized. Because chlorine gains electrons in the reaction, it is said to be reduced.

Types of redox reactions. Redox reactions are among the most common and most important chemical reactions in everyday life. The great majority of those reactions can be classified on the basis of how rapidly they occur. Combustion is an example of a redox reaction that occurs so rapidly that noticeable heat and light are produced. Corrosion, decay, and various biological processes are examples of oxidation that occurs so slowly that noticeable heat and light are not produced.

Combustion. Combustion means burning. Any time a material burns, an oxidation-reduction reaction occurs. The two equations below show what happens when coal (which is nearly pure carbon) and gasoline (C₈H₁₈) burn. You can see that the fuel is oxidized in each case:

C + O₂ → CO₂

2 C₈H₁₈ + 25 O₂ → 16 CO₂ + 18 H₂O

In reactions such as these, oxidation occurs very rapidly and energy is released. That energy is put to use to heat homes and buildings; to drive automobiles, trucks, ships, airplanes, and trains; to operate industrial processes; and for numerous other purposes.

Rust. Most metals react with oxygen to form compounds known as oxides. Rust is the name given to the oxide of iron and, sometimes, the oxides of other metals. The process by which rusting occurs is also known as corrosion. Corrosion is very much like combustion, except that it occurs much more slowly. The equation below shows perhaps the most common form of corrosion, the rusting of iron.

4 Fe + 3 O₂ → 2 Fe₂O₃

Decay. The compounds that make up living organisms, such as plants and animals, are very complex. They consist primarily of carbon, oxygen, and hydrogen. A simple way to represent such compounds is to use the letters x, y, and z to show that many atoms of carbon, hydrogen, and oxygen are present in the compounds.

When a plant or animal dies, the organic compounds of which it is composed begin to react with oxygen. The reaction is similar to the combustion of gasoline shown above, but it occurs much more slowly. The process is known as decay, and it is another example of a common oxidation-reduction reaction. The equation below represents the decay (oxidation) of a compound that might be found in a dead plant:

C_xH_yO_z + O₂ → CO₂ + H₂O

Biological processes. Many of the changes that take place within living organisms are also redox reactions. For example, the digestion of food is an oxidation process. Food molecules react with oxygen in the body to form carbon dioxide and water. Energy is also released in the process. The carbon dioxide and water are eliminated from the body as waste products, but the energy is used to make possible all the chemical reactions that keep an organism alive and help it to grow.

Oxidation-reduction reaction

Oxidation-reduction reactions are significant to many geochemical reactions (e.g., the production of natural gas ). In addition, oxidation-reduction reactions are critical in many carbon-based biological processes.

The term oxidation was originally used to describe reactions in which an element combines with oxygen . In contrast, reduction meant the removal of oxygen. By the turn of this century, it became apparent that oxidation always seemed to involve the loss of electrons and did not always involve oxygen. In general, oxidation-reduction reactions involve the exchange of electrons between two species.

An oxidation reaction is defined as the loss of electrons, while a reduction reaction is defined as the gain of electrons. The two reactions always occur together and in chemically equivalent quantities. Thus, the number of electrons lost by one chemical species (a variation of an element or chemical compound) is always equal to the number of electrons gain by another chemical species. The combination of the two reactions is known as a redox reaction. Chemical species that participate in redox reactions are described as either reducing or oxidizing agents. An oxidizing agent is a chemical species that causes the oxidation of another chemical species. The oxidizing agent accomplishes this by accepting electrons in a reaction. A reducing agent causes the reduction of another chemical species by donating electrons to the reaction.

In general, a strong oxidizing agent is a species that has an attraction for electrons and can oxidize another chemical species. The standard voltage reduction of an oxidizing agent is a measure of the strength of the oxidizing agent. The more positive the chemical species' standard reduction potential, the stronger the chemical species is as an oxidizing agent.

In reactions where the reactants and products are not ionic, there is still a transfer of electrons between chemical species. Chemists have devised a way to keep track of electrons during chemical reactions where the charge on the atoms is not readily apparent. Charges on atoms within compounds are assigned oxidation states (or oxidation numbers). An oxidation number is defined by a set of rules that describes how to divide up electrons shared within compounds. Oxidation is defined as an increase in oxidation state, while reduction is defined as a decrease in oxidation state. Because an oxidizing agent accepts electrons from another chemical species, a component atom of the oxidizing agent will decrease in oxidation number during the redox reaction.

There are many examples of oxidation-reduction reactions in the world. Important processes that involve oxidationreduction reactions include combustion reactions that convert energy stored in fuels into thermal energy, the corrosion of metals , and metabolic reactions.

Oxidation-reduction reactions occur in both physical and biological settings (where carbon-containing compounds such as carbohydrates are oxidized). The burning of natural gas is an oxidation-reduction reaction that releases energy [CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) + energy]. In many organisms, including humans, redox reactions burn carbohydrates that provide energy [C₆H₁₂O₆(aq) + 6O₂(g) → 6CO₂(g) + 6H₂O(l)]. In both examples, the carbon-containing compound is oxidized, and the oxygen is reduced.

See also Chemical bonds and physical properties; Chemical elements; Chemistry