Oxidation-Reduction Reaction

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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:

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.
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.
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.
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.

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

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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.