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 as fat 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.
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.
Halliwell, Barry. "Antioxidants: Sense or Speculation?" Nutrition Today, 29, no. 6, (November/December 1994): 15-19.
"Explosion on the Lady Delta." Video. Films for the Humanities and Sciences. P.O. Box 2053, Princeton, NJ 08453-2053.
Catherine Hinga Haustein
Science EncyclopediaScience & Philosophy: Overdamped to PeatOxidation-Reduction Reaction - History, Oxidation Numbers, Corrosion, Biological Processes, Current And Future Uses - Examples of oxidation-reduction reactions