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.
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, 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.
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:
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
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:
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. The space shuttle Challenger was destroyed by an explosion that killed all seven crew members. Cold temperatures before the launch fatigued 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 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.
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 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, 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.
Cells that use electricity can be used to deposit metals onto surfaces in a process known as electroplating, which 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 electrochemical 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 by simply 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.