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oxidation–reduction reactions
Sci-Tech Dictionary: oxidation-reduction reaction
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Home > Library > Science > Sci-Tech Dictionary
(′äk·sə′dā·shən ri′dək·shən rē′ak·shən)
(chemistry) An oxidizing chemical change, where an element's positive valence is increased
(electron loss), accompanied by a simultaneous reduction of an associated element (electron
gain).
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Science of Everyday Things: Oxidation-Reduction Reactions
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Home > Library > Science > Science of Everyday Things
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 "oxidationreduction 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. Oxidationreduction 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.
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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 standpont—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 oxidationreduction. 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 largescale 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.