Redox reactions everyday examples

1. REDOX REACTION- EVERYDAY EXAMPLES
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 (C 8 H 18 ) burn. You can see that the fuel is
oxidized in each case:
C + O 2 → CO 2
2 C 8 H 18 + 25 O 2 → 16 CO 2 + 18 H 2 O
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 O 2 → 2 Fe 2 O 3
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 x H y O z + O 2 → CO 2 + H 2 O
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.
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2. REDOX REACTIONS – EXAMPLES AND APPLICATIONS
C OMBUSTION AND E XPLOSIONS
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.
As one might expect from what has already been said about oxidation-reduction, the
oxygen is reduced while the carbon is oxidized.
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.
F ORMING A NEW SURFACE ON M ETAL - CORROSION
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 O 2 , and might corrode so quickly that they become useless,
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.
COINAGE METALS. –COPPER,SILVER, and GOLD
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 Ag 2 S. 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.
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

3. 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
Al
3+
ions (which have a high affinity for oxygen) to neutral aluminum atoms.
Aluminum oxidizes just like any other metal—and does so quite quickly, as a matter of
fact, by forming a coating of aluminum oxide (Al 2 O 3 ). 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.
E LECTROCHEMISTRY AND B ATTERIES
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.
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