Hydrogen peroxide (HY-druh-jin per-OK-side) is a clear,
colorless, somewhat unstable liquid with a bitter taste.
When absolutely pure, the compound is quite stable. Even
small amounts of impurities (such as iron or copper),
however, act as catalysts that increase its tendency to
decompose, sometimes violently, into water and nascent
oxygen (O). To prevent decomposition, small amounts of
inhibitors, such as acetanilide or sodium stannate are
added to pure hydrogen peroxide and hydrogen peroxide
solutions.
Hydrogen peroxide was discovered in 1818 by French
chemist Louis Jacques The´nard (1777–1857). It was first used
commercially in the 1800s, primarily to bleach hats. Today,
industrial processes make about 500 million kilograms
(1 billion pounds) of hydrogen peroxide annually for use in a
wide variety of applications ranging from whitening of teeth
to propelling rockets.
Hydrogen peroxide occurs in very small amounts in nature.
It is formed when atmospheric oxygen reacts with water
to form H2O2. Hydrogen peroxide is also present in plant and
animal cells as the byproduct of metabolic reactions that
occur in those cells.
The large amounts of hydrogen peroxide used in industry
are prepared in a complex series of reactions that begins
with any one of a family of compounds known as the alkyl
anthrahydroquinones, such as ethyl anthrahydroquinone.
The anthrahydroquinones are three-ring compounds that
can be converted back and forth between two or more similar
structures. During the conversion from one structure to
another, hydrogen peroxide is produced as a byproduct. The
anthraquinone is continuously regenerated during the production
of hydrogen peroxide, making the process very efficient.
Other methods for the preparation of hydrogen peroxide
are also available. For example, the electrolysis of sulfuric
acid results in the formation of a related compound, peroxysulfuric
acid (H2SO5), which then reacts with water to form
hydrogen peroxide. A third method of preparation involves
the heating of isopropyl alcohol [2-propanol; (CH3)2CHOH] at
high temperature and pressure, resulting in the formation of
hydrogen peroxide as one product of the reaction.
Most of hydrogen peroxide’s applications depend on the
fact that it tends to break down, releasing a single atom of
nascent oxygen (O):
H2O2 ! H2O + (O)
The term nascent oxygen refers to a single atom of
oxygen, a structure that is chemically very active. Nascent
oxygen tends to be a very strong oxidizing agent. For example,
the use of hydrogen peroxide with which most people are
probably familiar is as an antiseptic, a substance used to kill
germs. Hydrogen peroxide achieves this result because the
nascent oxygen it releases destroys bacteria, fungi, and other
microorganisms that cause disease.
The most important industrial application of hydrogen
peroxide—its use in the pulp and paper industry—also
depends on its oxidizing properties. In this case, it is used
to bleach the materials of which paper is made, converting
colored compounds to colorless compounds. About 55 percent
of all hydrogen peroxide made in the United States is used
for this purpose. Another nine percent is used in the bleaching
of other materials, such as textiles, furs, feathers, and
hair. Another important application of hydrogen peroxide is
in water and sewage treatment plants, where its antibacterial
action destroys disease-causing organisms in the water. Some
additional uses of hydrogen peroxide include:
• In bakeries to condition dough and make it easier to
work with;
• For cleaning metals;
• As a rocket propellant;
• In the preparation of other organic and inorganic compounds;
• As a neutralizing agent in the production of wines; and
• As a disinfectant in the treatment of seeds for agricultural
purposes.
The hydrogen peroxide solutions with which people come
into contact at home pose little or no health hazard because
the concentration of the compound is very low, usually about
3 percent. Prolonged use of hydrogen peroxide may cause
burns on the skin, however, and the more concentrated solutions
used in industry present more serious hazards. They can
be toxic if ingested and are explosive if not stored properly.
