Ethylene Glycol

   Ethylene glycol (ETH-uh-leen GLYE-kol) is clear, colorless,
syrupy liquid with a sweet taste. One should not attempt
to confirm the compound’s taste, however, as it is toxic. In
recent years, more than 4 billion kilograms (9 billion pounds)
of ethylene glycol has been produced in the United States
annually. The compound is used primarily as an antifreeze
and in the manufacture of a number of important chemical
compounds, including polyester fibers, films, bottles, resins,
and other materials.
   Ethylene glycol was first prepared in 1859 by the
French chemist Charles Adolphe Wurtz (1817–1884).
Wurtz’s discovery did not find an application, however,
until the early twentieth century, when the compound was
manufactured for use in World War I (1914–1918) in the
manufacture of explosives and as a coolant. By the 1930s,
a number of uses for the compound had been found, and the
chemical industry began producing ethylene glycol in large
quantities.
   The primary method of producing ethylene glycol
involves the hydration of ethylene oxide, a ring compound
consisting of two methylene (-CH2) groups and one oxygen
atom. Hydration is the process by which water is added to a
compound. The hydration of ethylene oxide is conducted at a
temperature of about 383F (195C) without a catalyst, or at
about 50C to 70C (122F to 158F) with a catalyst, usually a
strong acid, either process resulting in a yield of at least 90
percent of ethylene glycol.
   Other methods of preparation are also available. For
example, the compound can be produced directly from synthesis
gas, a mixture of carbon monoxide and hydrogen; or by
treating ethylene (CH2=CH2) with oxygen in an acetic acid
solution using a catalyst of tellurium oxide or bromide ion.
One of the first major uses of ethylene glycol was as a
radiator coolant in airplanes. The compound actually made
possible a change in the design of airplanes. At one time,
plain water was used as the coolant in airplane radiators. The
faster the airplane flew, the greater the risk that its radiator
would boil over. Adding ethylene glycol to the water raised
the boiling point of the coolant and allowed airplanes to fly
faster with smaller radiators. This change was especially
useful in the construction of military airplanes used in
combat.

   Ethylene glycol is still used extensively as a coolant and
antifreeze in cooling systems. It is also used as a deicing
fluid for airport runways, cars, and boats. Brake fluids and
shock-absorber fluids often contain ethylene glycol as protection
against freezing. About 26 percent of all the ethylene
glycol made in the United States is used for some kind of
cooling or antifreeze application.

   The largest single use of ethylene glycol today is in the
manufacture of a plastic called polyethylene terephthalate
(PET). PET’s primary application is in the manufacture of
plastic bottles, an application that accounts for about a third
of all the ethylene glycol made in the United States. Large
amounts of PET are also used in the manufacture of polyester
fibers and films. Some additional uses of the compound
include:

• As a humectant (a substance that attracts moisture) in
keeping some food, tobacco, and industrial products
dry;
• As a solvent in some paints and plastics;
• In the dyeing of leathers and textiles;
• In the manufacture of printing inks, wood stains, ink
for ball-point pens, and adhesives;
• In the production of artificial smoke and fog for theatrical
productions;
• As a stabilizer in the soybean-based foam sometimes
used to extinguish industrial fires; and
• In the manufacture of specialized types of explosives.
   Ethylene glycol poses a number of potential health and
safety hazards. It is very flammable and highly toxic. Ingestion
of the compound may cause nausea, vomiting, abdominal
pain, weakness, convulsions, and cardiac problems.
Higher doses can result in severe kidney damage that leads
to death.

Amino Acid

amino acid An organic molecule possessing both
acidic carboxylic acid (–COOH) and basic amino
(–NH2) groups attached to the same tetrahedral carbon
atom.
Amino acids are the principal building blocks of
proteins and enzymes. They are incorporated into
proteins by transfer RNA according to the genetic
code while messenger RNA is being decoded by ribo-
somes. The amino acid content dictates the spatial
and biochemical properties of the protein or enzyme
during and after the final assembly of a protein.
Amino acids have an average molecular weight of
about 135 daltons. While more than 50 have been dis-
covered, 20 are essential for making proteins, long
chains of bonded amino acids.
Some naturally occurring amino acids are alanine,
arginine, asparagine, aspartic acid, cysteine, glutamine,
glutamic acid, glycine, histidine, isoleucine, leucine,
lysine, methionine, phenylalanine, proline, serine, thre-
onine, tryptophan, tyrosine, and valine.
The two classes of amino acids that exist are
based on whether the R-group is hydrophobic or
hydrophilic. Hydrophobic or nonpolar amino acids
tend to repel the aqueous environment and are located
mostly in the interior of proteins. They do not ionize
or participate in the formation of hydrogen bonds. On
the other hand, the hydrophilic or polar amino acids
tend to interact with the aqueous environment, are
usually involved in the formation of hydrogen bonds,
and are usually found on the exterior surfaces of pro-
teins or in their reactive centers. It is for this reason
that certain amino acid R-groups allow enzyme reac-
tions to occur.
The hydrophilic amino acids can be further subdi-
vided into polar with no charge, polar with negatively
charged side chains (acidic), and polar with positively
charged side chains (basic).

Chlorophyll : The Green

  Chlorophyll (KLOR-uh-fill) is the pigment that gives
plants, algae, and cyanobacteria their green color. The name
comes from a combination of two Greek words, chloros,
meaning ‘‘green’’ and phyllon, meaning ‘‘leaf.’’ Chlorophyll is
the substance that enables plants to create their own food
through photosynthesis.
At least five forms of chlorophyll exist. They are:
• chlorophyll a (also known as a-chlorophyll), with a formula
of C55H72O5N4Mg
• chlorophyll b (also known as b-chlorophyll), with a formula
of C55H70O6N4Mg
• Chlorophyll c1, with a formula of C35H30O5N4Mg
• Chlorophyll c2, with a formula of C35H28O5N4Mg
• Chlorophyll d, with a formula of C54H70O6N4Mg
Chlorophyll a occurs in all types of plants and in algae.
Chlorophyll b is found primarily in land plants. Chlorophyll
c1 and chlorophyll c2 are present in various types of algae.
Chlorophyll d is found in red algae.

   All forms of chlorophyll have a similar chemical structure.
They have a complex system of rings made of carbon
and nitrogen known as a chlorin ring. The five forms of
chlorophyll differ in the chemical groups attached to the
chlorin ring. These differences result in slightly different
colors of the five chlorophylls.
   French chemists Pierre-Joseph Pelletier (1788–1842) and
Joseph-Bienaime´ Caventou (1795–1877) first isolated chlorophyll
in 1817. In 1865, German botanist Julius von Sachs
(1832–1897) demonstrated that chlorophyll is responsible
for photosynthetic reactions that take place within the cells
of leaves. In the early 1900s, Russian chemist Mikhail Tsvett
(1872–1920) developed a technique known as chromatography
to separate different forms of chlorophyll from each
other. In 1929, the German chemist Hans Fischer (1881–
1945) determined the complete molecular structure, making
possible the first synthesis of the molecule in 1960 by the
American chemist Robert Burns Woodward (1917–1979).

   Plants make chlorophyll in their leaves using materials
they have absorbed through their roots and leaves. The
synthesis of chlorophyll requires several steps involving
complex organic compounds. First, the plant converts a common
amino acid, glutamic acid (COOH(CH2)2CH(NH2)COOH)
into an alternative form known as 5-aminolevulinic acid
(ALA). Two molecules of ALA are then joined to form a ring
compound called porphobilinogen. Next, four molecules of
porphobilinogen are joined to form an even larger ring structure
with side chains. Oxidation of the larger ring structure
introduces double bonds in the molecule, giving it the ability
to absorb line energy. Finally, a magnesium atom is introduced
into the center of the ring and side chains are added to
the ring to give it its final chlorophyll configuration.
   Plants store chlorophyll in their chloroplasts, organelles
(small structures) that carry out the steps involved in photosynthesis.
Each chloroplast contains many clusters of several
hundred chlorophyll molecules called photosynthetic units.
When a photosynthetic unit absorbs light energy, chlorophyll
molecules move to a higher energy state, initiating
the process of photosynthesis. The overall equation for the
process of photosynthesis is 6CO2 + 6H2O ! C6H12O6 + 6O2.
That simple equation does not begin to suggest the complex
nature of what happens during photosynthesis. Botanists
divide that process into two major series of reactions: the light
reactions and the dark reactions. In the light reactions, plants
use the energy obtained from sunlight to make two compounds,
adenosine triphosphate (ATP) and nicotinamide adenine
dinucleotide phosphate (NADPH). ATP and NADPH are
not themselves components of carbohydrates, the final products
of photosynthesis. Instead, they store energy that is
used to make possible a series of thirteen different chemical
reactions that occur during the dark stage of photosynthesis
that result in the conversion of carbon dioxide and water
to the simple carbohydrate glucose (C6H12O6).

Riboflavin: The Vitamin B2

 

 

Riboflavin (REY-bo-FLAY-vin), commonly known as vitamin
B2, is an orange-yellow crystalline solvent with a bitter
taste. It is relatively stable when exposed to heat, but tends
to decompose in the presence of light for extended periods of
time. Riboflavin is used in the body for a variety of functions,
including the metabolism of carbohydrates for the
production of energy and the production of red blood cells.
Riboflavin was found in 1879 by Alexander Wynter Blyth (1844-1921) who noticed a compound in cow’s milk that glowed with a yellow fluorescence
when exposed to light. Blyth called the compound
lachtochrome (lachto- = ‘‘milk’’ and -chrome = color), but was
unable to determine its chemical composition or its chemical
properties. In fact, it was not until the 1930s that the
chemical nature of the compound was determined. The
Swiss chemist Paul Karrer (1889–1971) and the Austrian-
German chemist Richard Kuhn (1900–1967) independently
determined the chemical structure of riboflavin and first
synthesized the compound. The name riboflavin is derived
from the fact that the vitamin was first found in association
with the sugar ribose.

   Naturally, plants and microorganisms can synthesize riboflavin.
Some foods rich in riboflavin are brewer’s yeast, dark
green vegetables, mushrooms, legumes, nuts, milk and other
dairy products, sweet potatoes, and pumpkins. Bacteria that
live in the human digestive tract are also able to synthesize
some riboflavin, but not enough to meet the body’s requirement
for the vitamin.
   Riboflavin is produced synthetically using either the
genetically-modified bacterium Bacillus subtilis or a fungus
called Ashbya gossifyii. The bacteria or fungus are cultured
in a large vat that has been seeded with small amounts of
riboflavin. Over time, the organisms generate large quantities
of riboflavin until some desired amount of the compound
has been produced. The vat is then heated to a
temperature sufficient to kill the bacteria or fungi, leaving
crystalline riboflavin behind. The riboflavin is then separated
and purified.

   The human body needs riboflavin to use oxygen efficiently
in the metabolism of amino acids, fatty acids, and
carbohydrates. The vitamin is involved in the synthesis of
niacin (another B vitamin), it activates vitamin B6, and it
helps the adrenal gland to produce hormones. It helps the
body make antibodies to fight disease and infection, regulates
the thyroid gland, and is important in maintaining healthy
hair, nails, and skin. Riboflavin is especially important during
periods of rapid growth because it is involved in the formation
and growth of cells, especially red blood cells.
The human body needs riboflavin to use oxygen efficiently
in the metabolism of amino acids, fatty acids, and
carbohydrates. The vitamin is involved in the synthesis of
niacin (another B vitamin), it activates vitamin B6, and it
helps the adrenal gland to produce hormones. It helps the
body make antibodies to fight disease and infection, regulates
the thyroid gland, and is important in maintaining healthy
hair, nails, and skin. Riboflavin is especially important during
periods of rapid growth because it is involved in the formation
and growth of cells, especially red blood cells.
most likely to suffer from riboflavin deficiency problems are
those with anorexia (a condition in which people refuse to
eat adequate amounts of food), older people with poor diets,
alcoholics (because alcohol impairs a person’s ability to
absorb and use the vitamin), and newborn babies being treated
for jaundice by exposure to ultraviolet light (because
light destroys riboflavin).

Dimethyl Ketone

Dimethyl ketone (DYE-meth-el KEY-tone) is a clear, colorless,
highly volatile and highly flammable liquid with a
characteristic sweet odor and taste. The compound is almost
universally known in chemistry laboratories and industrial
applications by its common name of acetone.
  Acetone was apparently first prepared in 1610 by the
French alchemist Jean Be´guin (c. 1550–c. 1650). Be´guin
obtained acetone by heating lead acetate (also known as
Saturn’s salt) to a high temperature. He obtained a sweetsmelling,
very flammable liquid that he named ‘‘burning
spirit of Saturn.’’ One of the first uses to which the substance
was put was as a solvent in the extraction of the active
constituents of opium. In 1833, the French chemist Antoine
Bussy (1794–1882) gave the compound its modern name of
acetone. The correct chemical formula for acetone was determined
independently in 1832 by the French chemist Jean
Baptiste Andre´ Dumas (1800–1884) and the German chemist
Justus von Liebig (1803–1873).
 
Most of the acetone produced today is made by one of
four methods:
• In the Hock process, cumene [C6H5CH(CH3)2] is first
oxidized to produce cumene hydroperoxide
[C6H5C(CH3)2COOH], which is then reduced to produce
acetone and phenol (C6H5OH); or
• Isopropyl alcohol (2-propanol; CH3CHOHCH3) is oxidized
over a catalyst to obtained acetone; or
• Butane (C4H10) is oxidized to obtain acetone; or
• Acetone is obtained as a by-product of the manufacture
of glycerol [C3H5 (OH)3].

  Acetone’s primary applications are based on its ability to
dissolve such a wide array of organic substances. It is used as
a solvent for paints, varnishes, lacquers, inks, glues, rubber
cements, fats, oils, waxes, and various types of rubber and
plastics. It is perhaps best known to the average person as
the primary ingredient in nail polish remover. The largest
single use of the compound is as a raw material in the
manufacture of other organic chemicals, such as chloroform,
acetic acid, iodoform, bromoform, isoprene, rayon, and photographic
film. It also finds application in storing acetylene
gas (because it absorbs up to 24 times its own weight of the
gas), to clean and dry chemical equipment and electronic
parts, and for the extraction of components of plant and
animal tissues.
  The primary safety concern about acetone is its extreme
flammability. Workers who handle the compound must use
great care to prevent its coming into contact or even being in
the vicinity of open flames. Under the proper conditions,
acetone is also explosive. Exposure of the skin, eyes, and
respiratory system to acetone may produce mild symptoms,
such as dizziness, headaches, and disorientation and irritation
of the eyes and skin. Such conditions are rare, however,
and no long-term health effects of the compound have as yet
been discovered.

Retinol: The Vitamin A

  Retinol (RET-uh-nol) is the scientific name for vitamin A,
a vitamin found only in animals. It occurs as a yellowish to
orange powder with a slight brownish cast and is a relatively
stable compound. Retinol is converted in the body from an
alcohol to the corresponding aldehyde, retinal (C20H28O), one
of the primary chemical compounds involved in the process
by which light is converted to nerve impulses in the retina of
the eye. Vitamin A is also required for a number of other
biochemical reactions in the body, including growth and
development of tissue and maintenance of the immune system

  Vitamin A is synthesized in animal bodies through a
variety of pathways. One important source of vitamin A is a
group of related compounds called the carotenes, substances
responsible for the yellowish or orangish appearance of
fruits and vegetables such as carrots, sweet potatoes, squash,
cantaloupe, apricots, pumpkin, and mangos. Some leafy
green vegetables, such as collard greens, spinach, and kale,
are also good sources of the carotenes. The most important of
the carotenes is b-carotene (beta-carotene), C40H56. The oxidation
of carotenes in animal bodies converts them to retinol.

  The chemical structure of retinol was determined in 1931
by Swiss chemist Paul Karrer (1889–1971), and the compound
was first prepared synthetically shortly thereafter by Austrian-
German chemist Richard Kuhn (1900–1967). The first
successful process for producing retinol commercially was
developed in the mid-1940s by German chemist Otto Isler
(1920–1992), then employed at the pharmaceutical company
Roche, located in Sissein, Germany. Isler’s process involved a
complex series of reactions that begins with the combination
of a fourteen carbon hydrocarbon and a six carbon hydrocarbon
to create the fundamental backbone from which the
retinol molecule is constructed. Regular production of vitamin
A began in 1948 with a projected output of 10 kilograms
per month, which before long was raised to 50 kilograms per
month. The Roche plant at Sissein continues to produce
retinol today.
  Vitamin A is probably best known for its role in maintaining
normal vision. Deficiencies of the compound are
likely to manifest themselves earliest in a variety of eye
problems, most commonly night blindness. Night blindness
is a condition in which one loses the ability to distinguish
objects in reduced light. If left untreated, vitamin A deficiencies
may lead to decreased ability to see in normal light and,
eventually, to complete blindness.
  But vitamin A has been shown to have a number of other
functions in the body. It is essential for the maintenance of
growth, bone formation, reproduction, proper immune system
function, and healing of wounds. A number of additional
claims have been made for the compound, although evidence
is not as strong as it is for the above functions. For example,
it may be effective in preventing or treating a variety of
conditions such as measles, intestinal parasites, osteoporosis,
inflammatory bowel disease, bone marrow disorders, certain
types of cancer, tuberculosis, peritonitis, osteoarthritis, food
poisoning, Alzheimer’s disease, miscarriage, and HIV/AIDS.
In each of these cases, evidence is not yet strong enough to
show a clear-cut connection between retinol and disease, but
research is being conducted to determine how strong the
association may be.
  Retinol is available commercially in a variety of formulations,
including tablets, capsules, and creams. Such products
usually contain a modified form of retinol that is more easily
absorbed by the body. For example, a product known as
tretinoin is a synthetic form of retinol known as all-trans
retinoic acid. The term all trans means that all of the double
bonds in retinoic acid are located on the same side of the
molecule. Products containing tretinoin are used to treat
acne, pimples, wrinkles, blackheads, freckles, sun-spots, and
even pre-cancerous lesions. They work by increasing the rate
with which the skin sheds old cells and replaces them with
new cells.
  Vitamin A supplements in pill or capsule form are available
in two formulations, those that contain retinol and
those that contain beta carotene. It is not possible to take
too much of the latter type of vitamin A. The body will not
convert excess amounts of carotene into retinol but will,
instead, excrete the excess in the urine or stool. An excess
of retinol-based vitamin A, by contrast, may result in certain
medical problems. Since the vitamin is fat soluble, in
may be stored in body fat and reach relatively high concentrations
if too much is ingested. An excess of retinol in the
body may be associated with liver damage, osteoporosis,
rash, fatigue, bone and joint pain, nausea, insomnia, and
personality changes.

Hydrogen Chloride

  Hydrogen chloride (HY-druh-jin KLOR-ide) is a colorless
gas with a strong, suffocating odor. The gas is not flammable,
but is corrosive, that is, capable of attacking and reacting
with a large variety of other compounds and elements.
Hydrogen chloride is most commonly available as an aqueous
solution known as hydrochloric acid. It is one of the most
important industrial chemicals in the world. In 2004, just
over 5 million metric tons (5.5 million short tons) of hydrogen
chloride were produced in the United States, making it
the eighteenth most important chemical in the nation for
that year.
  Hydrogen chloride has probably been known as far back
as the eighth century, when the Arabian chemist Jabir ibn
Hayyan (c. 721–c. 815; also known by his Latinized name of
Geber) described the production of a gas from common table
salt (sodium chloride; NaCl) and sulfuric acid (H2SO4). The
compound was mentioned in the writings of a number of
alchemists during the Middle Ages and was probably first

produced in a reasonably pure form by the German chemist
Johann Rudolf Glauber (1604–1670) in about 1625. The first
modern chemist to prepare hydrogen chloride and describe
its properties was the English chemist Joseph Priestley
(1733–1804) in 1772. Forty years later, in 1818, the English
chemistry and physicist Humphry Davy (1778–1829) showed
that the compound consisted of hydrogen and chlorine, giving
it the correct formula of HCl.
  Commercial production of hydrogen chloride had its
beginning in Great Britain in 1823. The method of production
most popular there and, later, throughout Europe was
one originally developed by the French chemist Nicholas
Leblanc (1742–1806) in 1783. Leblanc had invented the process
as a method for producing sodium hydroxide and sodium
carbonate, two very important industrial chemicals. Hydrogen
chloride was produced as a byproduct of the Leblanc
process, a byproduct for which there was at first no use.
The gas was simply allowed to escape into the air. The suffocating
and hazardous release of hydrogen chloride prompted
governments to pass legislation requiring some other means
of disposal for the gas. In England, that law was called the
Alkali Act and was adopted by the parliament in 1863. Unable
to release hydrogen chloride into the air, manufacturers
began dissolving it in water and producing hydrochloric acid.
Before long, a number of important commercial and industrial
uses for the acid itself were discovered. The ‘‘useless’’
byproduct of the Leblanc process soon became as important
as the primary products of the process, sodium hydroxide
and sodium carbonate.

  Hydrogen chloride is still sometimes made today by the
traditional process of reacting sodium chloride (NaCl) with a
sulfate, such as sulfuric acid or iron(II) sulfate (FeSO4). However,
more than 90 percent of the hydrogen chloride produced
throughout the world today comes as the byproduct of
the chlorination of organic compounds. Chlorination is the
process by which chlorine gas reacts with an organic compound,
usually replacing some of the hydrogen present in the
compound. Since a large number of important chlorinated
organic compounds are produced each year, large amounts of
hydrogen chloride gas are produced as a byproduct. That gas
is simply removed from the reaction and stored in cylinders
for future use. Other methods of producing hydrogen chloride
include the direct synthesis of hydrogen gas and chlorine
gas (producing a very pure product) and the reaction of
sodium chloride, sulfur dioxide, oxygen, and water with each
other at high temperatures (the Hargreaves process).

  Hydrogen chloride and hydrochloric acid have some uses
in common, and some that are different from each other. In
both dry and liquid form, the largest single use of hydrogen
chloride is in the synthesis of organic and inorganic chlorides.
A large number of compounds important in commerce
and industry contain chlorine, including most pesticides,
many pharmaceuticals, and a number of polymeric products.

  Hydrochloric acid is also used widely in the processing of
metallic ores and the pickling of metals. Pickling is the
process by which a metal is cleaned, usually with an acid,
to remove rust and other impurities that have collected on
the metal. Some additional uses of hydrogen chloride and
hydrochloric acid include the following:
• In the brining of foods and other materials. Brining is
the process by which a material is soaked in a salt
solution, usually in order to preserve the material;
• In the treatment of swimming pool water;
• As a catalyst in industrial chemical reactions;
• In the manufacture of semiconductors and other electronic
components;
• To maintain the proper acidity in oil wells (to keep oil
flowing smoothly);
• For the etching of concrete surfaces;
• In the production of aluminum, titanium, and a number
of other important metals.
  Both hydrogen chloride and hydrochloric acid pose serious
health risks to humans and other animals. The gas is an
irritant to the eyes and respiratory system, causing coughing,
choking, and tearing, as well as more serious damage to tissues.
Hydrochloric acid can burn the skin and mucous membranes.
Exposure of only five parts per million of the gas can
produce noticeable symptoms of distress, and exposure of
more than 2,000 parts per million can be fatal. If hydrochloric

acid gets into the eyes, blindness may result. Since hydrochloric
acid is present in many household products, users should
exercise great care when working with such materials.

Nitrous Oxide

 

  Nitrous oxide (NYE-truss OX-side) is also known as
dinitrogen oxide, dinitrogen monoxide, nitrogen monoxide,
and laughing gas. It is a colorless, nonflammable gas with a
sweet odor. Its common name of laughing gas is derived from
the fact that it produces a sense of light-headedness when
inhaled. The gas is widely used as an anesthetic, a substance
that reduces sensitivity to pain and discomfort.
  Nitrous oxide was probably first produced by the English
chemist and physicist Robert Boyle (1627–1691), although he
did not recognize the new compound he had found. Credit for
the discovery of nitrous oxide is, therefore, usually given to
the English chemist Joseph Priestley (1733–1804), who produced
the gas in 1772 and named it ‘‘nitrous air.’’ Other early
names used for the gas include ‘‘gaseous of azote’’ (nitrogen)
and ‘‘oxide of speton.’’ The most complete experiments on the
gas were conducted by the English chemist and physicist Sir
Humphry Davy (1778–1829), who tested nitrous oxide on
himself and his friends. He found that the gas could lessen

pain and discomfort and provided a sense of relaxation and
well-being. Before long, doctors were making use of Davy’s
discovery by using nitrous oxide as an anesthetic.
  The public found other uses for the gas as well. During
the Victorian period in England, members of the upper class
often held laughing gas parties at which people gathered to
inhale nitrous oxide as a recreational drug, rather than for
any therapeutic purpose. In the United States, the showman
P. T. Barnum (1810–1891) created a sideshow exhibit in which
people were invited to test the effects of inhaling nitrous
oxide. After seeing a demonstration of this kind, the American
dentist Horace Wells (1815–1848) first used nitrous oxide
as an anesthetic on his patients.
  In 1868, the American surgeon Edmund Andrews (1824–
1904) extended the use of nitrous oxide as an anesthetic for
his surgical patients. He mixed the gas with oxygen to
ensure that patients received enough oxygen while receiving
the anesthetic. The gas is still widely used by dentists as a
safe and relatively pleasant way of helping patients endure
the discomfort of drilling and other dental procedures.

  The most common commercial method of producing
nitrous oxide involves the controlled heating of ammonium
nitrate (NH4NO3). The compound decomposes to form nitrous
oxide and water. The reaction is essentially the same one
originally used by Priestley in 1772. Although an efficient
means of producing the gas, the reaction must be carried out
with extreme care as ammonium nitrate has a tendency to
decompose explosively when heated. Nitrous oxide can also
be produced by the decomposition of nitrates (compounds
containing the NO3 radical), nitrites (compounds containing
the NO2) radical, or nitriles (compounds containing the CH
radical).

  Nitrous oxide is best known and most widely used as an
anesthetic. Its use is limited primarily to dental procedures
and minor surgeries. Dentists favor nitrous oxide as an anesthetic
because the gas does not make patients completely
unconscious and does not require an anesthesiologist to
administer it. Nitrous oxide works as an anesthetic by blocking
neurotransmitter receptors in the brain, preventing pain
messages from being transmitted.
  Nitrous oxide is also used as a fuel additive in racing
cars, in which case it is often referred to as nitro. The gas is
injected into the intake manifold where it mixes with air and
fuel vapors. Since it breaks down at the high temperatures in
the car’s engine, it provides additional oxygen to increase the

efficiency with which the fuel burns. During World War II,
pilots used nitrous oxide for a similar purpose in their airplanes.
Some additional uses of nitrous oxide include:
• As a propellant in food aerosols;
• For the detection of leaks;
• As a packaging gas for potato chips and other snack
foods, preventing moisture from making the product
become stale;
• In the preparation of other nitrogen compounds; and
• As an oxidizing agent for various industrial processes.
  Nitrous oxide is safe to use in moderate amounts under
controlled conditions. Some people use the compound as a
recreational drug, however, hoping to get a ‘‘high’’ from inhaling
it. One risk of this practice is that the inhalation of
nitrous oxide may reduce the amount of oxygen a person
receives. Also, some long-term health effects, such as anemia
(low red blood cell count) and neuropathy (damage to the
nerves), have been associated with excessive use of the compound.
The use of nitrous oxide for recreational purposes is a
crime in some states.