Saccharin - Nonnutritive Sweeteners

Saccharin. A noncaloric sweetener that is about 300 times as sweet as sugar. The compound is manufactured on a large scale in several countries. It is made as saccharin, sodium saccharin, and calcium saccharin, as shown by formulas below.

Saccharin (ortho-benzosulfimide) was discovered in 1879 by I. Remsen and C. Fahlberg when they were researching the oxidation products of toluene sulfone amide. The most common forms of saccharin are sodium and calcium saccharin, although ammonium and other salts have been prepared and used to a very limited extent. The saccharins are white, crystalline powders, with melting points between 226 and 230◦C (438.8 and 446◦F). Soluble in amyl acetate, ethyl acetate, benzene, and alcohol; slightly soluble in water, chloroform, and ether. Saccharin is derived from a mixture of toluenesulfonic acids. They are converted into the sodium salts, then distilled with phosphorus trichloride and chlorine to obtain the orthotoluene sulfonyl chloride, which by means of ammonia is converted into ortho-toluenesulfamide. This is oxidized with permanganate, then treated with acid, and saccharin is crystallized out. In food formulations, saccharin is used mainly in the form of its sodium and calcium salts. Sodium bicarbonate may be added to provide improved water solubility.

Saccharin is used in conjunction with aspartame in carbonated beverages. Other uses include tabletop sweeteners, dry beverage blends, canned fruits, gelatin desserts, cooked and instant puddings, salad dressings, jams, jellies, preserves, and baked goods. For many years, saccharin has been under investigation by a number of countries. As of the late 1900s, some questions remained unresolved.

Niacin

Niacin (NYE-uh-sin) is a B vitamin (vitamin B3) that is essential to cell metabolism. It occurs in two forms, nicotinic acid and nicotinamide, also called niacinamide. The only structural difference between the two compounds is that a hydroxyl group (-OH) in nicotinic acid is replaced by an amino group (-NH2) group in nicotinamide. Lack of niacin causes a disease called pellagra. Pellagra was common throughout human history among poor people whose diet consisted almost entirely of corn products. Those corn products did not supply adequate amounts of niacin, causing symptoms such as diarrhea, scaly skin sores, inflamed mucous membranes, weakness, irritability, and mental delusions. In some cases, people with niacin deficiency develop reddish sores and rashes on their faces. Mental hospitals were full of people who seemed crazy, but who were actually suffering from a dietary deficiency. Thousands of people died from pellagra every year. Nicotinic acid was first isolated by the Polish-American biochemist Casimir Funk (1884–1967) in 1912. At the time,

Funk was attempting to find a cure for another dietary disease known as beriberi. Since nicotinic acid had no effect on beriberi, he abandoned his work with that compound. It was left, then, to the Austrian-American physician Joseph Goldeberger (1874–1929) to find the connection between nicotinic acid and deficiency diseases. In 1915, Goldberger conducted a series of experiments with prisoners in a Mississippi jail and found that he could produce pellagra by altering their diets. He concluded that the disease was caused by the absence of some factor, which he called the P-P (for pellagra-preventative) factor. The chemical structure of that factor was then discovered in 1937 by the American biochemist Conrad Arnold Elvehjem (1901–1962), who cured the disease in dogs by treating them with nicotinic acid.

Niacin is synthesized naturally in the human body beginning with the amino acid tryptophan. Tryptophan occurs naturally in a number of foods, including dairy products, beef, poultry, barley, brown rice, fish, soybeans, and peanuts. People whose diet consists mainly of corn products do not ingest adequate amounts of tryptophan, so their bodies are unable to make the niacin they need to avoid developing pellagra. It takes about 60 milligrams of tryptophan to produce 1 mg of niacin.

Niacin plays a number of essential roles in the body. It is necessary for cell respiration; metabolism of proteins, fats, and carbohydrates; the release of energy from foods; the secretion of digestive enzymes; the synthesis of sex hormones; and the proper functioning of the nervous system. It is also involved in the production of serotonin, an essential neurotransmitter in the brain. Niacin deficiency disorders occur as the result of an inadequate diet, consuming too much alcohol, and among people with certain types of cancer and kidney diseases. Physicians treat niacin deficiency diseases by prescribing supplements of 300 to 1,000 milligrams per day of the vitamin. Overdoses of niacin can cause a variety of symptoms, including itching, burning, flushing, and tingling of the skin.

The Benefit of Vitamin B3

· Required for energy metabolism, enzyme reactions, skin and nerve health, and digestion.

· High doses of nicotinic acid (3 g daily) can lower cholesterol (reduce LDL and triglycerides and increase HDL) and reduce the risk of heart attack and stroke; high dosages should be supervised by a physician.

· Defi ciency causes pellagra, the symptoms of which are skin rash, diarrhea, dementia, and death.

· Defi ciency may be caused by poor diet, malabsorption diseases, dialysis, and HIV.

· Drugs that deplete vitamin B3: antibiotics, isoniazid, and 5-Fluorouracil (chemotherapy).

· High-dose niacin, taken along with statin drugs (i.e., lovastatin), may increase the risk of rhabdomyolysis (muscle degeneration and kidney disease).

· Most people get adequate niacin from diet and/or a multivitamin; supplements may be recommended for those with high cholesterol.

Acetylsalicylic Acid : The Aspirin

   Acetylsalicylic acid (uh-SEE-till-sal-in-SILL-ik As-id, or uhse-
TEEL-sal-ih-SEEL-ik AS-id), more commonly known as
aspirin, is the world’s most commonly used therapeutic drug.
By one estimate, about 137 million aspirin tablets are taken
every day throughout the world. The drug is also known by
other names including: o-acetoxybenzoic acid; 2-(acetyloxy)-
benzoic acid; 2-carboxyphenyl acetate; and benzoic acid,
2-hydroxyacetate, in addition to about ten other systematic
names and many common names.
   The analgesic properties of willow tree bark, from which
salicylic acid comes, have been known for well over 3,500
years. They were first described in Egyptian scrolls dating to
about 1550 BCE and were later recommended by a number of
ancient authorities, including the famous Greek physician
Hippocrates (c. 460–370 BCE), the Roman encyclopedist Aulus
Cornelius Celsus (c. 10 BCE–date of death unknown), the
Roman philosopher Pliny the Elder (23 CE–CE), and the Greek
physician Pedanius Dioscorides (40–90 CE).

   In the period from 1828 to 1829, the active ingredient in
willow bark was first isolated by three individuals, the German
pharmacist Johann Bu¨chner (dates not available), the
French chemist Henri Leroux (dates not available), and the
the Italian chemist, Raffaele Piria (1815–1865). Bu¨chner gave
the name salicin to the bitter-tasting yellow crystals
extracted from willow bark after the Latin name for the
willow tree, Salix. In 1853, the French chemist Charles Frederick
Gerhardt (1816–1857) developed a method for reacting
salicylic acid (the active ingredient in salicin) with acetic
acid to make the first primitive form of aspirin.
   For many years the way aspirin works in the body was not
understood. Scientists now know that the compound’s helpful
effects come from its action on prostaglandins. Prostaglandins
are hormone-like substances released by cells that are injured.
They cause the body to release other substances that sensitize
nerve endings to pain and start the healing process. Aspirin
blocks prostaglandin production, thus relieving the sensation
of pain and the inflammation that are the body’s response to
injury. Aspirin reduces fever by acting on the region of the
brain that regulates body temperature and heart rate. Prostaglandins
block the body’s natural system for producing heat
so that by blocking the release of prostaglandins, aspirin
allows the regulation of body temperature to continue as
usual. Aspirin’s protection against heart attack and stroke
occur because of its effect on one special type of prostaglandin,
known as thromboxane A2. Thromboxane A2 promotes
the accumulation of cells that takes place when a blood clot
forms. By blocking or slowing down the production of thromboxane
A2, aspirin prevents the formation of blood clots and,
hence, the probability of heart attack and stroke.

   The modern method for making aspirin was developed in
1897 by the German chemist Felix Hoffman (1868–1946), an
employee of the German chemical manufacturer Bayer AG
Chemical Works. In this procedure, phenol (C6H5OH) is treated
with sodium hydroxide and carbon dioxide to make salicylic
acid. The salicylic acid is then reacted with acetic acid
(CH3COOH) to make acetylsalicylic acid, or aspirin. The preparation
of aspirin by this procedure is quite simple and is often
assigned to students in beginning high school and college chemistry
classes. Aspirin tablets themselves include only acetylsalicylic
acid, to which is added a small amount of water, starch
and lubricant that act as a binder to hold the tablet together.

The exclusive use of aspirin is as a medicine. It has three
important properties as a drug. It relieves pain, reduces
inflammation, and reduces fever. In addition to its effectiveness
in treating these medical symptoms, it is inexpensive
and available in a variety of forms, including chewable
tablets, extended-release formulations, effervescent tablets,
and even in chewing gums. Aspirin is often prescribed in low,
daily doses as a preventative measure for individuals at risk
for heart attack and stroke.
  While aspirin has many medical benefits, it is not without
risk for some individuals. Some people are allergic to the
compound and can not tolerate even a low dose. Such individuals
experience a number of symptoms if they ingest high
doses of aspirin, symptoms that include ringing in the ears,
nausea, vomiting, dizziness, confusion, hallucinations, coma,
seizures, rapid breathing, fever, and, in the most severe
cases, death. Aspirin use is not recommended in children
under the age of twelve who show symptoms of viral infections
because it can lead to an extremely rare but deadly
complication known as Reye’s syndrome.

Isoprene

Isoprene (EYE-so-preen) is a clear, colorless, volatile
liquid that is both very flammable and quite explosive. It is
classified as a diene compound because its molecules contain
two (‘‘di-’’) double bonds (‘‘-ene’’). It is also a member of the
terpene family. The terpenes are a large family of organic
compounds that contain two or more isoprene units. An
example of a terpene is vitamin A, whose molecular formula
is C20H30O. Vitamin A contains four isoprene units. The
terpenes occur abundantly in nature in both plants and
animals.

Some common terpenes include geraniol, found in geraniums;
limonene, oil of orange; a-pinene, or oil of turpentine;
a-farnesene, oil of cintronella; zingiberene, oil of ginger;
farnesol, found in lily of the valley; b-selinene, oil of celery;
and caryophyllene, oil of cloves. Isoprene is also produced in
animal bodies and is said to be the most common hydrocarbon
present in the human body. By one estimate, a 70-kilogram
(150-pound) person produces about 17 milligrams of
isoprene per day. Probably the best-known source of isoprene
is natural rubber, which is a polymer consisting of long
chains of isoprene units joined to each other.

  A number of methods are available for preparing isoprene
from petroleum. Perhaps the most common process is
the cracking of hydrocarbons present in the naphtha portion
of refined petroleum. Cracking is the process by which large
hydrocarbons are broken down into smaller hydrocarbons
either with heat or over a catalyst, or by some combination
of heat and catalyst. The naphtha portion of petroleum consists
of hydrocarbons with boiling points between about
50C and 200C (120F and 400F). Other methods for the
preparation of isoprene include the dehydrogenation
(removal of hydrogen) of isopentene (CH3CH(CH3)CH=CH2),
the pyrolysis (decomposition by high heat) of methylpentene
(CH2=C(CH3)CH2CH2CH3), or the dehydration (removal of
water) of methylbutenol (CH3C(CH3)(OH)CH2CH3).

Natural rubber has been known to humans for hundreds
of years. Archaeologists have found that the Indians of
South and Central America were making rubber products
as early as the eleventh century. Until the end of the nineteenth
century, natural supplies of rubber obtained from
the rubber tree, Hevea brasiliensis, were sufficient to meet
consumer demand for the product. However, with the development
of modern technology—especially the invention of
the automobile—natural supplies of the product proved to
Interesting Facts

• Isoprene and other terpenes are now known
to undergo reactions that contribute to the development
of pollutants, such as ozone and oxides
of nitrogen in the atmosphere.
• Isoprene is a key intermediary in the synthesis
of cholesterol in the human body.
• The production of isoprene by plants seems to be
associated with the process of photosynthesis
and is affected by temperature, sunlight, other gases, and other
factors.
• The polymer of isoprene is called polyisoprene. It
exists in two forms, cis- and trans-polyisoprene. The two
forms are called geometric isomers. They have the
same kind and number of atoms, but the atoms
are arranged differently in the two forms. Natural
rubber consists of transpolyisoprene, while another product found in rubber plants, gutta percha, is made of cis-polyisoprene.

be insufficient to meet growing demand. Chemical researchers
began to look for ways of producing synthetic forms of
rubber.

One approach was to attempt making synthetic rubber
with exactly the same chemical composition as that of natural
rubber, that is, a polymer of trans-polyisoprene. As early
as the 1880s, British chemist Sir William Augustus Tilden
(1842–1926) was successful in achieving this objective. Tilden
found that he could make isoprene by heating turpentine
(C10H16 ). The isoprene then polymerized easily when exposed
to light. After more than twenty years of research, however,
Tilden decided that synthetic trans-polyisoprene could never
be made economically, and he encouraged his friends to
forget about the process.
Over the years, chemists did find ways of making other
types of synthetic rubber, and some never abandoned the
effort to make synthetic trans-polyisoprene. The critical
breakthrough needed in this research occurred in about
1953 when Swiss chemist Karl Ziegler (1898–1973) and Italian
chemist Giulio Natta (1903–1979) each found a way of
polymerizing isoprene in such a way that its geometric
structure matched that of natural rubber exactly. A year
later, chemists at two of the largest rubber companies in
the world, B. F. Goodrich and Firestone, announced that they
had developed methods for making synthetic trans-polyisoprene
using essentially the methods developed earlier by
Ziegler and Natta.
In the early twenty-first century, more than 95 percent
of the isoprene produced is used to make trans-polyisoprene
synthetic rubber. The remaining 5 percent is used to make
other types of synthetic rubber and other kinds of polymers.
A small amount of the compound is used as a chemical
intermediary, a substance from which other organic chemicals
is made.
Isoprene is a dangerous fire hazard. It also poses a risk to
human health and that of other animals. It is an irritant to
skin, eyes, and the respiratory system. Upon exposure, it
produces symptoms such as redness, watering, and itching
of the eyes and itching, reddening, and blistering of the skin.
If inhaled, it can irritate the lungs and respiratory system.
Isoprene is a known carcinogen.

Pectin

Pectin (PEK-tin) is a mixture, not a compound. Mixtures
differ from compounds in a number of important ways. The
parts making up a mixture are not chemically combined with
each other, as they are in a compound. Also, mixtures have
no definite composition, but consist of varying amounts of
the substances from which they are formed.
  Chemically, pectin is a polysaccharide, a very large molecule
made of many thousands of monosaccharide units joined
to each other in long, complex chains. Monosaccharides are
simple sugars. The most familiar monosaccharide is probably
glucose, the sugar from which the human body obtains
the energy it needs to grow and stay healthy. The monosaccharides
in pectin are different from and more complex than
glucose.

  Pectin occurs naturally in many fruits and vegetables.
It is most abundant in citrus fruits such as lemons, oranges,
and grapefruits, which may consist of up to 30 percent pectin.
In pure form it is a yellowish-white powder with virtually no
odor and a slightly gummy taste. When dissolved in water, it
forms a thick, jelly-like mass. This property explains one of
its primary purposes: the jelling of fruits when they are made
into jams and jellies.
  Pectin is made naturally in ripening fruit. It is obtained
commercially by treating the raw material (citrus peel or
apple pomace) with hot, acidified water. (Apple pomace is
the residue remaining after pressing of apples.) The pectin
in the peel or apple pomace dissolves in the hot water and is
then purified by repeated filtrations. It is extracted from the
water solution by adding alcohol or an aluminum salt to the
solution, causing the pectin to precipitate out of solution.
The precipitate is then dried and ground into a powder.
Additional steps are sometimes carried out to convert the
pectin produced by this method, called high ester pectin, to a
form that is more soluble: low ester pectin. To achieve this

Hydrogen Peroxide

   

   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.

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.

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).

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.

Ascorbic Acid (The Vitamin C)

  Ascorbic acid (as-KOR-bik AS-id), or vitamin C, is one of
the most important dietary vitamins for humans because it
plays a crucial role in building collagen, the protein that
serves as a support structure for the body. It is a watersoluble
vitamin, which means that the body excretes any
excess vitamin C in the urine and cannot store a surplus.
For that reason, humans must consume vitamin C in their
daily diets. Vitamin C is found in many fruits and vegetables
and most kinds of fresh meat. Citrus fruits, such as oranges
and lemons, are especially rich in the compound.

Humans have known about the consequences of vitamin C
deficiency for centuries. People traveling long distances
on land or by sea often came down with an illness called
scurvy. The same illness struck people living in their own
homes during long winters. The disease was characterized by
pain and weakness in the joints, fatigue, bleeding gums,
tooth loss, slow healing of wounds, and bruising. These symptoms
were caused as the body’s connective tissue broke down

and small blood vessels ruptured. These symptoms began to
disappear as fresh foods became more available. If they did
not get enough fresh food in their diets, people could die of
scurvy.
Scurvy was common enough that many people searched
for its cause and cure. Sailors were especially vulnerable to

the disease, and the first recorded investigations involving
vitamin C were done by seafaring men. In 1536, French
explorer Jacques Cartier (1491–1557) cured his sailors of
scurvy by following the advice of Indians in Newfoundland,
feeding them extract of pine needles. Scottish physician
James Lind (1716–1794) began investigating the disease in
1747. He read many historical accounts of the diseases and
combined that information with his own observations to
deduce that scurvy occurred only among people with very
limited diets. He went on a ten-week sea voyage and fed the
solders various foods to see which ones were best at curing
scurvy. Citrus fruits proved to be most effective in preventing
the disease, a result that Lind reported in 1753. Captain
James Cook (1728–1779) led expeditions to the South Seas in
the late 1700s and kept his crew healthy by feeding them
sauerkraut. In 1795 the British navy began serving its sailors
a daily portion of lime juice, and two things happened: British
sailors stopped getting scurvy, and people began calling
sailors ‘‘limeys.’’
  Many people refused to believe that scurvy was caused by
a dietary deficiency, suggesting that it was instead the result
of eating bad food or lack of exercise. In 1907, Norwegian
biochemists Alex Holst (1861–1931) and Theodore Frohlich
conducted a study in which guinea pigs were fed an experimental
diet that caused them to develop scurvy. The link
between the vitamin and the disease was firmly established
by this research. Ascorbic acid was first isolated independently
by the Hungarian-American biochemist Albert Szent-
Gyo¨rgi (1893–1986) and the American biochemist Charles
Glen King (1896–1988) in 1932. It was synthesized a year
later by the English chemist Sir Walter Norman Haworth
(1883–1950) and the Polish-Swiss chemist Tadeusz Reichstein
(1897–1996), again working independently of each other.

  Plants and most animals (humans and guinea pigs being
two exceptions) synthesize vitamin C in their cells through a
series of reactions in which the sugar galactose is eventually
converted to ascorbic acid. For many years, the compound has
been made commercially by a process known as the Reichstein
process, named after its inventor Tadeusz Reichstein. This
process begins with ordinary glucose, which is converted to
another sugar, sorbitol, which is then fermented to obtain

yet another sugar, sorbose. The sorbose is then converted
step-by-step into a series of other products, the last of which
is ascorbic acid.
  Chemists have long been searching for an alternative to
the Reichstein process because it uses so much energy and
produces by-products that are hazardous to the environment.
In the 1960s, Chinese scientists developed a method
that involves only two steps in the synthesis of ascorbic
acid, and in the early 2000s, Scottish scientists were
attempting to develop a method that involved only a single
step using fermentation. Currently, however, the Reichstein
process remains the most popular method for making the
compound.

  The best known use of vitamin C is as a nutritional
supplement, taken to ensure that one receives his or her
daily minimum requirement of the vitamin. The recommended
daily allowance (RDA) of vitamin C for adults is 60
milligrams per day. Anyone who eats a well-balanced diet
that includes citrus fruits, tomatoes, and green leafy vegetables
probably does not need to take a vitamin supplement.
However, the amount of vitamin C one normally receives
from a supplement is unlikely to cause any harm.

  In addition to its nutritional uses, ascorbic acid has a
number of other industrial applications, including:
• As a food preservative;
• As a reducing agent in chemical processes;
• As a preservative in foods;
• As a color fixing agent in meats, helping meats keep
their bright red appearance;
• As an additive to bread dough, where it helps increase
the activity of yeast used in the dough; and
• As a treatment for abscission in citrus plants, the tendency
for a plant to lose its leaves, flowers, and fruits.

Petroleum

  Petrolatum (peh-tro-LAY-tum) is a mixture, not a compound.

Mixtures differ from compounds in a number of

important ways. The parts making up a mixture are not

chemically combined with each other, as they are in a compound.

Also, mixtures have no definite composition, but

consist of varying amounts of the substances from which

they are formed.

   Petrolatum is a complex mixture of hydrocarbons

derived from the distillation of petroleum. Hydrocarbons

are compounds that contain only carbon and hydrogen. The

hydrocarbons that make up petrolatum belong to the

methane (saturated or alkane) family of hydrocarbons with

the general formula CnH2N+2. Some members of the family

include methane (CH4), ethane (C2H5), propane (C3H8), and

butane (C4H10).

   Petrolatum occurs in a semi-solid or liquid form. The

semi-solid form is also called petroleum jelly or mineral jelly

and is commercially available under a number of trade

names, including Kremoline, Pureline, Sherolatum, and

VaselineTM. It ranges in color from white to yellowish to

amber. It is practically odorless and tasteless. It melts over a

wide range, from about 38 C to about 55 C (100 F to 131 F).

The liquid form is also known as liquid paraffin, mineral

oil, or white mineral oil. Such products are sold commercially

under trade names such as Alboline, Drakeol, Frigol,

Kremol, and Paroleine. It is a colorless, tasteless, and odorless

oily liquid.

   

   Oil was first discovered in the United States in the 1850s

in western Pennsylvania. A chemist from Brooklyn, New

York, Robert Augustus Chesebrough (1837–1938), visited

the new wells and noticed a wax-like material sticking

to the petroleum drilling rods. He learned that oil workers

used the ‘‘rod wax’’ to heal burns on their skin. Chesebrough

eventually extracted and purified the substance—petrolatum—

from petroleum and began manufacturing it in 1870. He

received several patents for his discovery and in 1878, he

gave his product the trade name of VaselineTM. His product

quickly became popular as an ointment for wounds and

burns. Unlike the animal and vegetable oils then being used

for that purpose, petrolatum did not spoil. By the late 1870s,

VaselineTM was selling at the rate of one jar everyminute in

the United States. In 1880, it was added to the U.S. Pharmacopoeia,

a manual that lists drugs used in medical practice.

   Petrolatum is a product of the fractional distillation of

crude oil. Crude oil is a complex mixture of hundreds or

thousands of compounds. These compounds can be separated,

or distilled, from each other by heating crude oil to high

temperatures. As the temperature of the crude oil rises,

various groups or a ‘‘fraction’’ of compounds boil off. The first

group of compounds includes gaseous compounds dissolved

in crude oil. The next group of compounds includes compounds

with low boiling points. The next group of compounds

includes compounds with slightly higher boiling

points. And so on. Eventually, a tar-like mass of compounds

with very high boiling points is left behind in the distilling

tower. This residue is heated to separate liquids from solids

remaining behind. Some of these liquids and solids make up

the semi-solid and liquid forms of petrolatum.

  

   Petrolatum has a wide variety of uses, ranging from

personal care and medical applications to industrial uses.

The solid form, such as VaselineTM is used as a topical ointment

for the treatment of dry, cracked skin and to reduce the

risk of infection. It works as a moisturizing agent because it

reduces water loss from the skin, It helps prevent infection

because it creates a barrier over wounds that prevents disease-

causing organisms from entering the body. Solid petrolatum

is also an ingredient in many skin care and cosmetic

products, such as skin lotions, body and facial cleansers, antiperspirants,

lipsticks, lip balms, sunscreens, and after-sun

lotions. In hair products, it helps smooth frizzy hair by

allowing hair to retain its natural moisture. The formation

used in most of these products remains virtually unchanged

from that developed by Robert Chesebrough in the 1800s.

  

  Solid petrolatum is also used in industrial applications

for a variety of purposes, such as:

• As a softener in the production of rubber products;

• In the food processing industry, to coat raw fruits and

vegetables and to help products retain moisture;

• As a defoaming agent in the production of beet sugar

and yeasts;

• For the lubrication of firearms and machine parts;

• In the production of modeling clays;

• In the manufacture of candles, to prevent a candle from

shrinking as it cools after being burned;

• In the preparation of shoe polishes; and

• As an ingredient in rust preventatives.

   The primary use of liquid petrolatum is as a laxative, a

product that loosens the bowels. It also has a number of other

applications, such as an additive in foods such as candies,

confectionary products, and baked goods; as an ingredient in

personal care products, such as baby oil creams, hair conditioning

lotions, and ointments; in many different kinds of

pharmaceutical preparations; in the production of industrial

lubricants; as a softening agent in the manufacture of rubber,

textiles, fibers, adhesives, and machine parts; as dust

suppressants; and as dehydrating agents for a number of

industrial processes.

Sodium Hydroxide

  Sodium hydroxide (SO-dee-um hye-DROK-side) is a white
deliquescent solid commercially available as sticks, pellets,
lumps, chips, or flakes. A deliquescent material is one that
absorbs moisture from the air. Sodium hydroxide also reacts
readily with carbon dioxide in the air to form sodium carbonate.
Sodium hydroxide is the most important commercial caustic. A
caustic material is a strongly basic or alkaline material that
irritates or corrodes living tissue. The compound ranked number
11 among chemicals produced in the United States in 2004.

  Sodium hydroxide is produced commercially simultaneously
with chlorine gas by the electrolysis of a sodium
chloride solution. In this process, an electric current breaks
down sodium chloride into its component elements, sodium
and chlorine. The chlorine escapes as a gas, while the sodium
metal form reacts with water to form sodium hydroxide

2NaCl ! 2Na + Cl2
2Na + 2H2O ! 2NaOH + H2

  Sodium hydroxide can also be produced easily by means
of other chemical reactions. For example, the reaction
between slaked lime (calcium hydroxide; Ca(OH)2) and soda
ash (sodium carbonate; Na2CO3) produces sodium hydroxide:
Ca(OH)2 + Na2CO3 ! 2NaOH + CaCO3
None of these alternative methods can compete economically,
however, with the preparation by electrolysis.

Sodium hydroxide has a great variety of household and
industrial uses. It is the active ingredient in drain cleaners
such as Drano because it breaks up and dissolves the greasy
mass that is responsible for drain blockages. It is also an
ingredient in many other household products, including oven
cleaners, metal polishes, and hair straighteners. Sodium
hydroxide is also used in the preparation of homemade and
processed foods. It is used in the preparation of soft drinks,
chocolate, ice creams, caramel coloring, and cocoa. Hominy, a
starchy food similar to grits, is made by soaking corn kernels
in a solution of sodium hydroxide in water. Bakers glaze
pretzels and German lye rolls with a weak lye solution before
baking them. The lye gives baked goods a crisp crust. Some
people use lye to cure olives.
  The largest single use for sodium hydroxide is in the
production of organic compounds from which polymers are
made, such as propylene oxide and the ethylene amines, and
of the polymers themselves, including the polycarbonates
and epoxy resins. About a third of all the sodium hydroxide
produced in the United States goes to this application.
Another important use of sodium hydroxide is in the pulp
and paper industry, where it is used to digest (break down)
the raw materials from which pulp and paper are made.
About 13 percent of all the sodium hydroxide made in the

  United States goes to this application. Sodium hydroxide is
also an important raw material in the manufacture of soap.
The method by which soap is made has not changed very
much for thousands of years. A fat or oil is added to a boiling
solution of sodium hydroxide in water. The fat or oil hydrolyzes
into its component parts, glycerol and fatty acids. The
sodium hydroxide then reacts with the fatty acids, forming
sodium salts. The sodium salt of a fatty acid is a soap. Sodium
hydroxide is also an important raw material in the manufacture
of inorganic compounds, especially sodium and calcium
hypochlorite, sodium cyanide, and a number of sulfur-containing
compounds. Some other important uses of sodium
hydroxide include:

• In the manufacture of cellophane and rayon;
• As a neutralizing agent during the refining of petroleum;
• In the manufacture of aluminum metal;
• For the refining of vegetable oils;
• As an agent for peeling fruits and vegetables for processing;
• In the extraction of metals from their ores;
• For the processing of textiles;
• In water treatment facilities;
• For etching and electroplating operations; and
• In a wide variety of research laboratory applications.

  Sodium hydroxide is one of the most caustic substances
known and a strong irritant to the skin, eyes, and respiratory
system. Exposure to sodium hydroxide dust, powder, or solid
can cause burning of the skin and eyes, with possible permanent
damage to one’s vision. Ingestion of the compound

causes burning of the mouth, esophagus, and stomach, resulting
in nausea, diarrhea, internal bleeding, scarring, and permanent
damage to the lungs and gastrointestinal system.
More serious results, such as a drop in blood pressure and
collapse, are also possible.

Caffeine

  Caffeine (kaf-EEN) is an organic base that occurs naturally
in a number of plant products, including coffee beans,
tea leaves, and kola nuts. It occurs as a fleecy white crystalline
material, often in the form of long, silky needles. It
usually exists as the monohydrate, C8H10N4O2 H2O, although
it gives up its water of hydration readily when exposed
to air.
  Scientists believe that humans have been drinking beverages
that contain caffeine for thousands of years. The first
recorded reference to a caffeine drink can be found in a
Chinese reference to the consumption of tea by the emperor
Shen Nung in about 2700 BCE. Coffee is apparently a much
more recent drink, with the earliest cultivation of the coffee
tree dated at about 575 CE in Africa.
  Caffeine was first studied scientifically by two French
chemists, Joseph Bienaime´ Caventou (1795–1877) and Pierre
Joseph Pelletier (1788–1842), who were very interested in
the chemical properties of the alkaloids. Between 1817
and 1821, Caventou and Pelletier successfully extracted
caffeine, quinine, strychnine, brucine, chinchonine, and
chlorophyll (not an alkaloid) from a variety of plants. The
first synthesis of caffeine was accomplished in 1895 by the
German chemist Emil Hermann Fischer (1852–1919), who was
awarded the 1902 Nobel Prize in chemistry for his work on
the alkaloids.

  Caffeine belongs to a class of alkaloids called the methylxanthines.
Chocolate, from the cocoa tree Theobroma cacao
contains another member of the class, theobromine. Both
caffeine and theobromine are stimulants, that is, compounds
that act on the nervous system to produce alertness, excitement,
and increased physical and mental activity.
  Caffeine can be extracted from coffee, tea, and kola
plants by one of three methods. These methods are used
primarily to produce the decaffeinated counterparts of the
products: decaffeinated coffee, decaffeinated tea, or decaffeinated
soft drinks. A commercial variation of these procedures
is to treat the waste products of tea or coffee processing, such
as the dust and sweepings collected from factories, for the
extraction of caffeine.
  In the first of the three extraction methods, the natural
product (coffee beans, tea leaves, or kola beans) are treated
with an organic solvent that dissolves the caffeine from the
plant material. The solvent is then evaporated leaving behind
the pure caffeine. A second method follows essentially the
same procedure, except that hot water is used as the solvent
for the caffeine. A more recent procedure involves the use of
supercritical carbon dioxide for the extraction process.
  Supercritical carbon dioxide is a form of the familiar gas
that exists at high temperature and high pressure. It behaves
as both a liquid and a gas. Not only is the supercritical carbon
dioxide procedure an efficient method of extracting caffeine,
but it has virtually none of the harmful environmental and
health problems associated with each of the other two methods
of extraction.
  Caffeine is also made synthetically by heating a combination
of the silver salt of theobromine (C7H8N4O2Ag) with
methyl iodide (CH2I), resulting in the addition of one carbon
and two hydrogens to the theobromine molecule and converting
it to caffeine.
  Caffeine is used in foods and drinks and for medical
purposes. Its primary action is to stimulate the central nervous
system. People drink coffee, tea, or cola drinks to stay
awake and alert because caffeine creates a feeling of added
energy. It does this by increasing heart rate, improving blood
flow to the muscles, opening airways to aid breathing, and
releasing stored energy from the liver to provided added fuel
for the body. In large quantities, caffeine can also cause
nervousness, insomnia, and heart problems. The effects of
caffeine can linger in the body for more than six hours. In
medical applications, caffeine is sometimes used as a heart
stimulant for patients in shock, to treat apnea (loss of breathing)
in newborn babies, to counteract depressed breathing
levels as a result of drug overdoses, and as a diuretic.
  Caffeine stimulates the brain in two ways. First, because
it has a chemical structure similar to that of adenosine, it
attaches to adenosine receptors in the brain. Adenosine is a
substance that normally attaches to those receptors, slowing
brain activity and causing drowsiness. By blocking those
receptors, caffeine increases electrical activity in the brain,
creating a feeling of alertness. Caffeine also works in the
brain like drugs such as heroin and cocaine, although in a
much milder way. Like those drugs, caffeine increases dopamine
levels. Dopamine is a chemical present in the brain that
increases the body’s feeling of pleasure.
Studies have shown that caffeine can become addictive.
People who use the compound eventually need to take more
and more of it to get the same effect. When some people try to
stop using caffeine, they may suffer from headache, fatigue,
and depression, though these symptoms can be controlled
by gradually reducing the amount of caffeine consumed.
Either way, withdrawal symptoms end after about a week.

Glucose

  Glucose (GLOO-kose) is a simple sugar used by plants and
animals to obtain the energy they need to stay alive and to
grow. It is classified chemically as a monosaccharide, a compound
whose molecules consist of five- or six-membered
carbon rings with a sweet flavor. Other common examples
of monosaccharides are fructose and galactose. Glucose
usually occurs as a colorless to white powder or crystalline
substance with a sweet flavor. It consists in two isomeric
forms known as the D configuration and the L configuration.
Dextrose is the common name given to the D conformation of
glucose.
Credit for the discovery of glucose is often given to the
German chemist Andreas Sigismund Marggraf (1709–1782).
In 1747, Marggraf isolated a sweet substance from raisins
that he referred to as einer Art Z cker (a kind of sugar) that
we now recognize as glucose. More than 60 years later, the
German chemist Gottlieb Sigismund Constantine Kirchhof
(1764–1833) showed that glucose could also be obtained from
the hydrolysis of starch and that starch itself was nothing
other than a very large molecule (polysaccharide) composed
of many repeating glucose units. The molecular structure
for glucose was finally determined in the 1880s by German
chemist Emil Fischer (1852–1919), part of the reason for
which he was awarded the 1902 Nobel Prize in chemistry.
  Glucose is synthesized naturally in plants and some single-
celled organisms through the process known as photosynthesis.
In this process, sunlight catalyzes the reaction
between carbon dioxide and water that results in the formation
of a simple carbohydrate (glucose) and oxygen. The overall
reaction can be summarized by a rather simple chemical
equation:
6CO2 + 6H2O ! C6H12O6 + 6O2
However, photosynthesis actually involves a number of
complex reactions that occur in two general phases, the light
reactions and the dark reactions.
Glucose is produced commercially through the steam
hydrolysis of cornstarch or waste products containing cellulose
(a large molecule composed of glucose units) using a
dilute acid catalyst. The product thus obtained is typically
not very pure, but is contaminated with maltose (a disaccharide
consisting of two molecules of glucose joined to each
other) and dextrins (larger molecules consisting of a number
of glucose units joined to each other).

  Glucose is the primary chemical from which plants and
animals derive energy. In cells, glucose is broken down in a
complex series of reactions to produce energy with carbon
dioxide and water as byproducts.
  Glucose also has a number of commercial uses, nearly all
of them related to the food processing business. It is used in
the production of confectionary products; chewing gum; soft
drinks; ice creams; jams, jellies, and fruit preparations; baby
foods; baked products; and beers and ciders. A relatively small
amount is used for non-food purposes, primarily in the production
of other organic chemicals, such as citric acid, the amino
acid lysine, insulin, and a variety of antibiotics.
The most important health problem associated with
glucose is diabetes. Diabetes is a medical condition that
develops when the body either does not produce adequate
amounts of insulin or cannot use that compound properly.
Insulin is a hormone that controls the metabolism of glucose
in the body. If glucose is not metabolized properly, a
person’s body acts as if it is ‘‘starving.’’ Symptoms of diabetes
include excessive hunger, weight loss, and exhaustion.
If left untreated, the condition can result in coma
and death. Diabetics must have an artificial source of insulin
(usually from injections) and watch their diets to keep
these symptoms under control.

Urea

Urea (yoo-REE-uh) is a white crystalline solid or powder
with almost no odor and a salty taste. It is a product of the
decomposition of proteins in the bodies of terrestrial animals.
Urea is produced in the liver and transferred to the
kidneys, from which it is excreted in urine. The compound
was first identified as a component of urine by French chemist
Hilaire Marin Rouelle (1718–1799) in 1773. It was first
synthesized accidentally in 1828 by German chemist Friedrich
Wo¨hler (1800–1882). The synthesis of urea was one of the
most important historical events in the history of chemistry.
It was the first time that a scientist had synthesized an
organic compound. Prior to Wo¨hler’s discovery, scientists
believed that organic compounds could be made only by the
intervention of some supernatural force. Wo¨hler’s discovery
showed that organic compounds were subject to the same set
of natural laws as were inorganic compounds (compounds
for non-living substances). For this reason, Wo¨hler is often
called the Father of Organic Chemistry.

The formation of urea is the evolutionary solution to the
problem of what to do with poisonous nitrogen compounds
that formed when proteins decompose in the body. Proteins
are large, complex compounds that contain relatively large
amounts of nitrogen. When they decompose, that nitrogen is
converted to ammonia (NH3), a substance that is toxic to
animals. If animals are to survive the decomposition of proteins
(as happens whenever foods are metabolized), some
method must be found to avoid the buildup of ammonia in
the body.

  That method involves a series of seven chemical reactions
called the urea cycle by which nitrogen from proteins

is converted into urea. Although high concentrations of urea
do pose a risk to animal bodies, the urea formed in these
reactions is normally excreted fast enough to avoid health
problems for an animal.
Urea is produced commercially by the direct synthesis
of liquid ammonia (NH3) and liquid carbon dioxide (CO2).
The product of this reaction is ammonium carbamate
(NH4CO2NH2):
2NH3 + CO2 ! NH4CO2NH2
Ammonia and carbon dioxide do not react with each
other under normal conditions of temperature and pressure.
If the pressure is raised to 100 to 200 atmospheres (1750 to
3000 pounds per square inch) and the temperature is raised
to about 200C (400C), however, the reaction proceeds efficiently
with the formation of ammonium carbamate. When
the pressure is then reduced to about 5 atmosphere (80
pounds per square inch), the ammonium carbamate decomposes
to form urea and water:
NH4CO2NH2 ! (NH2)2CO + H2O

  Urea is the sixteenth most important chemical in the
United States, based on the amount produced annually. In
2004, the chemical industry produced 5.755 million metric
tons (6.344 million short tons) of urea. Almost 90 percent of
that output was used in the manufacture of fertilizers. An
additional 5 percent went to the production of animal feeds.
In both fertilizers and animal feeds, urea and the compounds
from which it is made provide the nitrogen needed by growing
plants and animals for their good health and survival.
The other major use of urea is in the manufacture of various
types of plastics, especially urea-formaldehyde resins and
melamine.
Urea is also used:
• In the production of personal care products, such as
hair conditioners, body lotions, and dental products;
• In certain pharmaceutical and medical products, such as
creams to treat wounds and damaged skin;
• As a stabilizer in explosives, a compound that places
limits on the rate at which an explosion proceeds;
• In the manufacture of adhesives;
• For the flame-proofing of fabrics;
• For the separation of products produced during the
refining of petroleum;
• In the production of sulfamic acid (HOSO2NH2), an
important raw material in many chemical processes;
• As a coating for paper products; and
• In the production of deicing agents.

Penicillin

  The penicillins (pen-uh-SILL-ins) are a class of antibiotic
compounds derived from the molds Penicillium notatum and
Penicillium chrysogenum. The class contains a number of
compounds with the same basic bicyclic structure to which
are attached different side chains. That basic structure consists of two amino acids, cysteine and valine, joined to each
other to make a bicyclic (‘‘two-ring’’) compound. The different
forms of penicillin are distinguished from each other by
adding a single capital letter to their names. Thus: penicillin
F, penicillin G, penicillin K, penicillin N, penicillin O, penicillin
S, penicillin V, and penicillin X. A number of other
antibiotics, including ampicillin, amoxicillin, and methicillin,
have similar chemical structures.

  Penicillin was discovered accidentally in 1928 by the
Scottish bacteriologist Alexander Fleming (1881–1995). Fleming
noticed that a green mold, which he later identified as
Penicillium notatum, had started to grow on a petri dish that
he had coated with bacteria. As the bacteria grew towards the mold, they began to die. At first, Fleming saw some promise
in this observation. Perhaps the mold could be used to kill
the bacteria that cause human disease. His experiments
showed, however, that the mold’s potency declined after a
short period of time He was also unable to isolate the antibacterial
chemical produced by the mold. He decided that
further research on Penicillium was probably not worthwhile.

As a result, it was not until a decade later that Penicillium’s
promise was realized. In 1935, English pathologist
Howard Florey (1898–1968) and his biochemist colleague
Ernst Chain (1906–1970) came across Fleming’s description
of his experiment and decided to see if they could isolate the
chemical product produced by Penicillium with anti-bacterial
action. They were eventually successful, isolating and purifying
a compound with anti-bacterial action, and, in 1941, began trials
with human subjects to test its safety and efficacy (ability to kill
bacteria). The successful conclusion of those trials not only
provided one of the great breakthroughs in the human battle
against infectious diseases, but also won for Florey, Chain, and
Fleming the 1945 Nobel prize for Physiology or Medicine.

Penicillins are classified as biosynthetic or semisynthetic.
Biosynthetic penicillin is natural penicillin. It is produced by culturing molds in large vats and collecting and
purifying the penicillins they produce naturally. There are
six naturally occurring penicillins. The specific form of penicillin
produced in a culturing vat depends on the nutrients
provided to the molds. Of the six natural penicillins, only
penicillin G (benzylpenicillin) is still used to any extent.
  Semi-synthetic penicillins are produced by making chemical
alterations in the structure of a naturally occurring
penicillin. For example, penicillin V is made by replacing the
-CH2C6H5 group in natural penicillin G with a -CH2OC6H5
group.

  Penicillins are prescription medications used to treat a
variety of bacterial infections, including meningitis, syphilis,
sore throats, and ear aches. They do so by inactivating an
enzyme used in the formation of bacterial cell walls. With the
enzyme inactivated, bacteria can not make cell walls and die
off. Penicillins do not act on viruses in the same way they do
on bacteria, so they are not effective against viral diseases,
such as the flu or the common cold.
A number of side effects are related to the use of penicillin.
These side effects include diarrhea, upset stomach, and
vaginal yeast infections. In those individuals who are allergic to penicillins, side effects are far more serious and include
rash, hives, swelling of tissues, breathing problems, and
anaphylactic shock, a life-threatening condition that requires
immediate medical treatment.
Penicillin may alter the results of some medical tests,
such as those for the presence of sugar in the urine. Penicillin
can also interact with a number of other medications,
including blood thinners, thyroid drugs, blood pressure
drugs, birth control pills, and other antibiotics, in some cases
decreasing their effectiveness.
  Once promoted as wonder drugs, the use of penicillins
has declined slowly because of the spread of antibiotic resistance.
Antibiotic resistance occurs when new strains of bacteria
evolve that are resistant to existing types of penicillin.
One reason that antibiotic resistance has become a problem
is the extensive and often unnecessary use of penicillins.
When they are prescribed for colds and the flu, for example,
they have no effect on the viruses that cause those diseases,
but they encourage the growth of bacteria more able to
survive against penicillins.

Saccharin - Sweet like Sugar

   Saccharin (SAK-uh-rin) is a synthetic compound whose
water solutions are at least 500 times as sweet as table sugar.
It passes through the human digestive system without being
absorbed, so it has an effective caloric value of zero. It is
used as a sugar substitute by diabetics or by anyone wishing
to reduce their caloric intake.
Saccharin was the first artificial sweetener discovered. It
was synthesized accidentally in 1879 when Johns Hopkins
researchers Constantine Fahlberg (1850–1910) and Ira Remsen
(1846–1927) were working on the development of new food
preservatives. The story is told that Fahlberg accidentally
spilled one of the substances being studied on his hand. Some
time later, he noticed the sweet taste of the substance and
began to consider marketing the product as an artificial sweetener.
Fahlberg and Remsen jointly published a paper describing
their work, but Fahlberg, without Remsen’s knowledge,
went on to request a patent for the discovery. He eventually
became very wealthy from proceeds of the discovery, none of

which he shared with Remsen. Remsen was later quoted as
saying that ‘‘Fahlberg is a scoundrel. It nauseates me to hear
my name mentioned in the same breath with him.’’
   which he shared with Remsen. Remsen was later quoted as
saying that ‘‘Fahlberg is a scoundrel. It nauseates me to hear
my name mentioned in the same breath with him.’’
   of an artificial sweetener. By 1902, it had become so popular
in Germany that the German sugar industry lobbied for laws
limiting production of saccharin. Similar actions occurred in
the United States in 1907 and, by 1911, the federal government
restricted use of the compound to overweight invalids.
A shortage of sugar during World War I (1914–1918) led
to the reintroduction of saccharin as a sweetening agent in
foods. Another sugar shortage during World War II (1939–
1945) saw a new boom in saccharin production. This time, the
compound’s popularity continued after the war ended.
Questions about saccharin’s safety have been raised a
number of times in the past. In 1969, for example, the sweetener
cyclamate was found to be carcinogenic, and its use was
banned in the United States. Doubts over saccharin’s safety
resurfaced, partly since it was often mixed with cyclamate in
artificial sweeteners. In 1972, studies with rats suggested
that saccharin too might be carcinogenic, and the U.S. Food
and Drug Administration imposed restrictions on the sweetener’s
use. Later studies attempting to reproduce the 1972
research were largely unsuccessful, and the status of saccharin
as a carcinogen remain unsettled.
In 1977, the Canadian government decided that sufficient
evidence existed to ban the use of saccharin except for use
with diabetics and others with special medical problems. The
U.S. government considered taking similar action, but, after
more than a decade of reviewing the evidence, decided to
allow the use of saccharin among the general public. Nonetheless,
the status of saccharin as a potential health hazard
remains the subject of an active debate in the United States
and other parts of the world.
  
   A number of methods are available for the synthesis of
saccharin. For many years, the most popular process was one
developed by the Maumee Chemical Company of Toledo,
Ohio, in 1950. This method begins with anthranilic acid
(o-aminobenzoic acid; C6H4(NH2)COOH), which is treated successively
with nitrous acid (HNO2), sulfur dioxide (SO2),
chlorine (Cl2), and ammonia (NH3) to obtain saccharin.
Another process discovered in 1968 starts with o-toluene,
which is then treated with sulfur dioxide and ammonia to
obtain saccharin.
   Saccharin is not very soluble, so it is commonly made
into its sodium or calcium salt (sodium saccharin or calcium
saccharin), both of which readily dissolve in water, for use in
drinks and cooking. Saccharin is often blended with other
sweeteners to reduce its metallic aftertaste.

   Saccharin is used almost exclusively as an artificial
sweetener in food and drinks to replace sugar. Its lack of
calories makes it suitable for diet products and for medical
preparations designed for people who must reduce their caloric
intake. It also finds some small application as a food
preservative, as an antiseptic agent, and as a brightening
agent in electroplating procedures.
Raw saccharin can be an irritant to the skin, eyes, and
respiratory system. If ignited, it burns with the release of
irritating fumes. Only individuals who come into contact

with large quantities of saccharin are likely to be concerned
about such safety problems, however.