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.

Beta-Carotene

Beta-carotene (b-carotene; BAY-tuh KARE-oh-teen) belongs to a family of organic compounds called the carotenoids. The carotenoids are all brightly pigmented (colored) compounds found in a number of plants, bacteria, algae, and fungi. Betacarotene is responsible for the yellowish to orange color of pumpkins, apricots, sweet potatoes, nectarines, and, most notably, carrots. The compound also occurs in spinach and broccoli, but in such small concentrations that the green chlorophyl presentmasks the orange color of beta-carotene. In its pure form, beta-carotene occurs as purple crystals shaped like thin leaflets.

In plants, algae, and photosynthetic bacteria, beta-carotene plays an important role in photosynthesis, the process by which plants convert water and carbon dioxide into carbohydrates and oxygen. In nonphotosynthetic bacteria and fungi, beta-carotene protects the organism against the harmful effects of light and oxygen.

Animals require beta-carotene for normal growth and development, but are unable to manufacture the compound themselves. As a result, they must ingest some beta-carotene from plant sources in order to stay healthy. The compound is a provitamin, a substance that is converted in the body to a vitamin. Beta-carotene is converted into vitamin A, whose role in the body is the maintenance of strong bones and teeth and healthy skin and hair. Beta-carotene also acts as an antioxidant, a substance that attacks free radicals in the body that may cause cancer. It may also protect against heart disease and strengthen the body’s immune system.

Beta-carotene was first isolated by the German chemist Heinrich Wilhelm Ferdinand Wackenroder (1789–1854), who extracted the compound from carrot roots in 1831. The compound was first synthesized in 1950 by the Swiss chemist Paul Karrer (1889–1971).

Beta-carotene can be obtained from natural sources by crushing or pulverizing the source (such as carrots) and adding a solvent that will dissolve the organic components of the plant. These components can then be separated from each other by chromatographic techniques. A major commercial source of beta-carotene obtained by this method is the algae Dunaliella salina, which grows in large salt lakes in Australia. The compound can also be prepared synthetically by one of two methods, the BASF and the Roche methods, both named after the pharmaceutical firms where they were developed. Both methods of preparation begin with long-chain hydrocarbons containing about twenty carbon atoms each. These hydrocarbons are then joined to each other to form the 40-carbon beta-carotene compound.

Beta-carotene has two uses: in vitamin supplements and as a food additive. Anyone who eats a healthy diet that includes foods rich in vitamin A, such as fish oil, liver, eggs, butter, and orange or yellow vegetables and fruits, will get adequate amounts of beta-carotene. However, many people take vitamin supplements to ensure that they have enough beta-carotene (as well as other vitamins) in their daily diet. Although some warnings have been issued about taking too much vitamin A, there is no clinical evidence that an overdose of the vitamin does any long-term harm to a person.

Beta-carotene is used as a food additive to increase the color intensity of a product. It is used primarily with yellow and orange foods, such as butter and margarine, although it is sometimes added to ice cream and fruit juices as well. Beta-carotene is used in only very small amounts as a food additive. In these amounts, it poses no health hazard to humans or other animals. The compound has also been used in experiments to test its effectiveness against certain diseases, such as lung cancer. In such cases, it has been found to be more harmful than beneficial, increasing the risk of cancer and death among people participating in the studies.

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.