Chemical & Chemical Engineering News (80th Anniversary Issue), Vol. 81, No. 36, 2003, Sept. Edited by X. Lu Introduction

Fluorochemicals touch millions of peoples' lives in meaningful ways every day, especially in health care and quality-of-life products. At the top of a very long list: the inorganic fluorides used in drinking water and dental products, the one out of every five active pharmaceutical products that is fluorinated, and the synthetic blood substitutes and inhalation drug delivery systems that use fluorocarbons.

Fluorochemicals also underscore a wide range of commercial successes. Growth in the industrial and household refrigeration and air conditioning industries is based largely on the use of low-toxicity, nonflammable, and energy-efficient fluorocarbon fluids. Fluoropolymers and fluoroelastomers are used widely in homes, buildings, automobiles, aerospace applications, and wherever high performance is required--performance such as excellent thermal, flame, electrical, chemical, and solvent resistance and low oxygen and moisture permeability. Other low-molecular-weight perfluoroalkyl-based materials provide oil-, water-, and soil-repellent surface properties for textile, fiber, and paper coatings; and similar materials are used as surfactants to stabilize aqueous fire-fighting foams. Fluorocarbons are also used as fire extinguishants in aerospace and other critical areas. Modern high-energy-density lithium-ion batteries used in handheld electronic devices rely on LiPF₆.

And there's even more. Consider, for example, agrochemicals, where about 17% of active materials used as pesticides and fungicides are based on fluorinated compounds. The manufacture of silicon chips relies on the wet and dry etch processes utilizing materials such as ultra-high-purity HF and NF₃. Elemental fluorine is used to prepare UF₆, used by the nuclear industry for uranium enrichment. Electrical utilities rely on the high dielectric strength of SF₆, also manufactured using direct elemental fluorination, for high-voltage circuit breakers and transformers. And within the chemical processing industries, catalysts include BF₃ and SbF₅, KF is used as a fluorinating agent, and AlF₃ is used to process aluminum.

Today, the innovative use and application of fluorine chemistry and fluorinated materials continues unbounded, especially in growth areas of optoelectronics, electronics, life sciences, and high-performance materials. And what's on the horizon is even more exciting: Imagine "smart dust" where supermicrosensors allow us to gather vital data on a scale previously unimaginable; new materials capable of making more effective repairs to the human body; and pharmaceuticals customized for use, compatibility, and delivery.

It's hard not to get caught up in the future potential of fluorine chemistry. Recently, following a speech I gave to an industry group, a chemist half-jokingly told me, "I'm glad to see the fluorine (F) pin on your jacket lapel, because without it, you're not fulfilling your role as a leader in the chemical industry." I couldn't agree more. In the 117 years since it was first isolated by Henri Moissan, fluorine has become a powerful foundation for chemical exploration, discovery, and innovation. Who would have imagined that fluorine would serve as the foundation for so many of today's modern marvels and tomorrow's most promising innovations?

Nance Dicciani is president and CEO of Specialty Materials, a strategic business group of Honeywell. She holds a Ph.D. in chemical engineering and has more than 25 years of experience in the chemical industry.

FLUORINE AT A GLANCE

Name: From the Latin fluere, flow.

Atomic mass: 19.00.

History: Fluorine was first identified by Karl W. Scheele in 1771. It was first isolated in 1886 by French chemist F. Henri Moissan. His reward was the Nobel Prize for Chemistry in 1906.

Occurrence: Found in fluorspar, cryolite, and other minerals.

Appearance: Pale yellow gas. The free element has a characteristic odor.

Behavior: Fluorine is the most reactive element and reacts with practically all substances. It reacts with water to produce oxygen and ozone. Fluorine gas is corrosive and toxic.

Uses: An essential trace element for mammals. Fluorine and its compounds are used in producing uranium and more than 100 commercial fluorochemicals. Certain fluorocarbons were used in air conditioning and refrigeration but have been phased out because they damaged the ozone layer.

CHLORINE

CYNTHIA BURROWS, UNIVERSITY OF UTAH

As an anion, chlorine is rather ordinary. Sodium grabs all the press from NaCl as the dietary culprit in hypertension, while silver's the key to AgCl's action in photography. Arguably, chloride plays a more important role in the antitumor drug cisplatin, Pt(NH₃)₂Cl₂, although its function there is to get lost (that is, hydrolyze) so that platinum has room to bind to DNA. Sadly, the –1 oxidation state of chlorine is often glossed over as the necessary counterion to an exotic metal ion or complex cation, the necessary yin to complement the yang.

Conversely, molecular chlorine, Cl₂, has held a starring role in history, both for its benefits to human health and for its detrimental effects on the environment. Carl Wilhem Scheele, a Swedish pharmacist, first described the greenish yellow gas in 1774 after dropping hydrochloric acid onto manganese dioxide. Sir Humphry Davy recognized the gas as an element in 1810 and named it based on the Greek word for its color, khloros.

Chlorine was by this time already in use. In the small town of Javelle, France, chlorine added to alkaline water created l'eau Javelle ("bleach" in English, NaOCl in chemspeak) that was used in the fabric industry in the late-18th century.

SHORE LEAVE Morton Salt maintains a plant near the south end of Utah's Great Salt Lake.

The mid-19th century saw one of the most dramatic improvements in human health: Bleach began to be used as a disinfectant in hospitals, and chlorination of the water supply in London during a cholera outbreak in 1850 saved many lives. Chlorine continues to be the disinfectant of choice in the food industry, in swimming pools, and in the drinking water supply in most developed countries.

The Chlorine Chemistry Council argues that the element also has an enormous economic impact, contributing 2 million jobs and around $50 billion to the annual U.S. economy in one way or another. Only a fraction of that is in the form most consumers would easily recognize: household bleach and swimming pool chemicals. We tend to overlook the fact that the C in PVC (polyvinyl chloride) pipes is chloride, and without chlorine we wouldn't have Saran wrap, nylon, microprocessors, soccer balls, or plastic toys. Even less obvious to most is the role of chlorine in the wood and paper industry (as a bleach) and in the processing of metals and the production of other materials such as titanium dioxide. Chlorocarbon compounds range from the good (chloroquine, an antimalarial) to the bad (DDT and chlorofluorocarbons) to the downright ugly (polychlorinated biphenyls). All are synthesized by chlorination of hydrocarbon precursors.

Where does all this chlorine come from? I can literally see tons of it out my window. Elemental chlorine does not exist naturally on our planet but is manufactured by electrolysis of seawater. The vast deposits of salt created during millions of years of continental upheaval and slow evaporation of the ancient Lake Bonneville are mined on the shores of the present-day Great Salt Lake. Through elaborate extraction procedures, the various chloride salts can be separated. Some of this salt ends up on your french fries (NaCl), and some you throw on your sidewalk in the winter (CaCl₂). The MgCl₂ is electrolyzed to produce Mg⁰, a lightweight metal used in the auto industry. Of course, the by-product of magnesium production is elemental chlorine, which can be responsibly used for all of the above-mentioned health and manufacturing applications.

The dark side of Cl₂ production is that too much of it is released directly into the atmosphere. According to the Environmental Protection Agency's Toxics Release Inventory, the biggest U.S. point source of atmospheric Cl₂ is 50 miles upwind of my house on the western shore of Great Salt Lake. Magnesium Corp. of America released 42 million lb of Cl₂ into the skies of Utah's West Desert in 2000, about 90% of the U.S. total for that year.

A little bit of chlorine is a great way to kill bacteria, but higher concentrations turn Dr. Jekyll to Mr. Hyde. More than twice as dense as air, chlorine can settle to the ground as it did in Ypres, France, in April 1915, accounting for thousands of fatalities. Responsible use of chlorine will ensure its continued applications toward improvement of human health and lifestyle without waging war on the environment.



Cynthia Burrows is a professor of chemistry at the University of Utah in Salt Lake City and senior editor of ACS's Journal of Organic Chemistry. She and her family enjoy camping and rockhounding in Utah's mineral-rich West Desert.

CHLORINE AT A GLANCE

Name: From the Greek khloros, greenish yellow.

Atomic mass: 35.45.

History: Discovered, yet misidentified as a compound, by Swedish chemist Carl Wilhelm Scheele in 1774. Identified as an element by Sir Humphry Davy in 1810.

Occurrence: Found in nature dissolved in salts in seawater and in the deposits of salt mines. Today, most chlorine is produced through the electrolysis of aqueous sodium chloride.

Appearance: Yellowish-green, dense, sharp-smelling gas.

Behavior: Liquid chlorine burns skin, and gaseous chlorine irritates mucous membranes. Breathing high concentration of the gas can be fatal; chlorine was used as a poison gas during World War I.

Uses: Chlorine is essential to living systems and is also one of the top chemicals manufactured in the U.S. for commercial uses. It is an excellent disinfectant for swimming pools and water supplies, and its compounds are used in plastics such as polyvinyl chloride (PVC), stain removers, and bleach. Sodium chloride is common table salt.

BROMINE

ARI GREENSPAN, P'TIL TEKHELET

My amazement at the bromine atom stems neither from a sophisticated analysis in the lab nor from its wide commercial applications, but from the increasing awareness of freezing cold water trickling into my wet suit while scuba diving. I was in search of a brilliant blue dye lost for 1,300 years in the sands of the sea.

People have always wanted to look great in clothing, but this was difficult to accomplish in ancient times. Almost all dyes then were plant based, which meant that colored fabrics would eventually fade. There were, in fact, only three dyes known in antiquity to be extremely permanent and intense. These were Tyrian purple (Argaman in the language of the Bible), royal blue (Tekhelet), and scarlet (Tola'at Shani). While scarlet was derived from an insect, both purple and blue were extracted from a snail, and the bromine atom plays a fascinating role in the creation of these colors.

Back then, Tyrian purple adorned the clothing of priests and kings, and was adored by the multitudes for its intensity and permanence. If you had even a stripe of purple on your garment, you would certainly be noticed in an otherwise drab sea of brownish, greenish, yellowish wraps. In ancient Rome, emperors were said to "take the purple upon themselves" as they dressed in royal togas dyed completely purple, and the historian Pliny cites "the mad lust for purple" at that time. But by the middle of the 4th century, if you weren't the caesar or one of his cadre, you could be put to death for wearing any purple at all. Thankfully, things changed and the fashion industry was born, all because a little snail in the Mediterranean had the ability to take bromine from the sea and bind it to indigo, forming dibromoindigo--Tyrian purple.

BIBLICAL Though extract from the Murex trunculus snail is yellow, the dibromoindigo dye is responsible for both the brilliant purple and blue of antiquity.

The permanence of this snail dye is astounding. I traveled to the organic repository of the Israeli Antiquities Authority to see a purple ball of wool from Qumran, dating back to Roman times. As dusty boxes of bone and parchment were pushed aside, a magnificent tuft of Murex-dyed wool was revealed--as vibrant today as the day it was dyed.

The Bible commands that a blue string be worn on the corners of the Israelites' prayer shawls, a blue dyed with mollusk extract. The process of making blue dye for this ancient Jewish rite was lost as a result of Roman edicts and restrictions over 1,300 years. But my buddies and I were determined to rediscover its secret. On a blustery fall day, we donned our wet suits and air tanks to search for the small, slow Murex trunculus snail that houses the exquisite chemistry of nature's art.

As we rose from the deep with a cache of 150 snails, in a cove near the Crusader fortress of Akko, I felt like a link in the chain of history in the quest for the biblical blue dye. The Arab children crowding around our hoard of snails were shouting their word for mollusk, "chilzun, chilzun." I trembled as I realized that their "chilzun" echoed the Talmud's Aramaic name for the creature--"chilazon." With great excitement we began to extract dye from the snail.

Our first glimpse of this proud bromine-based dye from hoary antiquity revealed a humble, clearish yellowish substance. Exposure to the air triggered a complex enzymatic reaction that transformed the liquid through the entire color spectrum until, within minutes, before our very eyes, the single drop of dye was an intense deep purple.

But if the snail we collected was identical to the one described by the rabbis over two millennia ago, why did it not produce the proscribed blue for our fringes? Why were we seeing only purple? The amazing answer to this conundrum, which baffled 20th-century scientists for decades, was discovered in the chemistry lab. In order to use this odoriferous dye, the snail extract must be reduced to achieve a solution. When this process is performed indoors, the result is a purple dye. But if, while in its reduced state, the dibromoindigo is exposed to the sun for a few minutes, the bromine invisibly breaks away from the molecule, leaving behind only indigo, the brilliant biblical blue.

The Talmud equated the color of the Tekhelet dye to the color of the depths of the ocean and heights of the sky. We now understand how the chemistry of a lowly sea snail and the exalted bromine atom yield a world rich in color, complexity, and permanent beauty.

Ari Greenspan is a dentist practicing in Jerusalem and the director of the P'til Tekhelet. (http://www.tekhelet.com). His interests range from biblical archaeology to medieval painted glass and from blacksmithing to gold-leaf illumination.

BROMINE AT A GLANCE

Name: From the Greek bromos, stench.

Atomic mass: 79.90.

History: Although it was first produced by a young German chemistry student, bromine's discovery is credited to French chemist Antoine-Jérôme Balard, who published a paper on it in 1826.

Occurrence: Occurs in seawater, underground salt mines, and deep brine wells, as well as some minerals.

Appearance: The only nonmetal that is a liquid at room temperature. Reddish-brown in color.

Behavior: Very volatile and extremely toxic. Bromine can cause severe burns on skin and its noxious vapors can irritate the nose and throat.

Uses: Mostly used as silver bromide in photographic film. Bromine was once used primarily in producing a leaded gasoline additive, ethylene dibromide, that prevents lead compounds from accumulating in engines, but the increased use of unleaded gas lowered demand for the additive. Also used in fire retardants, tear gas, fumigants, disinfectants, and pesticides.

IODINE

PETER J. STANG,UNIVERSITY OF UTAH

I first encountered iodine as a young child during the second World War, when my mother applied a mysterious purple-brown solution--with quite a sting--to disinfect a bruised, scraped knee. Little did I know then what iodine was or that nearly a half century later it would become one of my favorite elements.

My next encounter with iodine occurred as a teenager in Hungary, when I acquired it in the neighborhood pharmacy as part of my extensive home chemistry set and mixed it in small amounts with self-made smokeless gunpowder to impress my friends with its purple vapor.

Iodine, element 53 with a relative atomic mass of 126.90447, was first isolated by Bernard Courtois in 1811 from the ash of seaweed (by treating kelp with H₂SO₄). It was named by J. L. Gay Lussac in 1813, and its name derives from the Greek word iodes, meaning "violet-colored," reflecting the characteristic lustrous, deep purple color of resublimed crystalline iodine as well as the color of its vapor. Potassium iodide (KI) was used as a remedy for goiter (Derbyshire neck), an enlargement of the thyroid gland, as early as 1819. The thyroid is responsible for the production of thyroxine, a metabolism-regulating hormone. Iodine is an essential trace element for humans and plays an important role in many biological organisms. In modern times, KI is recommended for the treatment of radiation poisoning.