Table salt, bleach, fluoride in toothpaste, chlorine in swimming pools—what do all of these have in common? Add halogen lamps to the list, and the answer becomes more clear: all involve one or more of the halogens, which form Group 7 of the periodic table of elements. Known collectively by a term derived from a Greek word meaning "salt-producing," the halogen family consists of five elements: fluorine, chlorine, bromine, iodine, and astatine. The first four of these are widely used, often in combination; the last, on the other hand, is a highly radioactive and extremely rare substance. The applications of halogens are many and varied, including some that are dangerous, controversial, and deadly.
HOW IT WORKS
The Halogens on the Periodic Table
As noted, the halogens form Group 7 of the periodic table of elements. They are listed below, along with chemical symbol and atomic number:
- Fluorine (F) 9
- Chlorine (Cl): 17
- Bromine (Br): 35
- Iodine (I): 53
- Astatine (At): 85
On the periodic table, as displayed in chemistry labs around the world, the number of columns and rows does not vary, since these configurations are the result of specific and interrelated properties among the elements. There are always 18 columns; however, the way in which these are labeled differs somewhat from place to place. Many chemists outside the United States refer to these as 18 different groups of elements; however, within the United States, a somewhat different system is used.
In many American versions of the chart, there are only eight groups, sometimes designated with Roman numerals. The 40 transition metals in the center are not designated by group number, nor are the lanthanides and actinides, which are set apart at the bottom of the periodic table. The remaining eight columns are the only ones assigned group numbers. In many ways, this is less useful than the system of 18 group numbers; however, it does have one advantage.
ELECTRON CONFIGURATIONS AND BONDING.
In the eight-group system, group number designates the number of valence electrons. The valence electrons, which occupy the highest energy levels of an atom, are the electrons that bond one element to another. These are often referred to as the "outer shell" of an atom, though the actual structure is much more complex. In any case, electron configuration is one of the ways halogens can be defined: all have seven valence electrons.
Because the rows in the periodic table indicate increasing energy levels, energy levels rise as one moves up the list of halogens. Fluorine, on row 2, has a valence-shell configuration of 2s22p5; while that of chlorine is 3s23p5. Note that only the energy level changes, but not the electron configuration at the highest energy level. The same goes for bromine (4s24p5), iodine (4s24p5), and astatine (5s25p5).
All members of the halogen family have the same valence-shell electron configurations, and thus tend to bond in much the same way. As we shall see, they are inclined to form bonds more readily than most other substances, and indeed fluorine is the most reactive of all elements.
Thus it is ironic that they are "next door" to the Group 8 noble gases, the least reactive among the elements. The reason for this, as discussed in the Chemical Bonding essay, is that most elements bond in such a way that they develop a valence shell of eight electrons; the noble gases are already there, so they do not bond, except in some cases—and then principally with fluorine.
Characteristics of the Halogens
In terms of the phase of matter in which they are normally found, the halogens are a varied group. Fluorine and chlorine are gases, iodine is a solid, and bromine is one of only two elements that exists at room temperature as a liquid. As for astatine, it is a solid too, but so highly radioactive that it is hard to know much about its properties.
Despite these differences, the halogens have much in common, and not just with regard to their seven valence electrons. Indeed, they were identified as a group possessing similar characteristics long before chemists had any way of knowing about electrons, let alone electron configurations. One of the first things scientists noticed about these five elements is the fact that they tend to form salts. In everyday terminology, "salt" refers only to a few variations on the same thing—table salt, sea salt, and the like. In chemistry, however, the meaning is much broader: a salt is defined as the result of bonding between an acid and a base.
Many salts are formed by the bonding of a metal and a nonmetal. The halogens are all nonmetals, and tend to form salts with metals, as in the example of sodium chloride (NaCl), a bond between chlorine, a halogen, and the metal sodium. The result, of course, is what people commonly call "salt." Due to its tendency to form salts, the first of the halogens to be isolated—chlorine, in 1811—was originally named "halogen." This is a combination of the Greek words halos, or salt, and gennan, "to form or generate."
In their pure form, halogens are diatomic, meaning that they exist as molecules with two atoms: F2, Cl2, and so on. When bonding with metals, they form ionic bonds, which are the strongest form of chemical bond. In the process, halogens become negatively charged ions, or anions. These are represented by the symbols F−, Cl−, Br−, and I−, as well as the names fluoride, chloride, bromide, and iodide. All of the halogens are highly reactive, and will combine directly with almost all elements.
Due to this high level of reactivity, the halogens are almost never found in pure form; rather, they have to be extracted. Extraction of halogens is doubly problematic, because they are dangerous. Exposure to large quantities can be harmful or fatal, and for this reason halogens have been used as poisons to deter unwanted plants and insects—and, in one of the most horrifying chapters of twentieth century military history, as a weapon in World War I.
Chlorine is a highly poisonous gas, greenish-yellow in color, with a sharp smell that induces choking in humans. Yet, it can combine with other elements to form compounds safe for human consumption. Most notable among these compounds is salt, which has been used as a food preservative since at least 3000 b.c.
Salt, of course, occurs in nature. By contrast, the first chlorine compound made by humans was probably hydrochloric acid, created by dissolving hydrogen chloride gas in water. The first scientist to work with hydrochloric acid was Persian physician and alchemist Rhazes (ar-Razi; c. 864-c. 935), one of the most outstanding scientific minds of the medieval period. Alchemists, who in some ways were the precursors of true chemists, believed that base metals such as iron could be turned into gold. Of course this is not possible, but alchemists in about 1200 did at least succeed in dissolving gold using a mixture of hydrochloric and nitric acids known as aqua regia.
The first modern scientist to work with chlorine was Swedish chemist Carl W. Scheele (1742-1786), who also discovered a number of other elements and compounds, including barium, manganese, oxygen, ammonia, and glycerin. However, Scheele, who isolated it in 1774, thought that chlorine was a compound; only in 1811 did English chemist Sir Humphry Davy (1778-1829) identify it as an element. Another chemist had suggested the name "halogen" for the alleged compound, but Davy suggested that it be called chlorine instead, after the Greek word chloros, which indicates a sickly yellow color.
USES OF CHLORINE.
The dangers involved with chlorine have made it an effective substance to use against stains, plants, animals—and even human beings. Chlorine gas is highly irritating to the mucous membranes of the nose, mouth, and lungs, and it can be detected in air at a concentration of only 3 parts per million (ppm).
The concentrations of chlorine used against troops on both sides in World War I (beginning in 1915) was, of course, much higher. Thanks to the use of chlorine gas and other antipersonnel agents, one of the most chilling images to emerge from that conflict was of soldiers succumbing to poisonous gas. Yet just as it is harmful to humans, chlorine can be harmful to microbes, thus preserving human life. As early as 1801, it had been used in solutions as a disinfectant; in 1831, its use in hospitals made it effective as a weapon against a cholera epidemic that swept across Europe.
Another well-known use of chlorine is as a bleaching agent. Until 1785, when chlorine was first put to use as a bleach, the only way to get stains and unwanted colors out of textiles or paper was to expose them to sunlight, not always an effective method. By contrast, chlorine, still used as a bleach today, can be highly effective—a good reason not to use regular old-fashioned bleach on anything other than white clothing. (Since the 1980s, makers of bleaches have developed all-color versions to brighten and take out stains from clothing of other colors.)
Calcium hydrocholoride (CaOCl), both a bleaching powder and a disinfectant used in swimming pools, combines both the disinfectant and bleaching properties of chlorine. This and the others discussed here are just some of many, many compounds formed with the highly reactive element chlorine. Particularly notable—and controversial—are compounds involving chlorine and carbon.
CHLORINE AND ORGANIC COMPOUNDS.
Chlorine bonds well with organic substances, or those containing carbon. In a number of instances, chlorine becomes part of an organic polymer such as PVC (polyvinyl chloride), used for making synthetic pipe. Chlorine polymers are also applied in making synthetic rubber, or neoprene. Due to its resistance to heat, oxidation, and oils, neoprene is used in a number of automobile parts.
The bonding of chlorine with substances containing carbon has become increasingly controversial because of concerns over health and the environment, and in some cases chlorine-carbon compounds have been outlawed. Such was the fate of DDT, a pesticide soluble in fats and oils rather than in water. When it was discovered that DDT was carcinogenic, or cancer-causing, in humans and animals, its use in the United States was outlawed.
Other, less well-known, chlorine-related insecticides have likewise been banned due to their potential for harm to human life and the environment. Among these are chlorine-containing materials once used for dry cleaning. Also notable is the role of chlorine in chlorofluorocarbons (CFCs), which have been used in refrigerants such as Freon, and in propellants for aerosol sprays. CFCs tend to evaporate easily, and concerns over their effect on Earth's atmosphere have led to the phasing out of their use.
Fluorine has the distinction of being the most reactive of all the elements, with the highest electronegativity value on the periodic table. Because of this, it proved extremely difficult to isolate. Davy first identified it as an element, but was poisoned while trying unsuccessfully to decompose hydrogen fluoride. Two other chemists were also later poisoned in similar attempts, and one of them died as a result.
French chemist Edmond Fremy (1814-1894) very nearly succeeded in isolating fluorine, and though he failed to do so, he inspired his student Henri Moissan (1852-1907) to continue the project. One of the problems involved in isolating this highly reactive element was the fact that it tends to "attack" any container in which it is placed: most metals, for instance, will burst into flames in the presence of fluorine. Like the others before him, Moissan set about to isolate fluorine from hydrogen fluoride by means of electrolysis—the use of an electric current to cause a chemical reaction—but in doing so, he used a platinum-iridium alloy that resisted attacks by fluorine. In 1906, he received the Nobel Prize for his work, and his technique is still used today in modified form.
PROPERTIES AND USES OF FLUORINE.
A pale green gas of low density, fluorine can combine with all elements except some of the noble gases. Even water will burn in the presence of this highly reactive substance. Fluorine is also highly toxic, and can cause severe burns on contact, yet it also exists in harmless compounds, primarily in the mineral known as fluorspar, or calcium fluoride. The latter gives off a fluorescent light (fluorescence is the term for a type of light not accompanied by heat), and fluorine was named for the mineral that is one of its principal "hosts".
Beginning in the 1600s, hydrofluoric acid was used for etching glass, and is still used for that purpose today in the manufacture of products such as light bulbs. The oil industry uses it as a catalyst—a substance that speeds along a chemical reaction—to increase the octane number in gasoline. Fluorine is also used in a polymer commonly known as Teflon, which provides a non-stick surface for frying pans and other cooking-related products.
Just as chlorine saw service in World War I, fluorine was enlisted in World War II to create a weapon far more terrifying than poison gas: the atomic bomb. Scientists working on the Manhattan Project, the United States' effort to develop the bombs dropped on Japan in 1945, needed large quantities of the uranium-235 isotope. This they obtained in large part by diffusion of the compound uranium hexafluoride, which consists of molecules containing one uranium atom and six fluorine anions.
FLUORIDATION OF WATER.
Long before World War II, health officials in the United States noticed that communities having high concentration of fluoride in their drinking water tended to suffer a much lower incidence of tooth decay. In some areas the concentration of fluoride in the water supply was high enough that it stained people's teeth; still, at the turn of the century—an era when dental hygiene as we know it today was still in its infancy—the prevention of tooth decay was an attractive prospect. Perhaps, officials surmised, it would be possible to introduce smaller concentrations of fluoride into community drinking water, with a resulting improvement in overall dental health.
After World War II, a number of municipalities around the United States undertook the fluoridation of their water supplies, using concentrations as low as 1 ppm. Within a few years, fluoridation became a hotly debated topic, with proponents pointing to the potential health benefits and opponents arguing from the standpoint of issues not directly involved in science. It was an invasion of personal liberty, they said, for governments to force citizens to drink water which had been supplemented with a foreign substance.
During the 1950s, in fact, fluoridation became associated in some circles with Communism—just another manifestation of a government trying to control its citizens. In later years, ironically, antifluoridation efforts became associated with groups on the political left rather than the right. By then, the argument no longer revolved around the issue of government power; instead the concern was for the health risks involved in introducing a substance lethal in large doses.
Fluoride had meanwhile gained application in toothpastes. Colgate took the lead, introducing "stannous fluoride" in 1955. Three years later, the company launched a memorable advertising campaign with commercials in which a little girl showed her mother a "report card" from the dentist and announced "Look, Ma! No cavities!" Within a few years, virtually all brands of toothpaste used fluoride; however, the use of fluoride in drinking water remained controversial.
As late as 1993, in fact, the issue of fluoridation remained heated enough to spawn a study by the U.S. National Research Council. The council found some improvement in dental health, but not as large as had been claimed by early proponents of fluoridation. Furthermore, this improvement could be explained by reference to a number of other factors, including fluoride in toothpastes and a generally heightened awareness of dental health among the U.S. populace.
Another controversial application of fluorine is its use, along with chlorine and carbon, in chlorofluorocarbons. As noted above, CFCs have been used in refrigerants and propellants; another application is as a blowing agent for polyurethane foam. This continued for several decades, but in the 1980s, environmentalists became concerned over depletion of the ozone layer high in Earth's atmosphere.
Unlike ordinary oxygen (O2), ozone or O3 is capable of absorbing ultraviolet radiation from the Sun, which would otherwise be harmful to human life. It is believed that CFCs catalyze the conversion of ozone to oxygen, and that this may explain the "ozone hole," which is particularly noticeable over the Antarctic in September and October.
As a result, a number of countries signed an agreement in 1996 to eliminate the manufacture of halocarbons, or substances containing halogens and carbon. Manufacturers in countries that signed this agreement, known as the Montreal Protocol, have developed CFC substitutes, most notably hydrochlorofluorocarbons (HCFCs), CFC-like compounds also containing hydrogen atoms.
The ozone-layer question is far from settled, however. Critics argue that in fact the depletion of the ozone layer over Antarctica is a natural occurrence, which may explain why it only occurs at certain times of year. This may also explain why it happens primarily in Antarctica, far from any place where humans have been using CFCs. (Ozone depletion is far less significant in the Arctic, which is much closer to the population centers of the industrialized world.) In any case, natural sources, such as volcano eruptions, continue to add halogen compounds to the atmosphere.
Bromine is a foul-smelling reddish-brown liquid whose name is derived from a Greek word meaning "stink." With a boiling point much lower than that of water—137.84°F (58.8°C)—it readily transforms into a gas. Like other halogens, its vapors are highly irritating to the eyes and throat. It is found primarily in deposits of brine, a solution of salt and water. Among the most significant brine deposits are in Israel's Dead Sea, as well as in Arkansas and Michigan.
Credit for the isolation of bromine is usually given to French chemist Antoine-Jérôme Balard (1802-1876), though in fact German chemist Carl Löwig (1803-1890) actually isolated it first, in 1825. However, Balard, who published his results a year later, provided a much more detailed explanation of bromine's properties.
The first use of bromine actually predated both men by several millennia. To make their famous purple dyes, the Phoenicians used murex mollusks, which contained bromine. (Like the names of the halogens, the word "Phoenicians" is derived from Greek—in this case, a word meaning "red" or "purple," which referred to their dyes.) Today bromine is also used in dyes, and other modern uses include applications in pesticides, disinfectants, medicines, and flame retardants.
At one time, a compound containing bromine was widely used by the petroleum industry as an additive for gasoline containing lead. Ethylene dibromide reacts with the lead released by gasoline to form lead bromide (PbBr2), referred to as a "scavenger," because it tends to clean the emissions of lead-containing gasoline. However, leaded gasoline was phased out during the late 1970s and early 1980s; as a result, demand for ethylene dibromide dropped considerably.
The name "halogen" is probably familiar to most people because of the term "halogen lamp." Used for automobile headlights, spotlights, and floodlights, the halogen lamp is much more effective than ordinary incandescent light. Incandescent "heat-producing" light was first developed in the 1870s and improved during the early part of the twentieth century with the replacement of carbon by tungsten as the principal material in the filament, the area that is heated.
Tungsten proved much more durable than carbon when heated, but it has a number of problems when combined with the gases in an incandescent bulb. As the light bulb continues to burn for a period of time, the tungsten filament begins to thin and will eventually break. At the same time, tungsten begins to accumulate on the surface of the bulb, dimming its light. However, by adding bromine and other halogens to the bulb's gas filling—thus making a halogen lamp—these problems are alleviated.
As tungsten evaporates from the filament, it combines with the halogen to form a gaseous compound that circulates within the bulb. Instead of depositing on the surface of the bulb, the compound remains a gas until it comes into contact with the filament and breaks down. It is then redeposited on the filament, and the halogen gas is free to combine with newly evaporated tungsten. Though a halogen bulb does eventually break down, it lasts much longer than an ordinary incandescent bulb and burns with a much brighter light. Also, because of the decreased tungsten deposits on the surface, it does not begin to dim as it nears the end of its life.
First isolated in 1811 from ashes of seaweed, iodine has a name derived from the Greek word meaning "violet-colored"—a reference to the fact it forms dark purple crystals. During the 1800s, iodine was obtained commercially from mines in Chile, but during the twentieth century wells of brine in Japan, Oklahoma, and Michigan have proven a better source.
Among the best-known properties of iodine is its importance in the human diet. The thyroid gland produces a growth-regulating hormone that contains iodine, and lack of iodine can cause a goiter, a swelling around the neck. Table salt does not naturally contain iodine; however, sodium chloride sold in stores usually contains about 0.01% sodium iodide, added by the manufacturer.
Iodine was once used in the development of photography: during the early days of photographic technology, the daguerreotype process used silver plates sensitized with iodine vapors. Iodine compounds are used today in chemical analysis and in synthesis of organic compounds.
Just as fluorine has the distinction of being the most reactive, astatine is the rarest of all the elements. Long after its existence was predicted, chemists still had no luck finding it in nature, and it was only created in 1940 by bombarding bismuth with alpha particles (positively charged helium nuclei). The newly isolated element was given a Greek name meaning "unstable."
Indeed, none of astatine's 20 known isotopes is stable, and the longest-lived has a half-life of only 8.3 hours. This has only added to the difficulties involved in learning about this strange element, and therefore it is difficult to say what applications, if any, astatine may have. The most promising area involves the use of astatine to treat a condition known as hyperthyroidism, related to an overly active thyroid gland.
WHERE TO LEARN MORE
"The Chemistry of the Halogens." Purdue University Department of Chemistry (Web site). <http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch10/group7.html> (May 20, 2001).
"Halogens." Chemical Elements.com (Web site). <http://www.chemicalelements.com/groups/halogens.html> (May 20, 2001).
"Halogens." Corrosion Source (Web site). <http://www.corrosionsource.com/handbook/periodic/halogens.htm> (May 20, 2001).
"Halogens" (Web site). <http://registrar.ies.ncsu.edu/ol_2000/module6/halogen/halogen/htm> (May 20, 2001).
"The Halogens" (Web site). <http://www.nidlink.com/~jfromm/elements/halogen.htm> (May 20, 2001).
Knapp, Brian J. Chlorine, Fluorine, Bromine, and Iodine. Henley-on-Thames, England: Atlantic Europe, 1996.
Oxlade, Chris. Elements and Compounds. Chicago: Heinemann Library, 2001.
Stwertka, Albert. A Guide to the Elements. New York: Oxford University Press, 1998.
"Visual Elements: Group VII—The Halogens" (Web site). <http://www.chemsoc.org/viselements/pages/data/intro_groupvii_data.html> (May 20, 2001).
The negative ion that results when an atom gains one or more electrons. An anion (pronounced "AN-ie-un") of an element is never called, for instance, the chlorine anion. Rather, an anion involving a single element is named by adding the suffix -ide to the name of the original element—in this case, "chloride." Other rules apply for more complex anions.
The number of protons in the nucleus of an atom. Since this number is different for each element, elements are listed on the periodic table of elements in order of atomic number.
A one-or two-letter abbreviation for the name of an element.
A term describing an element that exists as molecules composed of two atoms. All of the halogens arediatomic.
The use of an electrical current to cause a chemical reaction.
The length of time it takes a substance to diminish to one-half its initial amount.
Group 7 of the periodic table of elements, including fluorine, chlorine, bromine, iodine, and astatine. The halogens are diatomic, and tend to form salts; hence their name, which comes from two Greek terms meaning "salt-forming."
An atom that has lost or gained one or more electrons, and thus has a net electric charge.
A form of chemical bonding that results from attractions between ions with opposite electricalcharges.
Atoms that have an equal number of protons, and hence are of the same element, but differ in their number of neutrons. This results in a difference ofmass. Isotopes may be either stable or unstable—that is, radioactive.
PERIODIC TABLE OF ELEMENTS:
A chart that shows the elements arranged in order of atomic number. Vertical columns within the periodic table indicate groups or "families" of elements with similar chemical characteristics.
A large molecule containing many small units that hook together.
A term describing a phenomenon whereby certain materials are subject to a form of decay brought about by the emission of high-energy particles. "Decay" in this sense does not mean "rot"; instead, radioactive isotopes continue to emit particles, changing into isotopes of other elements, until they becomestable.
A compound formed by the reaction of an acid with a base. Salts are usually formed by the joining of a metal and a nonmetal.
Electrons that occupy the highest energy levels in anatom, and are involved in chemical bonding. The halogens all have seven valence electrons.
"Halogens." Science of Everyday Things. . Encyclopedia.com. (November 24, 2017). http://www.encyclopedia.com/science/news-wires-white-papers-and-books/halogens-0
"Halogens." Science of Everyday Things. . Retrieved November 24, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/news-wires-white-papers-and-books/halogens-0
The halogens are the five chemical elements that make up Group 17 on the periodic table: fluorine, chlorine, bromine, iodine, and astatine. The term halogen comes from Greek terms meaning "to produce sea salt." The halogens are all chemically active. For that reason, none occur naturally in the form of elements. However—with the exception of astatine—they are very widespread and abundant in chemical compounds. The most widely known of these compounds is sodium chloride, or common table salt.
Fluorine and chlorine are gases. Bromine is one of only two liquid elements. Iodine is a solid. Astatine is radioactive and is one of the rarest of the chemical elements. Fluorine is the most reactive of all known elements. Chemical reactivity decreases throughout this family of elements, with fluorine being the most reactive of all known elements, and chlorine, bromine, and iodine being relatively less reactive, respectively.
Simple compounds of the halogens are called halides. When a halogen becomes part of a compound with one other element, its name is changed to an -ide ending; for example, a chloride.
The name fluorine comes from the name of the mineral in which the element was found, fluorspar. Fluorine was one of the last common elements to be isolated. It is so reactive that chemists searched for more than 70 years to find a way to extract the element from its compounds. Then, in 1886, French chemist Henri Moissan (1852–1907) found a way to produce fluorine by passing an electric current through a liquid mixture of potassium fluoride and hydrogen fluoride. Moissan's method is still used today, with some modifications, for the production of fluorine.
Properties and uses. Fluorine is one of the most dangerous chemicals known. It attacks the skin and throat, causing serious burns and respiratory problems at very low concentrations. It is also very reactive chemically. It attacks most chemicals vigorously at room temperature and reacts explosively with water.
An indication of fluorine's reactivity is that it even forms compounds with the family of elements known as the inert gases. The inert gases include helium, neon, argon, krypton, and xenon. They get their name from the fact that they generally do not combine with any other element. However, compounds of xenon and fluorine and krypton and fluorine have been produced. They are the only known compounds of the inert gases to have been discovered.
Words to Know
Chemical activity: The tendency to form chemical compounds.
Compound: The combination of two or more elements in a definite mass ratio.
Radioactive: The tendency of an element to break down spontaneously into one or more other elements.
Synthesized: Prepared by scientists in a laboratory; not a naturally occurring process.
Because it is so reactive, fluorine itself has few uses. One exception is its role as an oxidizing (burning) agent in rocket fuels. The vast majority of fluorine, however, is used to make compounds. One of the most interesting of those compounds is hydrofluoric acid. This compound has been used since the 1600s to etch glass.
Perhaps the most familiar application of fluorine compounds is in toothpaste additives. Scientists have discovered that the addition of tiny amounts of fluoride in a person's diet can decrease the number of dental caries (cavities) that develop. Today, many kinds of toothpastes include stannous fluoride to improve a person's dental health.
For many years, the most important group of fluorine compounds used commercially were the chlorofluorocarbons (CFCs). The CFCs were developed and used as refrigerants, blowing agents for polyurethane foam, and propellants in spray cans. At one time, more than 700 million kilograms (1.5 billion pounds) of CFCs were produced in a single year.
In the 1980s, however, scientists found that CFCs break down in the atmosphere. The chlorine formed as a result of this breakdown attacks the ozone layer in Earth's stratosphere (the part of Earth's atmosphere that extends 7 to 31 miles [11 to 50 kilometers] above the surface). The loss of the ozone layer is a serious problem for humans since ozone screens out radiation that causes skin cancer and other damage to plants and animals on Earth. Today, scientists are exploring the use of another class of fluorine compounds—the hydrochlorofluorocarbons, or HCFCs—as replacements for CFCs.
Chlorine was first prepared in the 1770s by Swedish chemist Carl Wilhelm Scheele (1742–1786), who thought it was a compound. It was later identified as an element by English chemist Humphry Davy (1778–1829). Davy suggested the name of chlorine for the element because of its greenish-yellow color. (The Greek word for "greenish-yellow" is chloros. )
Chlorine occurs most abundantly in sodium chloride, which is obtained from seawater and from underground deposits of rock salt formed from seas that have dried up. To obtain chlorine, an electrical current is passed through brine, a water solution of sodium chloride.
Properties and uses. Chlorine gas is toxic. It attacks the respiratory tract, causing coughing, congestion, and flu-like symptoms. In high doses, it can be fatal. For this reason, chlorine was used as a chemical weapon during World War I (1914–18).
Chlorine is also very reactive, although less so than fluorine. It forms compounds with almost every other element. Among the most important of those compounds are sodium chloride (table salt), potassium chloride, hydrochloric acid, and calcium chloride.
Chlorine consistently ranks among the top ten chemicals produced in the United States. Some of chlorine's uses depend on its toxic effects. For example, chlorine is now widely used as a disinfectant in municipal water systems, swimming pools, and sewage treatment plants. Many organic (carbon-containing) compounds of chlorine are used as pesticides, herbicides, and fungicides. These compounds kill unwanted insects, weeds, fungi, and other plants and animals. The use of these compounds is often associated with undesirable environmental effects, however. DDT, for example, is a chlorine-containing compound that was once one of the most popular pesticides ever produced. But its harmful effects on fish, birds, and other animals in the environment eventually led to bans on its use in many industrialized nations.
The mention of chlorine brings summertime and swimming pools to mind for most people. Chlorine is added to pools and spas to kill bacteria in water that might otherwise cause disease. The process of adding chlorine to a swimming pool is called chlorination. Chlorination is used for other purification purposes also, as in the purification of public water supplies.
Chlorination can be done in various ways. In some cases, gaseous chlorine is pumped directly into water. In other cases, a compound containing chlorine, such as sodium or calcium hypochlorite, is added to water. When that compound breaks down, chlorine and other purifying substances are released to the water.
The term chlorination applies more generally to any chemical reaction in which chlorine is added to some other substance. For example, chlorine and methane gas can be reacted with each other to form a series of chlorine-containing compounds. The best known of that series are trichloromethane (also known as chloroform; used as an anesthetic) and tetrachloromethane (also known as carbon tetrachloride; used as a solvent and a refrigerant).
Chlorine is also used in the bleaching of paper, pulp, and textiles. The largest single application of the element is in the preparation of a large variety of compounds, including organic chlorides that are the starting point in the manufacture of plastics and other kinds of polymers (chemical compounds that consist of repeating structural units). One of the most important of these polymers is polyvinyl chloride (PVC), from which plastic pipe and many other plastic products are made. Another is neoprene, a synthetic form of rubber that is resistant to the effects of heat, oxidation, and oils. Neoprene is widely used in automobile parts.
Bromine was discovered in 1826 by French chemist Antoine-Jérôme Balard (1802–1876). Balard chose the name bromine from the Greek word for "stink," because of its strong and disagreeable odor. Like chlorine, bromine is obtained from brine. Chlorine gas is used to convert bromide compounds in brine to elemental bromine.
Properties and uses. Bromine is a beautiful reddish-brown liquid that vaporizes (changes to a gas) easily. The vapors are irritating to the eyes and throat. The liquid is highly corrosive and can cause serious burns if spilled on the skin. Bromine is chemically less active than fluorine and chlorine but more active than iodine.
Like chlorine, bromine can be used as a disinfectant. In fact, some water treatment systems have converted from chlorination to bromination as a way of purifying water. For many years, one of the most important compounds of bromine was ethylene dibromide, an additive in leaded gasolines. Since leaded gasoline has been removed from the market, this use has declined.
The product in which most people are likely to encounter compounds of bromine is in photographic film. Tiny crystals of silver bromide undergo a chemical change when exposed to light. This change is responsible for the image produced when photographic film is used to take a picture.
Bromine is also used to make a number of organic products that function as pesticides. The most popular of these currently is methyl bromide, a fumigant (another word for a substance used to destroy pests). Methyl bromide is used as a spray for potatoes, tomatoes, and other agricultural crops.
The halons are a group of organic compounds that contain bromine along with at least one other halogen. The halons are popular as flame retardants. However, scientists have found that, like the CFCs, they appear to cause damage to Earth's ozone layer. For that reason, their use has been largely reduced throughout the world.
Iodine was discovered accidentally in 1811 by French chemist Bernard Courtois (1777–1838). Courtois was burning seaweed to collect potassium nitrate when he noticed that a beautiful violet vapor was produced. When the vapor cooled, it changed to dark, shiny, metallic-like crystals. Humphry Davy later suggested the name iodine for the element from the Greek word iodos, for "violet."
As with chlorine and bromine, iodine is obtained from seawater. It can also be produced from Chile saltpeter (sodium nitrate), in which it occurs as an impurity in the form of sodium iodate (NaIO3).
Properties and uses. Iodine vapor is irritating to the eyes and respiratory system. It is highly toxic if ingested. Iodine is the least active of the common halogens (not counting astatine).
The human body uses iodine to make thyroxine, an important hormone (chemical messenger) produced by the thyroid gland. (The thyroid is a gland located in the neck that plays an important role in metabolism—a term used to describe processes of energy production and use by the body.) If insufficient amounts of iodine are present in the diet, a person may develop a condition known as goiter, a sometimes noticeable enlargement of the thyroid gland. Once the relationship between iodine and goiter were first discovered, manufacturers of table salt began to add iodine (in the form of sodium iodide) to their product (iodized salt). This practice has largely eliminated the problem of goiter in modern developed nations.
Iodine is also used commercially in a variety of products including dyes, specialized soaps, lubricants, photographic film, medicines, and pharmaceuticals.
Astatine is generally regarded as one of the rarest naturally occurring elements. According to some estimates, no more than 44 milligrams of the element are to be found in Earth's crust. It is hardly surprising, then, that the element was first produced synthetically. In 1940, three physicists at the University of California at Berkeley—D. R. Corson, K. R. Mackenzie, and Emilio Segrè (1905–1989)—made astatine by bombarding the element bismuth with alpha particles in a cyclotron (a particle accelerator or atom-smasher).
About 24 isotopes (forms) of astatine exist, all of them radioactive. The most long-lived has a half-life of 8.3 hours, meaning that half of a sample of the element disappears in 8.3 hours. Because it is so rare and has such a short half-life, astatine is one of the most poorly understood of all chemical elements. It has no practical applications at this time.
[See also Element, chemical; Organic chemistry; Periodic table ]
"Halogens." UXL Encyclopedia of Science. . Encyclopedia.com. (November 24, 2017). http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/halogens-1
"Halogens." UXL Encyclopedia of Science. . Retrieved November 24, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/halogens-1
The halogens are the family of chemical elements that includes fluorine (atomic symbol F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The halogens make up Group VIIA of the Periodic Table of the elements. Elemental halogens are diatomic molecules. However, due to their high reactivity, the halogens are never found in nature in native form. The family name means "salt-forming," from the Greek for salt, halos, and for generating genes. The salinity of the oceans on Earth is due in large part to such halogen salts (halides) as sodium chloride (NaCl) and potassium iodide (KI).
Halogens display physical and chemical properties typical of nonmetals. They have relatively low melting and boiling points that increase steadily down the group. Near room temperature, the halogens span all of the physical states: Fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. All of the elements are colored, with the color becoming more intense moving down the group. Fluorine gas is pale yellow, and chlorine gas is a yellowish green. Liquid bromine and its vapors are brownish red. Solid iodine appears as shiny, dark gray crystals, and the vapors are a deep purple. The halogens are poor thermal and electrical conductors in all phases, and as solids they are brittle and crumbly. The halogens have distinctive, unpleasant odors, will burn exposed flesh, and are toxic.
The neutral atoms of the halogens possess seven outer electrons. An additional electron can be added to halogen atoms to form singly charged negative ions. These ions have a closed outer-shell configuration. Electronegativity is a measure of the ability of an atom of one element to remove an electron from an atom of another element. As a group, the halogens are among the most electronegative elements. Fluorine has the highest electronegativity of all the elements. Halogens are so reactive that all the elements except helium and neon have been found to react with at least one of the halogens. Fluorine is always assigned a formal oxidation number of –1, whereas the other halogens can exhibit a range of oxidation numbers.
The ability of halogens to form chemical compounds with all metals and most nonmetals has led to a wide variety of uses for these elements. Chlorine is used as a bleach and a disinfectant. Iodine has been used as a topical microbicide. Iodine and bromine are added to halogen lamps to lengthen
the life of the filament and prevent darkening of the bulb. Chloride and iodide are essential dietary minerals for humans. Organic compounds containing halogens are used as fire-retardants (halons), as refrigerants (Freons), and in nonstick coatings (Teflon). Silver bromide and silver iodide have been used in photographic emulsions since the early days of photography. Many halogenated compounds are toxic. A well-known example is DDT (dichlorodiphenyltrichloroethane), once a widely applied pesticide that was banned in the United States after severe environmental effects were discovered. All known isotopes of astatine are radioactive, with the longest-lived isotope having a half-life of about eight hours. Relatively little is known of the physical and chemical properties of astatine. However, it is predicted to have properties similar to iodine.
see also Bromine; Chlorine; Fluorine; Iodine.
John Michael Nicovich
Lide, David R., ed. (2003). The CRC Handbook of Chemistry and Physics, 84th edition. Boca Raton, FL: CRC Press.
Winter, Mark. "WebElements™ Periodic Table." The University of Sheffield and WebElements, Ltd., U.K.. Available from <http://www.webelements.com>.
"Halogens." Chemistry: Foundations and Applications. . Encyclopedia.com. (November 24, 2017). http://www.encyclopedia.com/science/news-wires-white-papers-and-books/halogens
"Halogens." Chemistry: Foundations and Applications. . Retrieved November 24, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/news-wires-white-papers-and-books/halogens