Halogens

Type of physical science: Chemistry

Field of study: Chemistry of the elements

The halogens are the elements fluorine, chlorine, bromine, iodine, and astatine. They all have seven electrons in their outer shell, one less than a noble gas. Therefore, they are all very reactive and make useful compounds such as Teflon and Freon. In nature, they are never found as a free element, but often are found as a salt.

Overview

The halogens are known as the "salt generators," as these elements are typically found as their salt. The halogens, group VIIA on the periodic table, are fluorine, chlorine, bromine, iodine, and astatine. Chlorine makes 0.03 percent of the crust of the earth, whereas bromine is 0.0002 percent and iodine is merely 0.0001 percent of the crust. Astatine has been detected in extremely small amounts in the natural decay series of thorium and uranium. Astatine was discovered as a synthetic element.

Because the halogens are the group VIIA elements, they have many close similarities and trends in properties. Each halogen has seven electrons in its outer shell, only one less than inert gases. Thus, they are very reactive. Halogens are never found free in nature, because of their high reactivities. They have been studied mostly in the form of their inorganic salts. The halogens, however, have use in many organic compounds, such as DDT (dichloro-diphenyl-trichloroethane), chloroform, polychlo-rinated biphenyls (PCBs), and so forth. Many of the halogenated compounds have been found to have an adverse effect on the environment.

As halogen atoms each have seven valence electrons and need only one electron to fill the outer electron shell to have a stable noble gas configuration; they are all very reactive. In attempting to complete the outer p sublevel of the electron shells, all the halogens behave as nonmetals and as oxidizing agents, thus they are able to accept any free electron. Each halogen, from fluorine downward, decreases in its reactivity, though all are reactive. In physical states, they go from gaseous (fluorine and chlorine) to liquid (bromine) to solid crystals (iodine) at room temperature. In the gaseous state, all halogens exist as diatomic molecules, thus having shared, filled outer electron shells. Halogens are salts when they form binary compounds with metals. The halides, the binary compounds of halogens with metals, contain the halogen in the -1 oxidation state.

Fluorine, with atomic number 9, is the most reactive of the halogens. At room temperature, fluorine exists as a pale yellow gas. As the temperature is depressed, it becomes a liquid at its boiling point of -190 degrees Celsius. At -220 degrees Celsius, fluorine comes to its melting point, becoming solid. At atomic mass 18.9984, fluorine exists in nature only as the isotope with atomic mass 19. The five other isotopes of fluorine have been prepared artificially.

These isotopes are radioactive, unstable, and have short half-lives.

Fluorine is found in the mineral salt fluorite. This salt of calcium fluoride is common and widespread. It is usually found in hydrothermal veins and is common in dolomites and limestone. Fluorspar, commercial fluorite, is mined worldwide. Other minerals containing fluorine include cryolite, a sodium aluminum fluoride found in important deposits in Greenland; and fluoroapatite, a calcium phosphate fluoride.

Of the halogens, chlorine is the most abundant in nature. Like fluorine, it never exists in a free state, but exists as a diatomic molecule or a mineral. The most common form of chlorine is in the mineral halite, which is the salt sodium chloride. With atomic number 17, chlorine is not as reactive as fluorine but is still very reactive. Chlorine exists as a green gas at room temperature, providing the root of its name, green. At -35 degrees Celsius, the green gas condenses at the boiling point of chlorine liquid. Further cooling to -101 degrees Celsius is the melting point of solid chlorine. Of the eleven isotopes of chlorine, two occur naturally. Chlorine 35 and chlorine 37 exist in a 3:1 ratio. The atomic mass of chlorine is 35.453.

Chlorine is found mostly in the mineral halite. Halite is dissolved in the waters of salt springs, salt lakes, and the ocean. It is the major salt in deposits of enclosed basins. The deposits of salt have been formed by the gradual evaporation and ultimate drying up of enclosed bodies of salt water. Salt beds range from a few meters to more than 60 meters in thickness. Other chloride containing minerals include sylvite, cerargyite, and atacamite. None of these minerals is as common as halite.

Bromine is a halogen found in very low concentrations in seawater. Though bromine is less reactive than chlorine and much less reactive than fluorine, most elements combine directly with bromine. With atomic number 35, bromine is unique among nonmetals as it is a liquid at room temperature. With a temperature rise to 58.78 degrees Celsius, bromine boils to a gaseous diatomic state. At -7.2 degrees Celsius, bromine freezes to a solid state. As a liquid, it is relatively dense, with a specific gravity of 3.12. The liquid evaporates rapidly, giving a red-brown vapor with an irritating odor. The root of bromine means "stinking." Bromine has an atomic mass of 79.909.

Iodine is commercially found by the burning of seaweed ashes, as seaweed concentrates the iodine from the seawater. Iodine is also found in some nitrate deposits. Iodine exists as a black, lustrous solid crystal at room temperature. It can volatize easily to give a violet vapor. This provides the root of the name, violet. A temperature of 113.5 degrees Celsius is required to melt iodine and 184.4 degrees Celsius will boil it. Iodine is less reactive than the halogens above it in group VIIA. With atomic number 53, the atomic mass of iodine is 126.904.

Astatine was the second synthetic element discovered after technetium. With atomic number 85 and atomic mass 210, astatine was synthesized in the laboratory in 1940. As it has existed only as a radioactive gas, it cannot be made into anything that can be seen or detected in any way except by its radioactivity. Astatine means "unstable." Astatine appears to be chemically similar to iodine, and is also volatile. In nature, only astatine 219 with a half-life of 0.9 minute is reported to exist within uranium ores.

Applications

The saline halides are those that consist of infinite lattices of separate metal and halogen ions in the solid state. The halides of the alkali and alkali earth metals, most of the lanthanides, and the transition metals in their lower oxidation states are saline in character. In most saline halides, the number of halide ions coordinated to each metal ion is much greater than the number of halide ions that could be bound to the metal ion in any molecular configuration.

Halogen atoms easily gain an electron to form the halide ions, and their compounds with metals are quite common. Most metal halides have ionic bonds, in which the electron from the metal is transferred to the valence shell of the halide, provided the metal is in a low oxidation state. When the metal is in a high oxidation state, polarization of the anion often produces a covalently bonded molecule, in which a pair of electrons is shared between the two atoms. The halide will crystallize in a molecular lattice only when, in the molecule, the number of halogen atoms bonded to the metal with low electronegativity equals the maximum coordination number of the electronegative atom for that halogen. Molecular halides, with covalent bonding, are characterized by low melting points, and saline halides, with ionic bonding, are characterized by high melting points. As the charge-to-radius ratio of the low electronegative atom increases, the probability of forming molecular halides increases.

Hydrogen halides are the binary hydrogen compounds of halogens that can be prepared by direct combination of the elements. The vigor of this combination varies substantially from fluorine to iodine. Fluorine reacts violently to hydrogen as soon as the two gases are mixed.

Chlorine and hydrogen react at a much slower rate, provided their mixtures are not heated or exposed to ultraviolet light. The light of heat splits the chloride molecules into chlorine atoms and initiates a chain reaction that is explosively fast. Chain mechanisms are also involved in the reaction of hydrogen with bromide and iodide, but reactions are less vigorous than with chloride.

Hydrogen halides can also be made from binary salts by reactions with a nonvolatile acid.

Many of the halogen compounds can be made with nonmetals. The structures of these substances can be predicted on the basis of the valence shell electron-pair repulsion theory. Most nonmetals form more than one compound with a given halogen. The number of halogen atoms that can become bound to any particular nonmetal can be related by two factors. One is the electronic structures of the elements that are combined and the other has to do with the sizes of the atoms.

Each halogen atom contains seven electrons in its valence shell and requires only one more to achieve the stable noble gas configuration. Thus, there is little tendency to form multiple bonds with other nonmetals. Halogens normally do not accept electrons in the formation of coordinate covalent bonds, because this would mean the addition of two electrons to a valence shell that already contains seven electrons, thereby exceeding the stable electron configuration by one electron.

At room temperature, fluorine is a pale yellow gas. It is the most reactive of all elements because of very low bond energy in the diatomic fluoride molecule. These thus split to be very reactive fluorine atoms. The low bond energy is caused by electron-electron repulsions between the small, compact electron-rich valence shells of the bonded fluorine atoms. Elemental fluorine is such an active oxidizing agent that it can be made only by electrolysis. Hydrogen fluoride is dissolved in molten potassium fluoride. Electrolysis produces fluoride gas at the anode (positive charge) and hydrogen gas at the cathode (negative charge). Fluoride is extremely reactant. In reactions with hydrogen, phosphorus, antimony, and some metals, the reaction is so vigorous that spontaneous ignition may occur. When fluoride reacts with some metals such as copper, the reaction stops after a while because surface deposits of metal fluoride protect the base metal from further attack. The high reactivity of the element contributes much to the stability of its compounds. One of the most inert plastics that is manufactured is the Teflon type of polymer, which is classified as a fluorocarbon. Teflon was originally created during work on the atomic bomb. Small amounts of fluorides in drinking water and in toothpastes have been found to be beneficial in preventing the formation of cavities, especially in the teeth of children. Two parts per million fluoride in drinking water may cause mottling of tooth enamel and larger quantities are toxic.

Chlorine is a pale green gas at room temperature. Chlorine molecules have slightly larger bond energy than fluoride, and chlorine atoms are less electronegative than fluorine. As a result, chlorine is less reactive than fluorine. Chlorine is produced by electrolysis of both molten and aqueous sodium chloride. More than 12 million tons of chlorine are produced each year in the United States, making it the eighth in total annual tons made by industry. It is largely a chemical intermediate and is used to make other chemicals. Chlorine is also used to treat municipal drinking water, to make solvents and plastics such as polyvinyl chloride (PVC), and to manufacture pesticides.

When salts are added to water, the salts dissolve and dissociate into ions. Seawater contains various ions; 55.07 percent of all ions in sea water are chlorine. The quantity of chlorine ions present in a water sample has been measured to establish salinity. Silver nitrate added to a water sample combines to make silver chloride. The silver also combines with bromine and iodine. The chlorine concentration measured in this way is termed "chlorinity," measured in grams chlorine to kilograms seawater. Chlorinity is thus defined as the quantity of silver required to remove all the halogens from 0.3285 kilogram of seawater.

At room temperature, bromine is a volatile, nonviscous, red liquid. Bromine is found in very low concentrations in seawater but can be easily removed by oxidation of bromine by chloride and because of the volatility of bromide. Most bromine is extracted from brines obtained from wells in Arkansas and Michigan, with brine ion concentration fifty to sixty times that of seawater. Half of the bromine produced each year is used to make ethylene dibromide, which is an additive to leaded gasoline. It prevents deposits of lead compounds from forming inside the engine. Bromine is also used to make silver bromide, for light-sensitive emulsions of photographic film and paper.

The recovery of iodine is expensive. Commercial quantities of iodine are recovered from the ashes of burned seaweed, in which iodine concentrations approach 1 percent. The iodine ion is oxidized to iodide molecule using chloride or another oxidizing agent. Iodine is used to make various medical products, and silver iodide is used for photographic films. Iodine is not nearly as poisonous as chlorine or bromine. Lack of iodine in the diet leads to the disease goiter.

Astatine was the second synthetic element. It was made by bombarding bismuth with α particles of sufficient energy to penetrate the nucleus. Alpha particles contain two protons and two neutrons. The α particles accelerated in a cyclotron, which starts the particles going around faster until they ultimately strike bismuth. When α particles enter the nucleus of a bismuth atom, two neutrons are ejected, leaving one atom of astatine. Astatine can be separated from bismuth by volatilizing it. The astatine vapor is invisible, but since it is radioactive, it can be measured with a Geiger-Muller counter.

Context

The history of the halogens is exciting in chemistry, and newer findings show halogens to be sometimes useful and dangerous. In 1886, French chemist Henri Moissan isolated fluorine.

Fluorine will attack anything, including glass and porcelain. At least two chemists died while investigating fluorine. Moissan used a platinum vessel, which resists fluorine for some time, and whatever else could be made of fluorspar (already a fluorine compound, it therefore cannot be attacked by fluorine any further). In spite of its violence, and partly because of it, fluorine is an important element. Some of its compounds are disinfectants. A useful but dangerous fluoride compound is the gas Freon, which is used in refrigeration. When Freon escapes, it combines with ozone, breaking down the earth's protective ozone layer.

French chemist Count Claude Louis Berthollet first thought chlorine was a compound of hydrochloric acid and oxygen. In order to recognize chlorine as an element, it first had to be proved that there were acids that do not contain oxygen. Once done, hydrogen was known to be an element, and chlorine, which had successfully resisted all attempts to break it down, had to be an element. Though chlorine has many beneficial uses, it is a deadly poison. Even small amounts of the gas destroy lung tissue. It was used as a poison gas in 1915.

Working with a mother liquor, which is seawater with all dissolved ions in it from several places, French chemist Antoine Jerome Balard noticed there was a yellowish layer above the layer of starch that had been covered by iodine. The yellowish layer not only had a different color but also had a distinct odor. Balard distilled the yellow portion and saw it turn into a thick red smoke that condensed into a liquid upon cooling. In his report, Balard stated that he had discovered a new element and that it could be obtained in two ways. One way was adding chlorine ion to the mother liquor, then distilling it and capturing the red smoke. The other was to add chlorine to the mother liquor and then ether. Bromine causes severe burns when in contact with the skin, but an application of glycerine is helpful if applied immediately. Vapors are poisonous, damaging the linings of the nose, throat, and lungs.

The process of obtaining saltpeter from seaweed ashes was called saltpeter boiling. The whole process was possible only because seaweed has the peculiarity of accumulating chemicals that are dissolved in seawater. Therefore, the mother liquor, as the result of the first burning of ashes was called, was a large collection of chemicals. Bernard Courtois added too many ashes to the boiler in 1811. The result was a cloud of violet vapor with an irritating odor. The vapor condensed into dark, shiny crystals. Courtois carried out a number of experiments and found the new substance, iodine, could not be decomposed by heat; it was reluctant to combine with oxygen and carbon. Iodine is used medically as an antiseptic.

Principal terms

ALPHA PARTICLES: helium nuclei of two protons and two neutrons with no electrons, carrying a +2 charge, emitted as radioactive decay, and may be used to bombard elements to create heavier elements

ATOMIC MASS: the sum of the atomic number and the number of neutrons in the nucleus of an atom

ATOMIC NUMBER: the number of protons found in the nucleus of an atom, with each element having its own atomic number

DIATOMIC MOLECULE: a molecule made of two of the same kinds of atoms, having a shared, filled outer electron shell

ELECTROLYSIS: the separation of ions in solution by passing an electric current through the solution

ELECTRONEGATIVITY: an atom's attraction for electrons in a bond

GROUP VIIA: the halogen elements of the periodic table, all with seven electrons in their outer electron shell

HALOGENS: those elements whose atoms have seven electrons in their outer electron shell, including fluorine, chlorine, bromine, iodine, and astatine, which are all very reactive

OXIDATION: a reaction in which electrons are lost

VOLATILE: a chemical reaction that is vigorous and produces a gaseous state

Bibliography

Applequist, Douglas, Charles Depuy, and Kenneth L. Rinehart. INTRODUCTION TO ORGANIC CHEMISTRY. 3d ed. New York: John Wiley & Sons, 1982. This book discusses reactions with halogen atoms in organic compounds such as Teflon, DDT, and chloroform.

Brady, James E., and Gerard E. Humiston. GENERAL CHEMISTRY PRINCIPLES AND STRUCTURES. 3d ed. New York: John Wiley & Sons, 1982. This freshman-level chemistry text covers all aspects of general chemistry. It was useful in specific properties of the halogens, illustrations, uses of halogens, chemical reactions with halogens, and terminology.

Duxbury, Alyn C., and Alison Duxbury. AN INTRODUCTION TO THE WORLD'S OCEANS. Reading, Mass.: Addison-Wesley, 1984. This is a survey of the science of oceanography, describing research methods, geology of the seafloor, sedimentation, ocean currents, and some marine life. This source was particularly useful in relation to salt water and to halogen ions within the salt water.

Hurlbut, Cornelius, Jr., and Cornelis Klein. MANUAL OF MINEROLOGY. 19th ed. New York: John Wiley & Sons, 1977. This book covers minerals in great depth. Describes crystal form, chemical and physical properties, and occurrence and uses of minerals. Useful in learning where halides were found.

Jolly, William L. THE CHEMISTRY OF THE NON-METALS. Englewood Cliffs, N.J.: Prentice-Hall, 1986. This book covers specific properties of each of the nonmetals, including the halogens. Covers the reactions of each of these elements and some important uses for each of the nonmetal elements.

Ley, Willy. THE DISCOVERY OF THE ELEMENTS. New York: Delacorte Press, 1968. Using the periodic table, Ley shows how the elements, including the halogens, were discovered and isolated. Explains how each element fits into the periodic table, the experiments done to discover these elements, and some basic properties of each of the elements.

Nechamkin, Howard. THE CHEMISTRY OF THE ELEMENTS. New York: McGraw-Hill, 1968. Chapter by chapter, Nechamkin discusses the various elements. The units on the halogens were most useful, giving physical and chemical properties of each, sources, discovery, important chemical reactions, and uses of each.

Seaborg, Glenn T., and Evens G. Valens. ELEMENTS OF THE UNIVERSE. New York: E. P. Dutton, 1958. Seaborg and Valens survey all the elements. This source was particularly useful with astatine, as Seaborg had experience with synthetic elements. Other elements were well treated, though somewhat dated.