Metals - Real-life applications
Alkali Metals
The members of the alkali metal family are distinguished by the fact that they have a valence electron configuration of s 1 This means that they tend to bond easily with other substances. Alkali metals tend to form positive ions, or cations, with a charge of 1+.
In contact with water, alkali metals form a negatively charged hydroxide ion (OH − ). When alkali metals react with water, one hydrogen atom splits off from the water molecule to form hydrogen gas, while the other hydrogen atom remains bonded to the oxygen, forming hydroxide. Where the heavier members of the alkali metal family are concerned, reactions can often be so vigorous that the result is combustion or even explosion. Alkali metals also react with oxygen to produce either an oxide, peroxide, or superoxide, depending on the particular member of the alkali metal family involved.
Shiny and soft enough to be cut with a knife, the alkali metals are usually white, though cesium is a yellowish white. When placed in a flame, most of these substances produce characteristic
The alkali metals, listed below by atomic number and chemical symbol, are:
- 3. Lithium (Li)
- 11. Sodium (Na)
- 19. Potassium (K)
- 37. Rubidium (Rb)
- 55. Cesium (Cs)
- 87. Francium (Fr)
Alkaline Earth Metals
Occupying Group 2 of the periodic table are the alkaline earth metals, which have a valence electron configuration of s 2 . Like the alkali metals, the alkaline earth metals are known for their high reactivity—a tendency for bonds between atoms or molecules to be made or broken in such a way that materials are transformed. Thus they are seldom
The alkaline earth metals are also like the alkali metals in that they have the properties of a base as opposed to an acid. An acid is a substance that, when dissolved in water, produces positive ions of hydrogen, designated symbolically as H + ions. It is thus known as a proton donor. A base, on the other hand, produces negative hydroxide ions when dissolved in water. These are designated by the symbol OH − . Bases are therefore characterized as proton acceptors.
The alkaline earth metals are shiny, and most are white or silvery in color. Like their "cousins" in the alkali metal family, they glow with characteristic colors when heated. Calcium glows orange, strontium a very bright red, and barium an apple green. Physically they are soft, though not as soft as the alkali metals; nor are their levels of reactivity as great as those of their neighbors in Group 1.
The alkaline earth metals, listed below by atomic number and chemical symbol, are:
- 4. Beryllium (Be)
- 12. Magnesium (Mg)
- 20. Cadmium (Ca)
- 38. Strontium (Sr)
- 56. Barium (Ba)
- 88. Radium (Ra)
Transition Metals
The transition metals are the only elements that fill the dorbitals. They are also the only elements that have valence electrons on two different principal energy levels. This sets them apart from the representative or main-group elements, of which the alkali metals and alkaline earth metals are a part. The transition metals also follow irregular patterns of orbital filling, which further distinguish them from the representative elements.
Other than that, there is not much that differentiates the transition metals, a broad grouping that includes precious metals such as gold, silver, and platinum, as well as highly functional metals such as iron, manganese, and zinc. Many of the transition metals, particularly those on periods 4, 5, and 6, form useful alloys with one another, and with other elements. Because of their differences in valence electron configuration, however, they do not always combine in the same ways, even within an element. Iron, for instance, sometimes releases two electrons in chemical bonding, and at other times three.
GROUPING THE TRANSITION METALS.
There is no easy way to group the transition metals, though some of these elements are traditionally categorized together. These do not constitute "families" as such, but they do provide useful ways to break down the otherwise rather daunting lineup of the transition metals. In two cases, there is at least a relationship between group number on the periodic table and the categories loosely assigned to a collection of transition metals. Thus the "coinage metals"—copper, silver, and gold—all occupy Group 9 on the IUPAC version of the periodic table. (These have traditionally been associated with one another because their resistance to oxidation, combined with their malleability and beauty, has made them useful materials for fashioning coins.
Likewise the members of the "zinc group"—zinc, cadmium, and mercury—on Group 10 on the IUPAC periodic table have often been associated as a miniature unit due to common properties. Members of the "platinum group"—platinum, iridium, osmium, palladium, rhodium, and ruthenium—occupy a rectangle on the table, corresponding to periods 5 and 6, and groups 6 through 8. What actually makes them a "group," however, is the fact that they tend to appear together in nature. Iron, nickel, and cobalt, found alongside one another on Period 4, may be grouped together because they are all magnetic to some degree.
The transition metals, listed below by atomic number and chemical symbol, are:
- 21. Scandium (Sc)
- 22. Titanium (Ti)
- 23. Vanadium (V)
- 24. Chromium (Cr)
- 25. Manganese (Mn)
- 26. Iron (Fe)
- 27. Cobalt (Co)
- 28. Nickel (Ni)
- 29. Copper (Cu)
- 30. Zinc (Zn)
- 39. Yttrium (Y)
- 40. Zirconium (Zr)
- 41. Niobium (Nb)
- 42. Molybdenum (Mo)
- 43. Technetium (Tc)
- 44. Ruthenium (Ru)
- 45. Rhodium (Rh)
- 46. Palladium (Pd)
- 47. Silver (Ag)
- 48. Cadmium (Cd)
- 57. Lanthanum (La)
- 72. Hafnium (Hf)
- 73. Tantalum (Ta)
- 74. Tungsten (W)
- 75. Rhenium (Re)
- 76. Osmium (Os)
- 77. Iridium (Ir)
- 78. Platinum (Pt)
- 79. Gold (Au)
- 80. Mercury (Hg)
- 89. Actinium (Ac)
- 104. Rutherfordium (Rf)
- 105. Dubnium (Db)
- 106. Seaborgium (Sg)
- 107. Bohrium (Bh)
- 108. Hassium (Hs)
- 109. Meitnerium (Mt)
- 110. Ununnilium (Uun)
- 111. Unununium (Uuu)
- 112. Ununbium (Uub)
The last nine, along with 11 of the actinides, are part of the list of metals known collectively as the transuranium elements—that is, elements with atomic numbers higher than 92. These have all been produced artificially, in most cases within a laboratory. Note that the last three had not been named as of 2001: the "names" by which they are known simply mean, respectively, 110, 111, and 112.
Lanthanides
The lanthanides are the first of two inner transition metal groups. Note that in the list of transition metals above, lanthanum (57) is followed by hafnium (72). The "missing" group of 14 elements, normally shown at the bottom of the periodic table, is known as the lanthanides, a family which can be defined by the fact that all fill the 4f orbital. However, because lanthanum (which does not fill an f orbital) exhibits similar properties, it is usually included with the lanthanides.
Bright and silvery in appearance, many of the lanthanides—though they are metals—are so soft they can be cut with a knife. They also tend to be highly reactive. Though they were once known as the "rare earth metals," lanthanides only seemed rare because they were difficult to extract from compounds containing other substances—including other lanthanides. Because their properties are so similar, and because they tend to be found together in the same substances, the original isolation and identification of the lanthanides was an arduous task that took well over a century.
The lanthanides (in addition to lanthanum, listed earlier with the transition metals), are:
- 58. Cerium (Ce)
- 59. Praseodymium (Pr)
- 60. Neodymium (Nd)
- 61. Promethium (Pm)
- 62. Samarium (Sm)
- 63. Europium (Eu)
- 64. Gadolinium (Gd)
- 65. Terbium (Tb)
- 66. Dysprosium (Dy)
- 67. Holmium (Ho)
- 68. Erbium (Er)
- 69. Thulium (Tm)
- 70. Ytterbium (Yb)
- 71. Lutetium (Lu)
Actinides
Just as lanthanum is followed on the periodic table by an element with an atomic number 15 points higher, so actinium (89) is followed by rutherfordium (104). The elements in the "gap" are the other inner transition family, the actinides, which all fill the 5f orbital. These are placed below the lanthanides at the bottom of the chart, and just as lanthanum is included with the lanthanides, even though it does not fill an f orbital, so actinium is usually lumped in with the actinides due to similarities in properties.
Most of the actinides tend to be unstable or radioactive, the later ones highly so. The first three members of the group (not counting actinium) occur in nature, while the other 11 are transuranium elements created artificially. In most cases, these were produced with a cyclotron or other machine for accelerating atoms, generally at the research center of the University of California at Berkeley during a period from 1940 to 1961. Einsteinium and fermium, however, were by-products of nuclear testing at Bikini Atoll in the south Pacific in 1952. The principal use of the actinides is in nuclear applications—weaponry or power plants—though some of them also have specialized uses as well.
The actinides (in addition to actinium, listed earlier with the transition metals), are:
- 90. Thorium (Th)
- 91. Protactinium (Pa)
- 92. Uranium (U)
- 93. Neptunium (Np)
- 94. Plutonium (Pu)
- 95. Americium (Am)
- 96. Curium (Cm)
- 97. Berkelium (Bk)
- 98. Californium (Cf)
- 99. Einsteinium (Es)
- 100. Fermium (Fm)
- 101. Mendelevium (Md)
- 102. Nobelium (No)
- 103. Lawrencium (Lr)
"Orphan" Metals
Seven other metals remain to be discussed. These are the "orphan" metals, which appear in groups 3, 4, and 5 of the North American periodic table (13, 14, and 15 of the IUPAC table), and they are called "orphans" simply because none belongs to a clearly defined family. Sometimes aluminum and the three elements below it in Group 3—gallium, indium, and thallium—are lumped together as the "aluminum family." This is not a widely recognized distinction, but will be used here only for the purpose of giving some form of organization to the "orphan" elements.
Though they do not belong to families, the "orphan" metals are nonetheless highly important. Aluminum, after all, is the most abundant of all metals on Earth, and has wide applications in daily life. Tin, too, is widely known, as is lead. Somewhat less recognizable is bismuth, but it appears in a product with which most Americans are familiar: Pepto-Bismol.
The seven "orphans" are listed below, along with atomic number and chemical symbol:
- 13. Aluminum (Al)
- 31. Gallium (Ga)
- 49. Indium (In)
- 50. Tin (Sn)
- 81. Thallium (Tl)
- 82. Lead (Pb)
- 83. Bismuth (Bi)
Aluminum
Named after alum, a salt known from ancient times and used by the Egyptians, Romans, and Greeks, aluminum is so reactive that it proved difficult to isolate. Only in 1825 did Danish physicist Hans Christian Ørsted (1777-1851) isolate an impure version of aluminum metal, using a complicated four-step process. Two years later, German chemist Friedrich Wöhler (1800-1882) obtained pure aluminum from a reaction of metallic potassium with aluminum chloride, or AlCl 3 .
For years thereafter, aluminum continued to be so difficult to produce that it acquired the status of a precious metal. European kings displayed treasures of aluminum as though they were gold or silver, and by 1855, it was selling for about $100,000 a pound. Then in 1886, French metallurgist Paul-Louis-Toussaint Héroult (1863-1914) and American chemist Charles Martin Hall (1863-1914) independently developed a process that made aluminum relatively easy to extract by means of electrolysis. Thanks to what became known as the Hall-Héroult process, the price of aluminum dropped to around 30 cents a pound.
America's aluminum comes primarily from mines in Alabama, Arkansas, and Georgia, where it often appears in a clay called kaolin, used in making porcelain. Aluminum can also be found in deposits of feldspar, mica, granite, and other clays. The principal sources of aluminum outside the United States include France, Surinam, Jamaica, and parts of Africa.
USES FOR ALUMINUM.
Though aluminum is highly reactive, it does not corrode in moist air. Instead of forming rust the way iron does, it forms a thin, hard, invisible coating of aluminum oxide. This, along with its malleability, makes it highly useful for producing cans and other food containers. The World Trade Center Towers in New York City gleam with a corrosion-resistant covering of aluminum. Likewise the element is used as a coating for mirrors: not only is it cheaper than silver, but unlike silver, it does not tarnish and turn black. Also, a thin coating of metallic aluminum gives Mylar balloons their silvery sheen.
Aluminum conducts electricity about 60% as well as copper, meaning that it is an extraordinarily good conductor, useful for transmitting electrical power. Because it is much more lightweight than copper and highly ductile, it is often used in high-voltage electric lines. Its conductivity also makes aluminum a favorite material for kitchen pots and pans. Unlike copper, it is not known to be toxic; however, because aluminum dissolves in strong acids, sometimes tomatoes and other acidic foods acquire a "metallic" taste when cooked in an aluminum pot.
Though it is soft in its pure form, when combined with copper, manganese, silicon, magnesium, and/or zinc, aluminum can produce strong, highly useful alloys. These find application in automobiles, airplanes, bridges, highway signs, buildings, and storage tanks. Aluminum's high ductility also makes it the metal of choice for foil wrappings—that is, aluminum foil. Compounds containing aluminum are used in a wide range of applications, including antacids; antiperspirants (potassium aluminum sulfate tends to close up sweat-gland ducts); and water purifiers.
Other Elements in the Aluminum Group
GALLIUM.
When Russian chemist Dmitri Ivanovitch Mendeleev (1834-1907) created the periodic table in 1869, he predicted the existence of several missing elements on the basis of "holes" between elements on the table. Among these was what he called "eka-aluminum," discovered in the 1870s by French chemist Paul Emile Lecoq de Boisbaudran (1838-1912). Bois-baudran named the new element gallium, after the ancient name of his homeland, Gaul.
Gallium will melt if held in the hand, and remains liquid over a large range of temperatures—from 85.6°F (29.76°C) to 3,999.2°F (2,204°C). This makes it highly useful for thermometers with a large temperature range. Used in a number of applications within the electronics industry, gallium is present in the compounds gallium arsenide and gallium phosphide, applied in lasers, solar cells, transistors, and other solid-state devices.
INDIUM AND THALLIUM.
As Bois-baudran would later do, German mineralogists Ferdinand Reich (1799-1882) and Hieronymus Theodor Richter (1824-1898) used a method called spectroscopy to discover one of the aluminum-group metals. Spectroscopy involves analyzing the frequencies of light (that is, the colors) emitted by an item of matter.
However, Reich faced a problem when studying the zinc ores that he believed contained a new element: he was colorblind. Therefore he called on the help of Richter, his assistant. In 1863, they discovered an element they called indium because of the indigo blue color it emitted. Reich later tried to claim that he made the discovery alone, but Richter's role is indisputable.
Indium is used in making alloys, one of which is a combination with gallium that is liquid at room temperature. Its alloys are used in bearings, dental amalgams, and nuclear control roles. Various compounds are applied in solid-state electronics devices, and in electronics research.
Two years before the discovery of indium, English physicist William Crookes (1832-1919) discovered another element using spectroscopic analysis. This was thallium, named after the Greek word thallos ("green twig"), a reference to a lime-green spectral line that told Crookes he had found a new element. Thallium is applied in electronic equipment, and the compound thallium sulfate is used for killing ants and rodents.
Three Other "Orphans"
TIN.
Known from ancient times, tin was combined with copper to make bronze—an alloy so significant it defines an entire stage of human technological development. Tin's chemical symbol, Sn, refers to the name the Romans gave it: stannum. Just as tin added to copper gave bronze its strength, alloys of tin with copper and antimony created pewter, a highly useful, malleable material for pots and dinnerware popular from the thirteenth century through the early nineteenth century.
Tin is found primarily in Bolivia, Malaysia, Indonesia, Thailand, Zaire, and Nigeria—a wide geographical range. Like aluminum today, it has been widely used as a coating for other metals to retard corrosion. Hence the term "tin can," referring to a steel can coated with tin. Also, tin has been used to coat iron or steel roofs, but because zinc is cheaper, it is now most often the coating of choice; nonetheless, these are still typically called "tin roofs."
LEAD.
Due to knowledge of its toxic properties, lead is not frequently used today, but this was not always the case. The Romans, for instance, used plumbum as they called it (hence the chemical symbol Pb) as a material for making water pipes. This is the root of our own word plumber. (The same root explains the verb plumb or the noun plumb bob. A plumb bob is a lead weight, hung from a string, used to ensure that the walls of a structure are "plumb"—that is, perpendicular to the foundation.) Many historians believe that plumbum in the Romans' water supply was one of the reasons behind the decline and fall of the Roman Empire.
The human body can only excrete very small quantities of lead a day, and this is particularly true of children. Even in small concentrations, lead can cause elevation of blood pressure; reduction in the synthesis of hemoglobin, which carries oxygen from the lungs to the blood and organs; and decreased ability to utilize vitamin D and calcium for strengthening bones. Higher concentrations of lead can effect the central nervous system, resulting in decreased mental functioning and hearing damage. Prolonged exposure can result in a coma or even death.
Before these facts became widely known in the late twentieth century, lead was applied as an ingredient in paint. In addition, it was used in water pipes, and as an anti-knock agent in gasolines. Increased awareness of the health hazards involved have led to a discontinuation of these practices. (Note that pencil "lead" is actually graphite, a form of carbon.)
Lead, however, is used for absorbing radiation—for instance, as a shield against X rays. Another important use for lead today is in storage batteries. Yet another "application" of lead relates to its three stable isotopes (lead−206, −207, and −208), the end result of radioactive decay on the part of various isotopes of uranium and other radioactive elements.
BISMUTH.
During the Middle Ages, when a number of unscientific ideas prevailed, scholars believed that lead eventually became tin, tin eventually turned into bismuth, and bismuth developed into silver. Hence they referred to bismuth as "roof of silver." Only in the sixteenth century did German metallurgist Georgius Agricola (George Bauer; 1494-1555) put forward the idea that bismuth was a separate element.
Bismuth appears in nature not only in its elemental form, but also in a number of compounds and ores. It is also obtained as a by-product from the smelting of gold, silver, copper, and lead. Because bismuth expands dramatically when it cools, early printers used a bismuth alloy to make metal type. When poured into forms cut with the shapes of letters and other characters, the alloy produced symbols with clear, sharp edges.
Wood's metal is a bismuth alloy involving cadmium, tin, and lead, which melts at a temperature of 158°F (70°C). This makes it useful for automatic sprinkler systems inside a building: when a fire breaks out, it melts the Wood's metal seal, and turns on the sprinklers. Bismuth compounds are applied in cosmetics, paints, and dyes, and prior to the twentieth century, compounds of bismuth were widely used for treating sexually transmitted diseases. Today, antibiotics have largely taken their place, but bismuth still has one well-known medicinal application: in bismuth subsalicylate, better known by the brand name Pepto-Bismol.
WHERE TO LEARN MORE
Aluminum Association (Web site). <http://www.aluminum.org> (May 28, 2001).
Kerrod, Robin. Matter and Materials. Illustrated by Terry Hadler. Tarrytown, NY: Benchmark Books, 1996.
"Lead Programs." Environmental Protection Agency (Web site). <http://www.epa.gov/opptintr/lead/> (May 28, 2001).
Mebane, Robert C. and Thomas R. Rybolt. Metals. Illustrated by Anni Matsick. New York: Twenty-First Century Books, 1995.
"Metals Industry Guide." About.com (Web site). <http://metals.about.com/industry/metals/mbody.htm> (May 28, 2001).
"Minerals and Metals." Natural Resources of Canada (Web site). <http://www.nrcan.gc.ca/mms/school/e_mine.htm> (May 28, 2001).
Oxlade, Chris. Metal. Chicago, IL: Heinemann Library, 2001.
Snedden, Robert. Materials. Des Plaines, IL: Heinemann Library, 1999.
WebElements (Web site). <http://www.webelements.com> (May 22, 2001).
Whyman, Kathryn. Metals and Alloys. Illustrated by Louise Nevett and Simon Bishop. New York: Gloucester Press, 1988.
Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.