Alkali metals

Type of physical science: Chemistry

Field of study: Chemistry of the elements

The alkali metals comprise a group of six very similar elements. Three of these elements, sodium, potassium, and lithium, are particularly important. The elements and their compounds are used extensively in many areas, including agriculture and medicine, and have numerous industrial applications.

Overview

The term "alkali metals" refers to the group I metallic elements lithium, sodium, potassium, rubidium, cesium, and francium. The elements are abbreviated by the chemical symbols Li, Na, K, Rb, Cs, and Fr, respectively. The symbols for sodium and potassium are derived from the Latin words for these elements--natrium and kalium, respectively.

Although the pure alkali metals were not isolated until the nineteenth century, many compounds of these elements have been known since ancient times. Reference to a substance called neter (sodium carbonate), obtained from plant ashes, can be found in the Old Testament.

Ordinary salt (sodium chloride) has been widely used since the ancient Roman era, and saltpeter (potassium nitrate) has been used by Europeans to make gunpowder since the twelfth century.

All the alkali metals are highly reactive and therefore do not occur in nature as free elements. Rather, they are combined with other elements and occur as compounds or in minerals.

With the exception of francium, they are all found in the waters of mineral springs, salt lakes, or oceans. Minerals containing alkali metals include spodumene (a mineral composed of lithium, aluminum, silicon, and oxygen), halite (sodium chloride), Chile saltpeter (sodium nitrate), sylvite (potassium chloride), and pollucite (composed of cesium, aluminum, silicon, oxygen, and water).

Alkali metals are commercially obtained from these sources.

Sodium and potassium are relatively common elements in the earth's crust (being the sixth and seventh most abundant elements, respectively), whereas lithium, rubidium, and cesium are rarer. In terms of percentage abundance, sodium, potassium, and rubidium have abundances of 2.6, 2.4, and 0.03 percent, respectively. Lithium and cesium both have abundances of less than 0.01 percent. Francium, which is radioactive, is extremely rare and occurs naturally only through the radioactive decay of other elements. It is also formed in nuclear reactors. It has been estimated that there is less than 25 grams of francium in the entire earth's crust.

Sodium and potassium were first isolated as the pure elements by the English chemist Sir Humphry Davy in 1807. Davy passed electricity through molten samples of sodium hydroxide (caustic soda) and potassium hydroxide (caustic potash), a process known as electrolysis, and succeeded in separating the metals from their compounds. Sodium derives its name from the Italian word soda, a term used during the Middle Ages referring to the alkaline (basic) substances obtained from plant ashes. Potassium originates from potasse, a French word for substances extracted from wood ashes.

In 1817, the Swedish chemist and mineralogist Johann Arfedson discovered lithium in several minerals. The pure metal was isolated some years later by Davy's method of electrolysis.

The element's name is derived from the Greek word lithos, meaning stone, signifying lithium's occurrence in minerals.

Between 1860 and 1861, Robert Bunsen and Gustav Kirchhoff, using an instrument called a spectroscope, discovered rubidium in the mineral lepidolite and cesium in mineral waters. A spectroscope analyzes the characteristic light emitted when an element is heated. When burned, compounds containing alkali metals generally impart a bright, characteristic color to flames. For example, lithium gives a vivid crimson color, sodium's light is bright yellow, and potassium produces a violet light. These colors also provide the basis for a simple qualitative test for the presence of these common alkali metals. Rubidium and cesium were named for the bright colors that Bunsen and Kirchhoff observed with the spectroscope: rubidus, which is the Latin word for red, and caesius, which is an ancient word for the blue color of the sky. The remaining member of group I elements, francium, was discovered in 1939 by Marguerite Perey of the Curie Institute in Paris and is named for Perey's native country.

Commercially, alkali metals can be obtained by electrolysis of molten compounds.

Lithium and sodium are obtained by the electrolysis of molten mixtures of either lithium chloride or sodium chloride with calcium chloride. Chlorine gas is a useful by-product of this process.

Potassium and the other, more reactive alkali metals can also be obtained by this procedure, but it is generally more difficult to do so because they have lower melting points and hence vaporize more readily. An alternative method involves treating molten chloride compounds, such as potassium chloride, with sodium vapor. Since they all have low boiling points, alkali metals may be purified by distillation.

Because alkali metals are so reactive, they readily combine with oxygen and water vapor in the air; therefore, they must be stored in an inert environment. Storage in a mineral oil, such as kerosene, or in an inert atmosphere of nitrogen or argon, minimizes their reaction with oxygen or water. Nevertheless, the metals readily lose their shiny metallic appearance because of surface reactions with even trace amounts of oxygen and, as a result, rapidly become dull. Yet, the bright metallic shine of the metals can be seen when they are freshly cut. This may be done easily with a sharp knife since the metals are extremely soft. Lithium, sodium, potassium, and rubidium are all silver-colored metals, while cesium has a gold appearance.

The softness and high reactivity of alkali metals distinguish them from most other metals, which generally are much harder and less reactive. Furthermore, alkali metals have low melting points (ranging from 28.4 degrees Celsius for cesium to 180.5 degrees Celsius for lithium) and low boiling points (similarly ranging from 669.3 degrees Celsius to 1,342 degrees Celsius). The densities of the alkali metals are also low when compared with most other metals.

In fact, the densities of lithium, sodium, and potassium are all less than that of water, on which these metals actually float. Like all metals, however, the group I elements are good conductors of heat and electricity.

The term "alkali metals" for the group I elements is derived from the fact that a number of compounds containing these metals are alkaline (basic). For example, in water the metals react spontaneously to produce a basic solution. In each case, the metals dissolve to form soluble basic compounds called hydroxides (historically referred to as alkalies). Hydrogen gas is also emitted when the alkali metals react with water.

The alkali metals react directly with most nonmetals except the noble gases (group O).

These reactions may occur at room temperature (for example, lithium reacts with nitrogen and all the alkali metals react directly with oxygen) or at elevated temperatures (for example, all react with hydrogen, sulfur, and chlorine) and result in the formation of binary compounds (compounds composed of two elements). Binary compounds containing an alkali metal and a nonmetal generally combine through the formation of positively charged ions (cations) and negatively charged ions (anions), respectively, giving rise to ionic compounds. The high reactivity of alkali metals results from the tendency of their atoms to lose one electron readily to form singly charged cations. Consequently, the metals are powerful reducing agents, meaning they readily donate electrons to other substances.

There is also variation in the chemical reactivity within the group I elements that can be related to the relative sizes of the alkali metal atoms. Both reactivity and atomic size increase down the group from lithium to francium. Since larger metal atoms lose their outer electrons (valence electrons) more readily than smaller atoms, larger atoms within a group generally have the greatest chemical reactivity. Thus, a small atom, such as lithium, tends to hold on to its electrons more strongly than a larger atom, such as cesium, and therefore is less reactive. This trend in reactivities is best demonstrated by the reaction of the alkali metals with water: Lithium reacts relatively slowly with water, sodium reacts vigorously, potassium inflames, and rubidium and cesium react explosively.

Compounds of the alkali metals are generally ionic (composed of cations and anions).

They are usually white solids, such as ordinary table salt (sodium chloride), and have high melting points. Most compounds dissolve in water to give solutions, which are electrolytic (conduct electricity), as a result of dissociation of the compound into its constituent ions.

Applications

Alkali metals, and especially their compounds, have numerous important uses and applications. The major use of the metals stems from their high reactivity and low melting points.

Metallic lithium and sodium have the widest applications, whereas the other metals are somewhat limited in their use because of their extreme reactivity. Great care must be taken when using any of the alkali metals since they react violently with many substances--for example, water and acids.

Although lithium is a soft metal, it may be mixed with harder metals to produce lightweight alloys. For example, lithium-magnesium and lithium-aluminum alloys have applications in the aviation and aerospace industries. Because it is a reactive metal, lithium is also used in the manufacture of other metals (iron, steel, and copper) to remove impurities such as oxygen, nitrogen, or sulfur. Similarly, a small quantity of lithium is placed inside vacuum tubes immediately before sealing, since it reacts with trace amounts of nitrogen or oxygen that may still be present. Lithium is also used in batteries and in ceramics and special glasses. For example, the 508-centimeter telescope at Mount Palomar contains a small amount of lithium.

Compounds of lithium are used to treat schizophrenia (lithium carbonate), in the production of pharmaceuticals, perfumes, and other organic substances (lithium hydride and lithium aluminum hydride), in the manufacture of greases and lubricants (lithium stearate), and in the removal of moisture from air conditioning and industrial drying systems (lithium chloride and lithium bromide).

Sodium is an important alkali metal with the most commercial applications. Like lithium, sodium is also used in the manufacture of many drugs, dyes, and other organic compounds, and the synthetic rubber industry consumes large amounts of the metal. Liquid sodium is used as a coolant in nuclear reactors. Sodium is used in metallurgy for the production of other metals such as titanium and the other alkali metals, potassium, rubidium, and cesium. A small amount of sodium is used in sodium vapor lamps, which emit a bright yellow light, and in mercury-sodium vapor lamps, which produce a white light and are used at road intersections and sports stadiums. Sodium compounds are used to make glass (sodium carbonate), paper, soap, detergents, and textiles (sodium carbonate and sodium hydroxide). They are also used as disinfectants and bleaching agents (sodium bisulfite and sodium hypochlorite), in effervescent salts, beverages, and baking powder (sodium bicarbonate), photography (sodium thiosulphate), and as fertilizers (sodium nitrate).

Potassium and potassium-sodium alloys are used as heat exchangers in nuclear reactors, and the alloy is also employed as a catalyst in organic reactions. Potassium compounds are used in explosives, gunpowder, and matches (potassium nitrate), bleaching (potassium permanganate), as sedatives (potassium bromide), as fertilizers (potassium chloride and potassium sulfate), and in special optical glasses (potassium carbonate).

Rubidium is used in photoelectric cells and to remove gases in sealed electron tubes.

Cesium ionizes readily when heated or struck by light and so is used in photomultiplier tubes to measure radioactivity. It is also used in the thousands of commercial atomic clocks that have been constructed since 1955. A further use for cesium is in thermionic devices, which convert heat directly into electricity by the ejection of electrons from a heated surface. These devices have potential as sources of electrical power in space vehicles and in other situations that require the use of a compact, portable source of electricity. Because both rubidium and cesium are easily ionized, their use as fuel in ion-propulsion engines for space vehicles has been studied. It is estimated that cesium could propel a spacecraft 140 times farther than the same amount of any known liquid or solid fuel.

Sodium and potassium, in the form of ions, are also essential nutrients for most forms of life. In humans, they play important roles in maintaining the normal flow of water between cells and body fluid. They are also involved in the transmission of nerve impulses and in muscle contraction. Potassium is particularly important for plant growth and is therefore an essential ingredient in most fertilizers.

Context

Alkali metals occur only in nature combined with other elements in compounds or minerals. Because they are so reactive, isolation of the metals from these compounds was not possible until the early nineteenth century when Davy's pioneering work on electrolysis led to the isolation of pure sodium, potassium, and later, lithium. As a result of Davy's work with the alkali metals, together with the later contributions of his student Michael Faraday, the science of electrochemistry--chemical changes produced by electricity--soon became an important process in science and industry.

In addition to isolating the pure alkali metals by the process of electrolysis, it was soon discovered that many other substances could also be made by this procedure, including hydrogen, oxygen, ozone, hydrogen peroxide, and chlorine. Electrochemistry was also employed in the extraction and purification of other metals--for example, copper, aluminum, and magnesium--and to electroplate metals and alloys to increase their resistance to corrosion.

Research with alkali metals also played a role in the evolution of spectroscopy, an area of science that enables the structure of molecules to be examined by studying their interactions with various types of electromagnetic radiation, such as light. The characteristic light emitted by heating alkali metals further demonstrated the great utility of spectroscopy and its application in the spectrochemical analysis of elements.

Principal terms

ALKALI: a basic substance that neutralizes acids

COMPOUND: a pure substance that is composed of at least two kinds of atoms

ELECTROLYSIS: a chemical change or reaction induced by electricity

ELEMENT: a pure substance that contains only one kind of atom

GROUP: a vertical column of elements in the periodic table, the members of which have similar chemical and physical properties

Bibliography

CHEMICALS FROM SODIUM CHLORIDE, PART 2. Princeton, N.J.: Films for the Humanities and Sciences, 1986. This is a program in the "Chemistry in Action" video series, which is designed to relate basic concepts of chemistry with commercially important technologies. The twenty-minute color videotape describes the electrolytic production of sodium from sodium chloride, as well as the metal's uses.

Emsley, John. THE ELEMENTS. New York: Oxford University Press, 1989. An excellent collection of data and facts on all the elements, each of which are summarized in two pages. Includes information on the chemical, physical, and nuclear properties of the elements, as well as their abundances, distributions, and biological roles.

Idhe, Aaron J. THE DEVELOPMENT OF MODERN CHEMISTRY. New York: Harper & Row, 1964. Chapter 5 gives a good account of the early development of electrochemistry. Of particular interest is the description of Davy's work on the isolation of the alkali metals.

McQuarrie, Donald A., and Peter A. Rock. DESCRIPTIVE CHEMISTRY. New York: W. H. Freeman, 1985. Chapter 3 contains a description of the physical properties, manufacture, and uses of alkali metals. The section contains a number of color photographs of the metals and their flame tests.

Weast, Robert C. HANDBOOK OF CHEMISTRY AND PHYSICS. 66th ed. Boca Raton, Fla.: CRC Press, 1986. Section B5 of this edition, published by the Chemical Rubber Company, contains about a half-page description of every element. Information on the discovery, occurrence, physical properties, manufacture, uses, and important compounds of the alkali metals can be found in this section. Revised and updated annually.

Weeks, Mary E. "Discovery of the Elements." JOURNAL OF CHEMICAL EDUCATION, 44 (1968). An easy-to-read historical account of the discovery and isolation of all the chemical elements.

Acids and Bases