Alkaline earths

Type of physical science: Chemistry

Field of study: Chemistry of the elements

The alkaline earths comprise a group of six metals: beryllium, magnesium, calcium, strontium, barium, and radium. The elements and their compounds have many important applications in alloys, metallurgy, agriculture, the building industry, and medicine. Magnesium and calcium are abundant in nature and play significant roles in biological processes.

Overview

The alkaline earths are a group of six metals--beryllium, magnesium, calcium, strontium, barium, and radium--that constitute group II of the periodic table of elements. The metals are abbreviated by the chemical symbols Be, Mg, Ca, Sr, Ba, and Ra, respectively. The elements were not isolated from their compounds until the nineteenth century, but several compounds of the elements have been known and used since ancient times. Lime, a compound of calcium, was used by the ancient Greeks to make mortar by mixing it with sand and water. The Chinese used mortar to lay stones in the Great Wall. Aquamarines and emeralds have been valued as precious gems throughout the ages and are minerals of beryllium. Milk of magnesia (a solution of magnesium hydroxide in water) has long been used as an antacid and is named for the ancient Turkish city of Magnesia, where magnesium minerals were found. The mineral alabaster (calcium sulfate) has been used since ancient times to make bowls, statues, and similar personal items.

The alkaline earths are reactive metals and are therefore not found in nature as the free elements; however, they are less reactive than the group I elements. Common minerals from which the alkaline earths are extracted include beryl (beryllium aluminum silicate), magnesite (magnesium carbonate), dolomite (a mixture of magnesium and calcium carbonates), chalk, limestone and marble (forms of calcium carbonate), fluorospar (calcium fluoride), gypsum (calcium sulfate), celestite (strontium sulfate), strontianite (strontium carbonate), barite (barium sulfate), witherite (barium carbonate), and pitchblende (a uranium ore that contains trace amounts of radium). In most cases, the ores containing the alkaline earths are converted to oxides (compounds containing the metal and oxygen) or halides (compounds containing the metal and a group VII element), which are then reduced electrochemically to the metal. Calcium and magnesium also occur in seawater at concentrations of about 0.4 and 1.3 grams per liter, respectively. Industrial plants have been constructed that extract magnesium from this source.

Calcium and magnesium are relatively abundant elements in the earth's crust (percentage abundances of 3.6 and 2.5, respectively), while barium, strontium, and beryllium are less common (0.05, 0.025, and 0.001 percent abundances, respectively). Radium is a highly radioactive and rare element. It is widely distributed in nature but at low concentrations (for example, a ton of pitchblende contains only a fraction of a gram of radium) and therefore is difficult to extract.

Throughout the eighteenth and nineteenth centuries, chemists recognized the existence of new elements in a variety of minerals and compounds. As practical skills developed, these new elements were isolated and purified. Since ancient times, it had been observed that certain minerals did not melt and remained unchanged when heated. These minerals were referred to as "earths." Some earths were also known to produce alkaline (basic) solutions in water, and so came to be known as the "alkaline earths." Initially it was believed that the alkaline earths were elements, but eventually it was recognized that they were actually compounds of new metallic elements and oxygen (for example, baryta was barium oxide, strontia was strontium oxide, and beryllia was beryllium oxide). The term "alkaline earths" was, however, retained for these metals. In 1808, the English chemist Sir Humphry Davy was the first to isolate the alkaline earth metals from their compounds by using electrolysis. This process involved passing electricity through molten compounds containing the alkaline earth metals. By this process, Davy obtained magnesium, calcium, strontium, and barium. Later, in 1828, the German chemist Friedrich Wohler isolated beryllium by the reaction of potassium with beryllium chloride. In 1898, radium was discovered and isolated in an impure form by Marie and Pierre Curie. From several tons of pitchblende residues, they succeeded in extracting 0.1 gram of radium chloride. Pure elemental radium was subsequently isolated in 1910 by Marie Curie by electrolysis of molten radium chloride. Commercially pure magnesium, calcium, strontium, and barium are now obtained by electrolysis of molten chloride compounds or by reduction of the oxides. Beryllium is obtained by reduction of beryllium fluoride with magnesium metal. Radium metal is also obtained by reduction of its compounds, but the element has no large-scale commercial use.

Two of the alkaline earth elements are named for towns in which minerals containing the elements were discovered: Strontian, a town in Scotland, and the ancient city of Magnesia in Turkey. Two other elements derive their names from minerals: calx (the Latin word for "lime") and beryl (the Greek word for "gem"). Formerly, beryllium was called glucinum (from the Greek glykys, meaning "sweet") since beryllium compounds were found to have a sweet taste (early chemists frequently engaged in the dangerous practice of tasting new substances as a means of characterizing them). "Barium" comes from the Latin word barys, meaning "heavy," and reflects the dense nature of barium and its compounds.

Radium was named from the Latin word radius, meaning "ray," since radium compounds glow in the dark.

The melting points and boiling points of the alkaline earths are generally low compared to most metals but are higher than those of the corresponding group I elements in the same period: magnesium has the lowest (melting point 650 degrees Celsius, boiling point 1,090 degrees Celsius) and beryllium the highest (melting point 1,278 degrees Celsius, boiling point 2,970 degrees Celsius). The densities are lower than those of most metals (except for the group I metals) and range from 1.55 grams per cubic centimeter for calcium to 5 grams per cubic centimeter for radium. Like all metals, the alkaline earths are good conductors of heat and electricity.

While less reactive than the group I metals, most of the alkaline earth elements react readily with oxygen and water. Calcium, strontium, and barium are therefore stored in an inert environment, such as in mineral oil, to prevent reaction with water or oxygen. These metals, together with radium, readily react with cold water, but the rates of reaction are slower than those of the group I metals. The reaction with water produces alkaline solutions, resulting from the formation of hydroxide compounds, and hydrogen gas. Magnesium is less reactive because of a thin protective coating of an inert magnesium compound (magnesium oxycarbonate), but when freshly cut it reacts slowly with steam or hot water to produce magnesium hydroxide and hydrogen. This reaction was used during World War II to produce hydrogen for weather balloons. Beryllium is unreactive at ordinary temperatures because of a coating of inert beryllium oxide. The metals readily lose their shiny metallic appearance because of surface reaction with even trace amounts of oxygen, and form a dull coating of the oxide. Although harder than the group I metals, the alkaline earths may be cut fairly easily to reveal the bright metallic luster characteristic of metals, but this soon dulls because of the surface reaction with oxygen.

Beryllium is a gray metal, and the other alkaline earths appear silvery-white when freshly cut.

Magnesium, calcium, strontium, and barium all react either at room temperature or at elevated temperatures with most nonmetals (for example, oxygen, hydrogen, nitrogen, carbon, sulfur, and chlorine) to form binary ionic compounds. These are compounds that contain two elements that combine through the formation of positive metal ions (cations) and negative nonmetal ions (anions). Cations of the alkaline earths normally have a +2 charge as a result of the loss of two valence electrons (these are the electrons furthest away from an atom's nucleus and are the first to be lost during a reaction). Because of its small size, beryllium's chemistry differs from the other group II elements. For example, it generally reacts with nonmetals (for example, oxygen, nitrogen, and chlorine) only at elevated temperatures to form binary covalent compounds. Ions are therefore absent in many beryllium compounds, and electrons are shared between metal and nonmetal atoms. Because most of the alkaline earths are so reactive, they readily replace other metals in compounds and are therefore good reducing agents, particularly magnesium and calcium. The chemical reactivity of the metals generally increases down the group, which is a result of the increase in relative atomic size going down the group. Larger atoms lose their valence electrons more easily and therefore are more chemically reactive. Thus, barium reacts vigorously with water, but beryllium is inert. With the exception of beryllium, which reacts only slowly, all the alkaline earths react vigorously with acids.

Compounds of the alkaline earths are usually white solids with high melting points.

Those that contain neutral anions and are water soluble produce neutral solutions; however, in water, soluble beryllium compounds are usually acidic.

Applications

There are many important uses and applications for the alkaline earths and their compounds. The major use for the metals results from their low densities (beryllium and magnesium) and high chemical reactivities (calcium, strontium, barium, and radium). Caution is advised when handling calcium, strontium, or barium because of their spontaneous reaction with water and acids. Beryllium is toxic, so exposure to the metal (and its compounds) should be minimized. Radium is also toxic, and its radioactive nature necessitates appropriate precautions.

Because beryllium is a very light metal with high mechanical strength and a high melting point, it is used to construct lightweight alloys for specific applications, particularly in the aerospace industry. Since it is an inert metal, it also enhances the corrosion resistance of its alloys. An alloy of beryllium and copper is used for electrically conducting springs, electrical contacts, and nonsparking tools. The latter are particularly useful for areas where explosive gases might be present. The metal is also used in gyroscopes and computer components, and in instruments in which lightness and stiffness are required. Beryllium absorbs fast neutrons and so finds applications in nuclear reactors. It is also transparent to X rays and is therefore an ideal material to construct the windows of X-ray tubes.

Magnesium became an important metal during World War II, when it was used as a lightweight alloy with aluminum, manganese, zinc, and other metals. It was particularly useful for aircraft parts and in incendiary bombs, in which the magnesium bomb case was ignited by a combustible mixture inside. Lightweight magnesium alloys now are also used for automobile parts, tools, and equipment. Because magnesium is a fairly chemically reactive metal, it is used for a variety of nonstructural purposes, including corrosion protection of other metals. Pieces of magnesium are placed near buried steel pipelines and water tanks, and are attached to the steel hulls of ships. Corrosive chemicals in the environment react selectively with the magnesium, rather than with the steel. It is therefore easier and more economical simply to replace the corroded magnesium rather than repair or replace the steel. Magnesium ignites easily, especially in powdered form, and burns with a brilliant white light. It is therefore used in fireworks and flares, and magnesium ribbon is used in camera flashbulbs. Since hot magnesium reacts with water, the latter should not be used on burning magnesium or magnesium fires. Magnesium is a strong reducing agent and so is used in metallurgy to displace other metals (for example, uranium, titanium, zirconium, and hafnium) from their compounds.

Small amounts of calcium, strontium, and barium are used in alloys for special purposes, and also to remove residual gases in vacuum tubes. Calcium is used as a reducing agent to prepare other metals such as thorium, uranium, and chromium. Only small amounts of metallic strontium and barium are produced annually.

Applications for radium result from the radioactive nature of the metal. It has been used in medicine for cancer treatment and in luminous paints for watch and instrument dials. Safer and cheaper substances are now preferred for these purposes, however.

Compounds of the alkaline earth metals--especially magnesium, calcium, strontium, and barium--have an enormous variety of applications. Because of their high toxicity, beryllium compounds have limited applications. Beryllium oxide is a good heat conductor but a poor electrical conductor, and therefore is used as a heat-conducting electrical insulator in electronic devices and lasers. Beryllium fluoride is used in the manufacture of certain glasses.

Magnesium and calcium compounds have many applications. They are used for fireproofing wood and as disinfectants (magnesium chloride), in talcum powder (magnesium oxide), as de-icers on roads (calcium chloride), as bleaching agents and algicides in swimming pools (calcium hypochlorite), and as food preservatives (calcium tartrate). In medicine, milk of magnesia (magnesium hydroxide) is a common antacid, and Epsom salts (magnesium sulfate) is used as a laxative. The building industry uses large amounts of plaster of Paris (calcium sulfate), and quicklime (calcium oxide) and slaked lime (calcium hydroxide) are used to make mortar, cement, and concrete. These materials absorb carbon dioxide from the air and are hardened through conversion to calcium carbonate. Magnesium and calcium compounds are also used in sugar refining, in fire extinguishers, in wine and beer manufacture, in tanning leather, in insulating materials, in fertilizers, and in the manufacture of pharmaceuticals.

Strontium compounds (for example, strontium nitrate) burn with a brilliant, crimson-red flame and so are used in flares and fireworks. They have also been used in luminous paints (strontium sulfide), in sugar refining (strontium hydroxide), and in medicine as sedatives and anticonvulsants (strontium bromide).

Barium compounds have several uses. Barium sulfate strongly absorbs X rays and is used in X-ray examinations to show an outline of a patient's digestive system. Although barium compounds are toxic, barium sulfate is highly insoluble and can therefore pass through the body unabsorbed. By contrast, barium carbonate is used as a rat poison. Compounds of barium such as barium nitrate burn with a green flame, and so are used in the manufacture of fireworks and signal flares. Other applications include use in motor oil detergents (barium oxide and barium hydroxide), in the manufacture of ceramics and special glass (barium carbonate), and in photocells and semiconductors (barium selenide).

Magnesium and calcium, in ionic form, are essential minerals for all plants and animals. Magnesium is a component of chlorophyll, the plant pigment responsible for photosynthesis (the process through which plants convert sunlight into energy). Magnesium is also involved in enzyme reactions that speed up certain biochemical reactions that control energy transfer processes in living cells. It is also involved in the replication of deoxyribonucleic acid (DNA) and ribonucleic acid (RNA), two substances that control heredity and genetic functions in living cells. Calcium is very important to animals that have shells or skeletons, and in higher animals is the most abundant mineral. It also plays a role in muscle contraction, vision, nerve excitation, and blood clotting. While calcium and magnesium are essential elements, the other alkaline earths are toxic. Since they are chemically similar to calcium and magnesium, they may be readily incorporated into living systems through environmental contamination. For example, radioactive strontium is a common product from nuclear fallout and is readily deposited in bones, where it can replace calcium. Radiation from this strontium can damage bone marrow, impair the formation of new cells, and ultimately may induce cancers. Radium can also concentrate in bone and release harmful radiation internally, with similar consequences.

Context

Magnesium is the most common alkaline earth metal used in its elemental state, and several hundred thousand tons are produced annually. Much of this is used to make alloys, and most aluminum contains about 5 percent magnesium to improve its mechanical properties and enhance its corrosion resistance. The widespread use of aluminum, a relatively new metal compared to metals such as iron and copper, which have been used since ancient times, has revolutionized contemporary industries.

The economic significance of alkaline earth compounds has been considerable. This is particularly true for the building industries, in which enormous amounts of calcium compounds are used to make mortar, cement, concrete, and plaster. Agriculture has also benefited from the use of calcium compounds such as lime, to improve soil conditions, and calcium phosphate fertilizers, which provide essential plant nutrients such as phosphorous. Calcium and magnesium are essential minerals for humans. Together with potassium and sodium, they constitute about 99 percent of the total mineral content of the human body.

Principal terms

ALKALINE: describes a substance that is basic

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

PERIOD: a horizontal row of elements in the periodic table

REDUCTION: the process that occurs when an atom, ion, or molecule gains electrons; substances that donate electrons are reducing agents

Bibliography

Idhe, Aaron J. THE DEVELOPMENT OF MODERN CHEMISTRY. New York: Harper & Row, 1964. This text gives a good account of the early development of electrochemistry and its use in the isolation of the alkaline earth metals (chapter 5), the early manufacture of magnesium (chapter 25), and the discovery of radium (chapter 18).

McQuarrie, Donald A., and Peter A. Rock. DESCRIPTIVE CHEMISTRY. New York: W. H. Freeman, 1985. Chapter 4 contains a description of the physical properties, manufacture, and uses of the alkaline earth metals. The section contains a number of color photographs of the metals.

Schubert, J. "Beryllium and Berylliosis." SCIENTIFIC AMERICAN 199 (August, 1958): 27-33. Although somewhat dated, this paper is easy to read and provides much general information about beryllium, its properties and occurrence in nature, and its toxicity.

Weast, Robert C. HANDBOOK OF CHEMISTRY AND PHYSICS. 66th ed. Boca Raton, Fla.: CRC Press, 1986. Section B5 of this edition contains a half-page or so description of every element. Information on the discovery, occurrence, physical properties, manufacture, uses, and important compounds of the alkaline earth metals can be found here. 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.

Williams, D. R. "Life's Essential Elements." EDUCATION IN CHEMISTRY 10 (1973): 56. This article discusses the elements that are essential for the growth and development of living organisms, and includes calcium and magnesium.

Acids and Bases