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Alkaline Earth Metals

ALKALINE EARTH METALS

CONCEPT

The six alkaline earth metalsberyllium, magnesium, calcium, strontium, barium, and radiumcomprise Group 2 on the periodic table of elements. This puts them beside the alkali metals in Group 1, and as their names suggest, the two families share a number of characteristics, most notably their high reactivity. Also, like the alkali metals, or indeed any other family on the periodic table, not all members of the alkali metal family are created equally in terms of their abundance on Earth or their usefulness to human life. Magnesium and calcium have a number of uses, ranging from building and other structural applications to dietary supplements. In fact, both are significant components in the metabolism of living thingsincluding the human body. Barium and beryllium have numerous specialized applications in areas from jewelry to medicine, while strontium is primarily used in fireworks. Radium, on the other hand, is rarely used outside of laboratories, in large part because its radioactive qualities pose a hazard to human life.

HOW IT WORKS

Defining a Family

The expression "families of elements" refers to groups of elements on the periodic table that share certain characteristics. These include (in addition to the alkaline earth metals and the alkali metals) the transition metals, halogens, noble gases, lanthanides, and actinides. (All of these are covered in separate essays within this book.) In addition, there are several larger categories with regard to shared traits that often cross family lines; thus all elements are classified either as metals, metalloids, and nonmetals. (These are also discussed in separate essays, which include reference to "orphans," or elements that do not belong to one of the families mentioned above.)

These groupings, both in terms of family and the broader divisions, relate both to external, observable characteristics, as well as to behaviors on the part of electrons in the elements' atomic structures. For instance, metals, which comprise the vast majority of elements on the periodic table, tend to be shiny, hard, and malleable (that is, they can bend without breaking.) Many of them melt at fairly high temperatures, and virtually all of them vaporize (become gases) at high temperatures. Metals also form ionic bonds, the tightest form of chemical bonding.

ELECTRON CONFIGURATIONS OF THE ALKALINE EARTH METALS.

Where families are concerned, there are certain observable properties that led chemists in the past to group the alkaline earth metals together. These properties will be discussed with regard to the alkaline earth metals, but another point should be stressed in relation to the division of elements into families. With the advances in understanding that followed the discovery of the electron in 1897, along with the development of quantum theory in the early twentieth century, chemists developed a more fundamental definition of family in terms of electron configuration.

As noted, the alkaline earth metal family occupies the second group, or column, in the periodic table. All elements in a particular group, regardless of their apparent differences, have a common pattern in the configuration of their valence electronsthe electrons at the "outside" of the atom, involved in chemical bonding. (By contrast, the core electrons, which occupy lower regions of energy within the atom, play no role in the bonding of elements.)

All members of the alkaline earth metal family have a valence electron configuration of s 2. This means that two electrons are involved in chemical bonding, and that these electrons move through an orbital, or range of probabilities, roughly corresponding to a sphere. The s orbital pattern corresponds to the first of several sub-levels within a principal energy level.

Whatever the number of the principal energy level which corresponds to the period, or row, on the periodic table the atom has the same number of sublevels. Thus beryllium, on Period 2, has two principal energy levels, and its valence electrons are in sublevel 2s 2. At the other end of the group is radium, on Period 7. Though radium is far more complex than beryllium, with seven energy levels instead of two, nonetheless it has the same valence electron configuration, only on a higher energy level: 7s 2.

HELIUM AND THE ALKALINE EARTH METALS.

If one studies the valence electron configurations of elements on the periodic table, one notices an amazing symmetry and order. All members of a group, though their principal energy levels differ, share characteristics in their valence shell patterns. Furthermore, for the eight groups numbered in the North American version of the periodic table, the group number corresponds to the number of valence electrons.

There is only one exception: helium, with a valence electron configuration of 1s 2, is normally placed in Group 8 with the noble gases. Based on that s 2 configuration, it might seem logical to place helium atop beryllium in the alkaline earth metals family; but there are several reasons why this is not done. First of all, helium is obviously not a metal. More importantly, helium behaves in a manner quite different from that of the alkaline earth metals.

Whereas helium, like the rest of the noble gases, is highly resistant to chemical reactions and bonding, alkaline earth metals are known for their high reactivitythat is, a tendency for bonds between atoms or molecules to be made or broken so that materials are transformed. (A similar relationship exists in Group 1, which includes hydrogen and the alkali metals. All have the same valence configuration, but hydrogen is never included as a member of the alkali metals family.)

Characteristics of the Alkaline Earth Metals

Like the alkali metals, the alkaline earth metals have the properties of a base, as opposed to an acid. 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, many of which can be cut with a knife.

Yet another similarity the alkaline earth metals have with the alkali metals is the fact that four of the themmagnesium, calcium, strontium, and bariumwere either identified or isolated in the first decade of the nineteenth century by English chemist Sir Humphry Davy (1778-1829). Around the same time, Davy also isolated sodium and potassium from the alkali metal family.

REACTIVITY.

The alkaline earth metals are less reactive than the alkali metals, but like the alkali metals they are much more reactive than most elements. Again like their "cousins," they react with water to produce hydrogen gas and the metal hydroxide, though their reactions are less pronounced than those of the alkali metals. Magnesium metal in its pure form is combustible, and when exposed to air, it burns with an intense white light, combining with the oxygen to produce magnesium oxide. Likewise calcium, strontium, and barium react with oxygen to form oxides.

Due to their high levels of reactivity, the alkaline earth metals rarely appear by themselves in nature; rather, they are typically found with other elements in compound form, often as carbonates or sulfates. This, again, is another similarity with the alkali metals. But whereas the alkali metals tend to form 1+ cations (positively charged atoms), the alkaline earth metals form 2+ cationsthat is, cations with a positive charge of 2.

BOILING AND MELTING POINTS.

One way that the alkaline earth metals are distinguished from the alkali metals is with regard to melting and boiling pointsthose temperatures, respectively, at which a solid metal turns into a liquid, and a liquid metal into a vapor. For the alkali metals, the temperatures of the boiling and melting points decrease with an increase in atomic number. The pattern is not so clear, however, for the alkaline earth metals.

The highest melting and boiling points are for beryllium, which indeed has the lowest atomic number. It melts at 2,348.6°F (1,287°C), and boils at 4,789.8°F (2,471°C). These figures are much higher than for lithium, the alkali metal on the same period as beryllium, which melts at 356.9°F (180.5°C) and boils at 2,457°F (1,347°C).

Magnesium, the second alkali earth metal, melts at 1,202°F (650°C), and boils at 1,994°F (1,090°C)significantly lower figures than for beryllium. However, the melting and boiling points are higher for calcium, third of the alkaline earth metals, with figures of 1,547.6°F (842°C) and 2,703.2°F (1,484°C) respectively. Melting and boiling temperatures steadily decrease as energy levels rise through strontium, barium, and radium, yet these temperatures are never lower than for magnesium.

ABUNDANCE.

Of the alkaline earth metals, calcium is the most abundant. It ranks fifth among elements in Earth's crust, accounting for 3.39% of the elemental mass. It is also fifth most abundant in the human body, with a share of 1.4%. Magnesium, which makes up 1.93% of Earth's crust, is the eighth most abundant element on Earth. It ranks seventh in the humanbody, accounting for 0.50% of the body's mass.

Barium ranks seventeenth among elementsin Earth's crust, though it accounts for only0.04% of the elemental mass. Neither it nor the other three alkali metals appear in the body insignificant quantities: indeed, barium and beryllium are poisonous, and radium is so radioactivethat exposure to it can be extremely harmful.

Within Earth's crust, strontium is present in quantities of 360 parts per million (ppm), which in fact is rather abundant compared to a number of elements. In the ocean, its presence is about 8 ppm. By contrast, the abundance of beryllium in Earth's crust is measured in parts per billion (ppb), and is estimated at 1,900 ppb. Vastly more rare is radium, which accounts for just 0.6 parts per trillion of Earth's crusta fact that made its isolation by French-Polish physicist and chemist Marie Curie (1867-1934) all the more impressive.

REAL-LIFE APPLICATIONS

Beryllium

In the eighteenth century, French mineralogist René Just-Haüy (1743-1822) had observed that both emeralds and the mineral beryl had similar properties. French chemist Louis-Nicolas Vauquelin (1763-1829) in 1798 identified the element they had in common: beryllium (Be), which has an atomic number of 4 and an atomic mass of 9.01 amu. Some three decades passed before German chemist Friedrich Wöhler (1800-1882) and French chemist Antoine Bussy (1794-1882), working independently, succeeded in isolating the element.

Beryllium is found primarily in emeralds and aquamarines, both precious stones that are forms of the beryllium alluminosilicate compound beryl. Though it is toxic to humans, beryllium nonetheless has an application in the health-care industry: because it lets through more x rays than does glass, beryllium is often used in x-ray tubes.

Metal alloys that contain about 2% beryllium tend to be particularly strong, resistant to wear, and stable at high temperatures. Copper-beryllium alloys, for instance, are applied in hand tools for industries that use flammable solvents, since tools made of these alloys do not cause sparks when struck against one another. Alloys of beryllium and nickel are applied for specialized electrical connections, as well as for high-temperature uses.

Magnesium

English botanist and physician Nehemiah Grew (1641-1712) in 1695 discovered magnesium sulfate in the springs near the English town of Epsom, Surrey. This compound, called "Epsom salts" ever since, has long been noted for its medicinal value. Epsom salts are used for treating eclampsia, a condition that causes seizures in pregnant women. The compound is also a powerful laxative, and is sometimes used to rid the body of poisonssuch as magnesium's sister element, barium.

For some time, scientists confused the oxide compound magnesia with lime or calcium carbonate, which actually involves another alkaline earth metal. In 1754, Scottish chemist and physicist Joseph Black (1728-1799) wrote "Experiments Upon Magnesia, Alba, Quick-Lime, and Some Other Alkaline Substances," an important work in which he distinguished between magnesia and lime. Davy in 1808 declared magnesia the oxide of a new element, which he dubbed magnesium, but some 20 years passed before Bussy succeeded in isolating the element.

Magnesium (Mg) has an atomic number of 12, and an atomic mass of 24.31 amu. It is found primarily in minerals such as dolomite and magnesite, both of which are carbonates; and in carnallite, a chloride. Magnesium silicates include asbestos, soapstone or talc, and mica. Not all forms of asbestos contain magnesium, but the fact that many do only serves to show the ways that chemical reactions can change the properties an element possesses in isolation.

AN IMPORTANT COMPONENT OF HEALTH.

Whereas magnesium is flammable, asbestos was once used in large quantities as a flame retardant. And whereas asbestos has been largely removed from public buildings throughout the United States due to reports linking asbestos exposure with cancer, magnesium is an important component in the health of living organisms. It plays a critical role in chlorophyll, the green pigment in plants that captures energy from sunlight, and for this reason, it is also used in fertilizers.

In the human body, magnesium ions (charged atoms) aid in the digestive process, and many people take mineral supplements containing magnesium, sometimes in combination with calcium. There is also its use as a laxative, already mentioned. Epsom salts, as befits their base or alkaline quality, are exceedingly bitterthe kind of substance a person only ingests under conditions of the most dire necessity. On the other hand, milk of magnesia is a laxative with a far less unpleasant taste.

MAGNESIUM GOES TO WAR.

It is a hallmark of magnesium's chemical versatility that the same element, so important in preserving life, has also been widely used in warfare. Just before World War I, Germany was a leading manufacturer of magnesium, thanks in large part to a method of electrolysis developed by German chemist R. W. Bunsen (1811-1899). When the United States went to war against Germany, American companies began producing magnesium in large quantities.

Bunsen had discovered that powdered magnesium burns with a brilliant white flame, and in the war, magnesium was used in flares, tracer bullets, and incendiary bombs, which ignite and burn upon impact. The bright light produced by burning magnesium has also led to a number of peacetime applicationsfor instance, in fireworks, and for flashes used in photography.

Magnesium saw service in another world war. By the time Nazi tanks rolled into Poland in 1939, the German military-industrial complex had begun using the metal for building aircraft and other forms of military equipment. America once again put its own war-production machine into operation, dramatically increasing magnesium output to a peak of nearly 184,000 tons (166,924,800 kg) in 1943.

STRUCTURAL APPLICATIONS.

Magnesium's principal use in World War I was for its incendiary qualities, but in World War II it was primarily used as a structural metal. It is lightweight, but stronger per unit of mass than any other common structural metal. As a metal for building machines and other equipment, magnesium ranks in popularity only behind iron and aluminum (which is about 50% more dense than magnesium).

The automobile industry is one area of manufacturing particularly interested in magnesium's structural qualities. On both sides of the Atlantic, automakers are using or testing vehicle parts made of alloys of magnesium and other metals, primarily aluminum. Magnesium is easily cast into complex structures, which could mean a reduction in the number of parts needed for building a carand hence a streamlining of the assembly process.

Among the types of sports equipment employing magnesium alloys are baseball catchers' masks, skis, racecars, and even horseshoes. Various brands of ladders, portable tools, electronic equipment, binoculars, cameras, furniture, and luggage also use parts made of this lightweight, durable metal.

Calcium

Davy isolated calcium (Ca) by means of electrolysis in 1808. The element, whose name is derived from the Latin calx, or "lime," has an atomic number of 20, and an atomic mass of 40.08. The principal sources of calcium are limestone and dolomite, both of which are carbonates, as well as the sulfate gypsum.

In the form of limestone and gypsum, calcium has been used as a building material since ancient times, and continues to find application in that area. Lime is combined with clay to make cement, and cement is combined with sand and water to make mortar. In addition, when mixed with sand, gravel, and water, cement makes concrete. Marbleonce used to build palaces and today applied primarily for decorative touchesalso contains calcium.

The steel, glass, paper, and metallurgical industries use slaked lime (calcium hydroxide) and quicklime, or calcium oxide. It helps remove impurities from steel, and pollutants from smokestacks, while calcium carbonate in paper provides smoothness and opacity to the finished product. When calcium carbide (CaC2) is added to water, it produces the highly flammable gas acetylene (C2H2), used in welding torches. In various compounds, calcium is used as a bleach; a material in the production of fertilizers; and as a substitute for salt as a melting agent on icy roads.

The food, cosmetic, and pharmaceutical industries use calcium in antacids, toothpaste, chewing gum, and vitamins. To an even greater extent than magnesium, calcium is important to living things, and is present in leaves, bone, teeth, shells, and coral. In the human body, it helps in the clotting of blood, the contraction of muscles, and the regulation of the heartbeat. Found in green vegetables and dairy products, calcium (along with calcium supplements) is recommended for the prevention of osteoporosis. The latter, a condition involving a loss of bone density, affects elderly women in particular, and causes bones to become brittle and break easily.

Strontium

Irish chemist and physician Adair Crawford (1748-1795) and Scottish chemist and surgeon William Cumberland Cruikshank (1790-1800) in 1790 discovered what Crawford called "a new species of earth" near Strontian in Scotland. A year later, English chemist Thomas Charles Hope (1766-1844) began studying the ore found by Crawford and Cruikshank, which they had dubbed strontia.

In reports produced during 1792 and 1793, Hope explained that strontia could be distinguished from lime or calcium hydroxide on the one hand, and baryta or barium hydroxide on the other, by virtue of its response to flame tests. Whereas calcium produced a red flame and barium a green one, strontia glowed a brilliant red easily distinguished from the darker red of calcium.

Once again, it was Davy who isolated the new element, using electrolysis, in 1808. Subsequently dubbed strontium (Sr), its atomic number is 38, and its atomic mass 87.62. Silvery white, it oxidizes rapidly in air, forming a pale yellow oxide crust on any freshly cut surface.

Though it has properties similar to those of calcium, the comparative rarity of strontium and the expense involved in extracting it offer no economic incentives for using it in place of its much more abundant sister element. Nonetheless, strontium does have a few uses, primarily because of its brilliant crimson flame. Therefore it is applied in the making of fireworks, signal flares, and tracer bulletsthat is, rounds that emit a light as they fly through the air.

One of the more controversial "applications" of strontium involved the radioactive isotope strontium-90, a by-product of nuclear weapons testing in the atmosphere from the late 1940s onward. The isotope fell to earth in a fine powder, coated the grass, was ingested by cows, and eventually wound up in the milk they produced. Because of its similarities to calcium, the isotope became incorporated into the teeth and gums of children who drank the milk, posing health concerns that helped bring an end to atmospheric testing in the early 1960s.

Barium

Aspects of barium's history are similar to those of other alkaline earth metals. During the eighteenth century, chemists were convinced that barium oxide and calcium oxide constituted the same substance, but in 1774, Swedish chemist Carl Wilhelm Scheele (1742-1786) demonstrated that barium oxide was a distinct compound. Davy isolated the element, as he did two other alkaline earth metals, by means of electrolysis, in 1808.

Barium (Ba) has an atomic mass of 137.27 and an atomic number of 56. It appears primarily in ores of barite, a sulfate, and witherite, a carbonate. Barium sulfate is used as a white pigment in paints, while barium carbonate is applied in the production of optical glass, ceramics, glazed pottery, and specialty glassware. One of its most important uses is as a drill-bit lubricantknown as a "mud" or slurryfor oil drilling. Like a number of its sister elements, barium (in the form of barium nitrate) is used in fireworks and flares. Motor oil detergents for keeping engines clean use barium oxide and barium hydroxide.

Beryllium is not the only alkaline earth metal used in making x rays, nor is magnesium the only member of the family applied as a laxative. Barium is used in enemas, and barium sulfate is used to coat the inner lining of the intestines to allow a doctor to examine a patient's digestive system. (Though barium is poisonous, in the form of barium sulfate it is safe for ingestion because the compound does not dissolve in water or other bodily fluids.) Prior to receiving x rays, a patient may be instructed to drink a chalky barium sulfate liquid, which absorbs a great deal of the radiation emitted by the x-ray machine. This adds contrast to the black-and-white x-ray photo, enabling the doctor to make a more informed diagnosis.

Radium

Today radium (Ra; atomic number 88; atomic mass 226 amu) has few uses outside of research; nonetheless, the story of its discovery by Marie Curie and her husband Pierre (1859-1906), a French physicist, is a compelling chapter not only in the history of chemistry, but of human endeavor in general. Inspired by the discovery of uranium's radioactive properties by French physicist Henri Becquerel (1852-1908), Marie Curie became intrigued with the subject of radioactivity, on which she wrote her doctoral dissertation. Setting out to find other radioactive elements, she and Pierre refined a large quantity of pitchblende, an ore commonly found in uranium mines. Within a year, they had discovered the element polonium, but were convinced that another radioactive ingredient was presentthough in much smaller amountsin pitch-blende.

The Curies spent most of their savings to purchase a ton of ore, and began working to extract enough of the hypothesized Element 88 for a usable sample0.35 oz (1 g). Laboring virtually without ceasing for four years, the Curiesby then weary and in financial difficultiesfinally produced the necessary quantity of radium. Their fortunes were about to improve: in 1903 they shared the Nobel Prize in physics with Becquerel, and in 1911, Marie received a second Nobel, this one in chemistry, for her discoveries of polonium and radium. She is the only individual in history to win Nobels in two different scientific categories.

Because the Curies failed to patent their process, however, they received no profits from the many "radium centers" that soon sprung up, touting the newly discovered element as a cure for cancer. In fact, as it turned out, the hazards associated with this highly radioactive substance outweighed any benefits. Thus radium, which at one point was used in luminous paint and on watch dials, was phased out of use. Marie Curie's death from leukemia in 1934 resulted from her prolonged exposure to radiation from radium and other elements.

WHERE TO LEARN MORE

"Alkaline Earth Metals." ChemicalElements.com (Web site). <http://www.chemicalelements.com/groups/alkaline.html>> (May 25, 2001).

"The Alkaline Earth Metals" (Web site). <http://www.nidlink.com/~jfromm/elements/alkaline.htm> (May 25, 2001).

Ebbing, Darrell D.; R. A. D. Wentworth; and James P. Birk. Introductory Chemistry. Boston: Houghton Mifflin, 1995.

Kerrod, Robin. Matter and Materials. Illustrated by Terry Hadler. Tarrytown, N.Y.: Benchmark Books, 1996.

Mebane, Robert C. and Thomas R. Rybolt. Metals. Illustrated by Anni Matsick. New York: Twenty-First Century Books, 1995.

Oxlade, Chris. Metal. Chicago, IL: Heinemann Library, 2001.

Snedden, Robert. Materials. Des Plaines, IL: Heinemann Library, 1999.

"Visual Elements: Group 1The Alkaline Earth Metals" (Web site). <http://www.chemsoc.org/viselements/pages/data/intro_groupii_data.html> (May 25, 2001).

KEY TERMS

ALKALINE EARTH METALS:

Group 2 on the periodic table of elements, with valence electron configurations of ns 2. The six alkaline earth metals, all of which are highly reactive chemically, are beryllium, magnesium, calcium, strontium, barium, and radium.

ALKALI METALS:

The elements in Group 1 of the periodic table of elements, with the exception of hydrogen. The alkali metals all have one valence electron in thes1 orbital, and are highly reactive.

CATION:

The positive ion that results when an atom loses one or more electrons. All of the alkaline earth metals tend to form 2+ cations (pronounced KAT-ieunz).

ELECTROLYSIS:

The use of an electric current to cause a chemical reaction.

ION:

An atom or group of atoms that has lost or gained one or more electrons, and thus has a net electric charge.

ISOTOPES:

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 unstablethat is, radioactive. Such is the case with the isotopes of radium, a radioactive member of the alkaline earth metals family.

ORBITAL:

A pattern of probabilities regarding the position of an electron for anatom in a particular energy state. The six alkaline earth metals all have valence electrons in an s 2 orbital, which describes a more or less spherical shape.

PERIODS:

Rows of the periodic table of elements. These represent successive principal energy levels in the atoms of the elements involved.

PRINCIPAL ENERGY LEVEL:

A value indicating the distance that an electron may move away from the nucleus of anatom. This is designated by a whole-number integer, beginning with 1 and moving upward. The higher the principal energy level, the greater the energy in the atom, and the more complex the pattern of orbitals.

RADIOACTIVITY:

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 changing into other isotopes until they become stable.

REACTIVITY:

The tendency for bonds between atoms or molecules to be made or broken in such a way that materials aretransformed.

SALT:

Generally speaking, a compound that brings together a metal and a nonmetal. More specifically, salts (along with water) are the product of a reaction between an acid and a base.

SHELL:

A group of electrons within the same principal energy level.

SUBLEVEL:

A region within the principal energy level occupied by electrons in anatom. Whatever the number n of the principal energy level, there are n sublevels. At each principal energy level, the first sublevel to be filled is the one corresponding to the s orbital patternwhere the alkaline earth metals all have their valence electrons.

VALENCE ELECTRONS:

Electrons that occupy the highest energy levels in anatom, and which are involved in chemical bonding.

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Alkaline Earth Metals

Alkaline earth metals

On the Periodic table , Group 2 (IIA) consists of beryllium, magnesium, calcium, strontium, barium, and radium. This family of elements is known as the alkaline earth metals , or just the alkaline earths. Although early chemists gave the name "earths" to a group of naturally occurring substances that were unaffected by heat and insoluble in water , the alkaline earth metals are also usually found in the continental crust . In contrast, Group 1 compounds and ions tend to concentrate in the ocean.

Calcium carbonate is geologically evident as limestone , marble , coral, pearls, and chalkall derived mainly from the shells of small marine animals. The weathering of calcium silicate rocks over millions of years converted the insoluble calcium silicate into soluble calcium salts, which were carried to the oceans . The dissolved calcium was used by marine organisms to form their shells. When the organisms died, the shells were deposited on the ocean floor where they were eventually compressed into sedimentary rock . Collisions of tectonic plates eventually allow this rock to rise above the ocean floor to become "land-based" limestone deposits.

Caverns throughout the world are formed by the action of atmospheric carbonic acid (water plus carbon dioxide ) on limestone to form the more soluble calcium bicarbonate. When the solution of calcium bicarbonate reaches the open cavern and the water evaporates, carbon dioxide is released and calcium carbonate remains. The calcium carbonate is deposited as stalagmites if the drops hit the ground before evaporating, or as stalactites if the water evaporates while the drop hangs from above.

Other minerals of alkaline earth metals are beryllium aluminum silicate (beryl), calcium magnesium silicate (asbestos), potassium magnesium chloride (carnallite), calcium magnesium carbonate (dolomite ), magnesium sulfate (epsomite), magnesium carbonate (magnesite), hydrogen magnesium silicate (talc), calcium fluoride (fluorspar), calcium fluorophosphate (fluorapatite), calcium sulfate (gypsum ), strontium sulfate (celestite), strontium carbonate (strontianite), barium sulfate (barite), and barium carbonate (witherite). Radium compounds occur in pitchblende, which is primarily uranium oxide, because radium is a product of the radioactive disintegration of U-238. Most pitchblende in the United States is found in Colorado.

The alkaline earth metals, like the alkali metals, are too reactive to be found in nature except as their compounds; the two valence electrons completing an s-subshell are readily lost, and ions with +2 charges are formed. The alkaline earth metals all have a silver luster when their surfaces are freshly cut, but, except for beryllium, they tarnish rapidly. Like most metals, they are good conductors of electricity .

Only magnesium and calcium are abundant in Earth's crust. Magnesium is found in seawater and as the mineral carnallite, a combination of potassium chloride and magnesium chloride. Calcium carbonate exists as whole mountain ranges of chalk, limestone, and marble. Its most abundant mineral is feldspar , which accounts for two-thirds of the earth's crust. Beryllium is found as the mineral beryl, a beryllium aluminum silicate. With a chromium-ion impurity, beryl is known as emerald. If iron ions are present, the gemstone is blue-green and known as aquamarine.

Beryllium is lightweight and as strong as steel. It is hard enough to scratch glass . Beryllium is used for windows in xray apparatus and in other nuclear applications, allowing the rays to pass through with minimum absorption.

Because beryllium is rather brittle, it is often combined with other metals in alloys. Beryllium-copper alloys have unusually high tensile strength and resilience, which makes them ideal for use in springs and in the delicate parts of many instruments. The alloy does not spark, and so finds use in tools employed in fire-hazard areas. Because beryllium-nickel alloys resist corrosion by salt water, they are used in marine engine parts.

Magnesium, alone or in alloys, replaces aluminum in many construction applications because the supply of this metal from seawater is virtually unlimited. Magnesium is soft and can be machined, cast, and rolled. Magnesium-aluminum alloys (trade name Dowmetal) are often used in airplane construction.

Magnesium hydroxide is used as milk of magnesia for upset stomachs. Epsom salts are magnesium sulfate. Soapstone, a form of talc, is used for laboratory table tops and laundry tubs. Magnesium oxide is used for lining furnaces.

Slaked lime, or calcium hydroxide, is the principal ingredient in plaster and mortar, in which the calcium hydroxide is gradually converted to calcium carbonate by reaction with the carbon dioxide in the air. Slaked lime is an important flux in the reduction of iron in blast furnaces. It is also used as a mild germ-killing agent in buildings that house poultry and farm animals, in the manufacture of cement and sodium carbonate, for neutralizing acid soil , and in the manufacture of glass.

Calcium carbide, made by reacting calcium oxide with carbon in the form of coke, is the starting material for the production of acetylene. Calcium propionate is added to foods to inhibit mold growth. Calcium carbonate and calcium pyrophosphate are ingredients in toothpaste.

Plaster of Paris is 2CaSO4H2O, which forms CaSO42H2O (gypsum), as it sets. Gypsum is used to make wallboard, or sheet rock. Asbestosno longer used as a building material in the United States because of concerns that exposure to asbestos fibers can cause canceris a naturally occurring mineral, a calcium magnesium silicate. Calcium and magnesium chlorides, byproducts of sodium chloride purification, are used in the de-icing of roads. Calcium chloride absorbs water from the air, so is used in the prevention of dust on roads, coal , and tennis courts and as a drying agent in the laboratory.

Florapatite, a calcium fluorophosphate, is an important starting material in the production of phosphoric acid, which, in turn, is used to manufacture fertilizers and detergents. The mines in Florida account for about one-third of the world's supply of this phosphate rock. Fluorspar, or calcium fluoride, is used as a flux in the manufacture of steel. It is also used to make hydrofluoric acid, which is then used to make fluorocarbons such as Teflon.

Calcium is involved in the function of nerves and in blood coagulation. Muscle contraction is regulated by the entry or release of calcium ions by the cell. Calcium phosphate is a component of bones and teeth. Hydroxyapatite, calcium hydroxyphosphate, is the main component of tooth enamel. Cavities are formed when acids decompose this apatite coating. Adding fluoride to the diet converts the hydroxyapatite to a more acid-resistant coating, fluorapatite or calcium fluorophosphate. Magnesium is the metal ion in chlorophyll, the substance in plants that initiates the photosynthesis process in which water and carbon dioxide are converted to sugars. Calcium ions are needed in plants for cell division and cell walls. Calcium pectinate is essential in holding plant cells together. Calcium and magnesium ions are required by living systems, but the other Group 2 elements are generally toxic.

The word barium comes from the Greek barys, meaning heavy. Barium salts are opaque to x rays, and so a slurry of barium sulfate is ingested in order to outline the stomach and intestines in x-ray diagnosis of those organs. Although barium ions are poisonous, the very low solubility of barium sulfate keeps the concentration low enough to avoid damage.

Both barium and strontium oxides are used to coat the filaments of vacuum tubes, which are still used in some applications. Because these elements act to remove traces of oxygen and nitrogen, a single layer of barium or strontium atoms on a filament may increase the efficiency more than a hundred million times.

Radium is a source of radioactive rays traditionally used in cancer treatment, though other radioactive isotopes are now more commonly used. A radioactive isotope of strontium, strontium-90, is a component of nuclear fallout.

The alkaline earths and their compounds burn with distinctive colors. The green of barium, the red of strontium, and the bright white of magnesium are familiar in fireworks. Strontium is also used in arc lamps to produce a bright red light for highway flares.

See also Chemical bonds and physical properties; Chemical elements; Geochemistry; Stalactites and stalagmites

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Alkaline Earth Metals

Alkaline earth metals

The elements that make up Group 2 of the periodic table are commonly called the alkaline earth metals. They include beryllium, magnesium, calcium, strontium, barium, and radium. All of these elements contain two electrons in the outermost energy level of their atoms, and they tend to have similar chemical and physical properties. Their properties can also be compared to those of the alkali metals, which lie next to them on the periodic table. They are shiny, relatively soft, and white or silvery in color.

Like the alkali metals, the alkaline earth metals react with acids and water to produce hydrogen gas. These reactions, however, are less intense than are those with the alkali metals. Alkaline earth metals also react vigorously with oxygen. Magnesium burns so actively in air, for example, that it is often used in flares because of the brilliant white light it produces during combustion.

Flame tests can be used to identify compounds of the alkaline earth metals. The characteristic colors of these elements are orangish-red for calcium, crimson for strontium, and apple-green for barium. The brilliant colors produced in fireworks displays are often produced by compounds of strontium and barium.

Beryllium

Beryllium ranks number 50 in abundance among the chemical elements. Interestingly enough, it occurs most commonly in gemstones and beautiful minerals such as beryl, emeralds, and aquamarine. The largest crystals of beryl are about a meter in length and weigh up to 60 metric tons.

The most important industrial application of beryllium is in the manufacture of alloys (metal mixtures). In very small amounts, the element adds strength, durability, and temperature stability to alloys. Copper-beryllium alloys make good hand tools in industries that use flammable solvents because the tools do not cause sparks when struck against other objects. Nickel-beryllium alloys are used for specialized electrical connections and various high temperature applications. Beryllium is used instead of glass in X-ray tubes because it lets through more of the X-radiation than glass would.

Beryllium is toxic to humans. Exposure to high concentrations can cause a pneumonia-like condition that can quickly result in death. Long-term exposure to even small concentrations can result in serious health problems, in particular a respiratory problem known as berylliosis.

Magnesium

Magnesium is the sixth most common element in Earth's crust. It occurs in minerals such as dolomite, magnesite, carnallite, asbestos, soapstone, mica, and spinel. The oceans also contain relatively high concentrations of magnesium chloride.

Magnesium performs a critical role in living things because it is a key component of chlorophyll. Chlorophyll is the green pigment that captures the energy of sunlight for storage in plant sugars during photo-synthesis. (Through the process of photosynthesis, plants use light to break down chemical compounds). Chlorophyll is a large molecule called a porphyrin; the magnesium occupies the center of the porphyrin molecule. In the animal kingdom, a similar porphyrin called heme allows hemoglobin to transport oxygen around in the bloodstream; in the case of heme, however, iron rather than magnesium occupies the central place in the porphyrin.

Elemental magnesium is a strong, light metal, particularly when alloyed, or mixed, with other metals like aluminum or zinc. These alloys have many uses in construction, such as in the manufacture of airplane

parts. Alloys of magnesium and the rare earth elements are so temperature resistant that they are used to make car engine parts.

Calcium

Calcium is the third most common metal on Earth, exceeded only by iron and aluminum, and the fifth most common element. Naturally occurring compounds of calcium include limestone, dolomite, marble, chalk, and iceland spar (all forms of calcium carbonate); gypsum (calcium sulfate); fluorite (calcium fluoride); and apatite (calcium fluorophosphate). Compounds of calcium are also found in sea water.

Calcium is an essential nutrient for living organisms. One of its function is the proper development of bones and teeth. Nutritionists say that growing children need about 1.5 grams of calcium every day to maintain good health. Calcium is also needed for the coagulation (clotting) of blood and for maintaining a normal heartbeat and blood pressure.

The industrial applications of calcium are numerous. Both limestone and gypsum have been used in building materials since ancient times; in general, gypsum was used in drier climates. Marble is also a good building material. Limestone and dolomite are the principle sources of slaked lime (calcium hydroxide) and quick lime (calcium oxide) for the steel, glass, paper, dairy, and metallurgical industries. Lime can act as an agent to remove impurities from steel, as a neutralizing agent for acidic industrial waste, as a reagent (a chemically active substance) for reclaiming sodium hydroxide from paper pulping waste, and as a scrubbing compound to remove pollutants from smokestack effluent. The paper industry uses calcium carbonate as an additive to give smoothness and opacity (the opposite of transparency) to the finished paper. The food, cosmetic, and pharmaceutical industries use it in antacids, toothpaste, chewing gum, and vitamins.

Strontium, barium, and radium

Strontium and barium are the fifteenth and fourteenth most abundant elements, respectively, in Earth's crust. They also occur in very small concentrations in the oceans. Radium is a radioactive element that occurs only in association with uranium, from which it is formed by radioactive decay. (A radioactive element is one that spontaneously gives off energy in the form of particles or waves by disintegration of their atomic nuclei.) This relationship between uranium and radium provides a reliable way to find the age of rocks. The larger the amount of radium in a rock, the longer decay has been taking place and the older the rock is.

Because of the brilliant red color they produce when burned, strontium compounds are widely used in fireworks and flares. Strontium carbonate is also a glass additive, and strontium hydroxide is a refining agent in the production of beet sugar. The most important commercial application of barium is in the form of barium sulfate, used as a lubricating mud in well-drilling operations. In the medical field, patients with gastrointestinal (stomach and intestinal) problems are often required to drink a chalky, white liquid form of barium sulfate before having X-ray examinations.

Radium was formerly used in medicine to treat various kinds of cancer and other conditions. Its use has declined, however, as safer radioactive materials have been discovered. Compounds of radium were also used to paint the luminous numbers on watch dials. That application has been stopped because of the health risks to workers who used the radium paint.

[See also Chlorophyll; Element, chemical; Periodic table ]

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Alkaline Earth Metals

Alkaline Earth Metals


Alkaline earth metals are the six elements forming Group IIa in the Periodic Table: beryllium (Be), magnesium (Mg), Calcium (Ca), Barium (Ba), Strontium (Sr), and Radium (Ra). Their oxides are basic (alkaline), especially when combined with water. "Earth" is a historical term applied to nonmetallic substances that are insoluble in water and stable to heating, and also the properties of the oxides. Hence, the term "alkali earths" is often used to describe these elements.

See Periodic Table in the For Your Reference section of Volume 1.

Each metal has the electron configuration of an inert (noble) gas plus two electrons in the next higher s orbital. Thus, Mg is 1s 22s 22p 63s 2 or alternatively (Ne)3s 2. The bonds of most compounds of alkali earths are ionic in nature because these outermost electrons are readily lost, forming stable divalent cations. Mg, however, can form compounds with both ionic and covalent bonds , whereas most compounds of Be are covalent. The heavier alkali earths are sometimes compared to Group IIb elements (zinc [Zn], cadmium [Cd], mercury [Hg]) that also have a filled s orbital (5s 2), but the filled 4d 10 orbitals and higher ionization energies of the latter make compounds of Group IIb elements markedly less ionic in character than those of alkali earths.

Mg and Ca are the eighth and sixth most abundant elements in Earth's crust at 2.5 and 3.6 percent, respectively. Be, Sr, and Ba comprise 0.001, 0.025, and 0.05 percent, respectively. Ra is radioactive, and since its longest-lived isotope 226Ra has a half-life of 1,600 years, there is very little Ra in Earth's crust. It is nonetheless present because 226Ra is continuously formed by the decay of uranium (238U). Alkali earth elements are very reactive and strongly reducing in character; thus, none occurs in a free state in the environment. They readily react with oxygen, and the pure metals tarnish in air, forming a surface layer of the oxide. The metals are soluble in liquid ammonia, forming covalent compounds with the general formula M(NH3)6. These solutions are strongly basic and frequently find application in industry.

Oxides of alkali earths were known in ancient times, calcium oxide being lime (from the Latin word calx ). Magnesium oxide or magnesia was also known, its name probably deriving from a district in Asia Minor. Oxides of the other alkali earths were identified in the eighteenth century. Barium oxide or baryta was found in the mineral called heavy spar and given the name barys (from the Greek, meaning "heavy"). Strontia or strontianite (strontium carbonate) was found in a lead mine at Strontian in Scotland. Beryllium oxide was extracted from the mineral beryl (from the Greek word bèryllos ). Be was originally called glucina (from the Greek glykys, meaning "sweet") because of its taste and is sometimes still referred to as glucinum in France.

The English chemist Sir Humphry Davy first isolated Mg, Ca, Sr, and Ba in 1808 by means of electrolysis. (Mg was originally called magnium since Davy had already applied the word "magnesium" to the element manganese.) Be was initially isolated from beryl by the French chemist Antoine Bussy and independently in Germany by Friedrich Wöhler in 1828. The discovery of Ra did not occur until 1898 when Marie and Pierre Curie purified it from barium using its radioactivity. They named it from the Latin word radius (meaning "ray") because the strength of its radioactivity was more than a million times that of uranium.

Because of their metallic properties and low mass, Be and Mg are used to form lightweight alloys for structural purposes. Ca sees less industrial use, although the phosphate is sometimes utilized in fertilizers. Sr and Ba have no significant industrial applications. Both Be and Ra are used in various devices, the former because it is quite transparent to x-rays and the latter because it is a ready source of both α - and γ -radiation. Mg and Ca are essential to all living systems for many reasons; the other alkali earths have no known biological roles.

see also Beryllium; Cesium; Curie, Marie Sklodowska; Davy, Humphry; Francium; Magnesium; Potassium; Rubidium; WÖhler, Friedrich.

Michael E. Maguire

Bibliography

Nechaev, I.; Jenkins, G. W.; and Van Loon, Borin (1997). Chemical Elements: The Exciting Story of Their Discovery and of the Great Scientists Who Found Them. Jersey City, NJ: Parkwest Publications.

Rossotti, Hazel (1998). Diverse Atoms: Profiles of the Chemical Elements. New York: Oxford University Press.

Internet Resource

Winter, Mark (2003). WebElements Periodic Table, Scholar Edition. WebElementsLtd. Additional information available from <http://www.webelements.com>.

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alkaline rock

alkaline rock Igneous rock containing a relatively high concentration of the alkali (lithium, sodium, potassium, rubidium, caesium, and francium) and alkaline earth metals (magnesium, calcium, strontium, barium, and radium). Both silica-saturated and silica-undersaturated varieties exist, expressed in the presence of alkali feldspars and feldspathoids respectively. Alkali ferromagnesian minerals are usually present, and their identity depends on the composition of the rock. Igneous rocks of the alkaline suite span the composition range from basic to acid, and may be intrusive or extrusive.

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alkaline rock

alkaline rock An igneous rock that contains a relatively high concentration of the alkali (lithium, sodium, potassium, rubidium, caesium, and francium) and alkaline earth metals (magnesium, calcium, strontium, barium, and radium). Both silica-saturated and silica-under-saturated varieties exist, expressed in the presence of alkali feldspars and feldspathoids respectively. Alkali ferromagnesian minerals are usually present, and their identity depends on the composition of the rock. Compare acid rock.

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alkaline-earth metals

alkaline-earth metals, metals constituting Group 2 of the periodic table. Generally, they are softer than most other metals, react readily with water (especially when heated), and are powerful reducing agents, but they are exceeded in each of these properties by the corresponding alkali metal. They form divalent compounds. In order of increasing atomic number the alkaline-earth metals are beryllium, magnesium, calcium, strontium, barium, and radium.

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alkaline-earth metals

alkaline-earth metals Bivalent metals forming Group II of the periodic table: beryllium, magnesium, calcium, strontium, barium, and radium. They are all light, soft and highly reactive. All, except beryllium and magnesium, react with cold water to form hydroxides (though magnesium reacts with hot water). Radium is important for its radioactive properties.

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alkaline earths

alkaline earths (ăl´kəlīn, –lĬn), oxides of the alkaline-earth metals, especially of calcium, strontium, and barium. They are not readily soluble in water and form solutions less basic than those of alkalies.

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