CHAPTER 1

The Alkaline Earths as Metals

OUTLINE

1.1. General Properties 1 1.2.3. Calcium 12 1.2.4. Strontium 15 1.2. Properties of the Alkaline Earth 1.2.5. Barium 18 Metals 4 1.2.6. Radium 19 1.2.1. Beryllium 4 1.2.2. Magnesium 8

The alkaline earth metals comprise Group 2 of the 1.1. GENERAL PROPERTIES periodic table and include the elements Be, Mg, Ca, Sr, Ba and Ra. These metals form cations with a formal Like other groups, the members of this family show charge of þ2 in solution and are the second most electro- specific patterns in their electron configuration, espe- positive metals of all of the elements (the alkali metals cially the outermost shells, that results in trends in are the most electropositive). The name of this specific chemical behavior (Table 1.1). group in the periodic table stems from the fact that their Another way to depict the electronic structure of oxides produce basic alkaline solutions and that these these elements is shown in Table 1.2. elements melt at such high temperatures that they All of the alkaline earth metals have two electrons in remain solid (earths) in fires. The alkaline earth metals their outer valence shell, so the energetically preferred provide a good example of group trends in chemical state of achieving a filled electron shell is to lose two þ properties within the periodic table, with well-character- electrons to form doubly charged cations, M2 . The alka- ized homologous behavior as one goes down the group. line earth metals are silver-colored, soft metals that react With the exception of Be and Mg, the metals have readily with halogens to form ionic salts. They also react a distinguishable flame color,ELSEVIER orange-red for Ca, with water, though not as rapidly as the alkali metals, to magenta-red for Sr, green for Ba and crimson-red for Ra. form strongly alkaline (basic) hydroxides. For example,

TABLE 1.1 TABLE 1.2

Z Element No. of electrons/shell Element Symbol Electronic configuration

4 Beryllium 2, 2 Beryllium Be [He]2s2 12 Magnesium 2, 8, 2 Magnesium Mg [Ne]3s2 20 Calcium 2, 8, 8, 2 Calcium Ca [Ar]4s2 38 Strontium 2, 8, 18, 8, 2 Strontium Sr [Kr]5s2 56 Barium 2, 8, 18, 18, 8, 2 Barium Ba [Xe]6s2

88 Radium 2, 8, 18, 32, 18, 8, 2 Radium Ra [Rn]7s2

Encyclopedia of the Alkaline Earth Compounds http://dx.doi.org/10.1016/B978-0-444-59550-8.00001-6 1 Copyright Ó 2013 Elsevier B.V. All rights reserved. 2 1. THE ALKALINE EARTHS AS METALS whereas Na and K react with water at room tempera- These elements are all found in the Earth’s crust, but ture, Mg reacts only with steam and Ca with hot water: not in the elemental form because they are so reactive. Instead, they are widely distributed in rock structures. MgðsolidÞþ2HOðgasÞ 2 The main minerals in which magnesium is found are 0 MgðOHÞ ðsolidÞþH ðgasÞ 2 2 “Carnellite”, “Magnesite” and “Dolomite”. Calcium is Be is an exception. It does not react with water or found in “Chalk”, “Limestone”, “Gypsum” and “Anhy- steam, and its halides are covalent. drite”. Magnesium is the eighth most abundant element The alkaline earth metals are named after their oxides, in the Earth’s crust, and calcium is the fifth. the alkaline earths, whose old-fashioned names were Some of the physical properties of the alkaline earth Beryllia, Magnesia, Lime, Strontia and Baryta. “Earth” metals are shown in Table 1.3. is the old term applied by early chemists to nonmetallic The metals of Group 2 are harder and denser than substances that were insoluble in water and resistant to sodium and potassium, and have higher melting points. heating, properties shared by these oxides. The realiza- These properties are due largely to the presence of two tion that these earths were not elements but compounds valence electrons on each atom, which leads to stronger is attributed to the chemist Antoine Lavoisier. In his metallic bonding than occurs in Group 1. “Traite´ E´ lementaire de Chemie” (Elements of Chemistry) Three of these elements give characteristic colors of 1789, he called them “salt-forming” earth elements. when heated in a flame: Later, he suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. Mg ¼ brilliant white Ca ¼ brick red In 1808, acting on Lavoisier’s idea, Humphrey Davy Sr ¼ crimson became the first to obtain samples of the metals by elec- trolysis of their molten “earths”. In all their compounds, these metals have an oxidation þ If the alkaline earths are compared to the alkalis, number of 2 and, with few exceptions, their compounds many similarities are apparent. The main difference is are ionic in nature. The reason for this can be seen by the electron configuration, which is ns2 for alkaline earth examination of the electron configuration, which always metals and ns1 for alkali metals. But for the alkaline has two electrons in an outer quantum level. These elec- earth metals, the nucleus also contains an additional trons are relatively easy to remove, but removing the positive charge. Also, the elements of Group 2 (alkaline third electron is much more difficult, as it is close to the nucleus and in a filled quantum shell. This results in earths) have much higher melting points and boiling 2þ points compared to those of Group 1 (alkali metals). the formation of M . The ionization energies reflect The alkalis are softer and more lightweight than the this electron arrangement. The first two ionization ener- alkaline earth metals that are much harder and denser. gies are relatively low, and the third very much higher. The second valence electron is very important when it In general, the chemical properties of Group 2 comes to comparing chemical properties of the alkaline elements are dominated by the strong reducing power earth and the alkali metals. The second valence electron of the metals. The elements become increasingly electro- is in the same “sublevel” as the first valence electron. positive as one descends within the Group. In direct contact with oxygen or chlorine gas, little or no reaction Therefore, the Zeff is much greater. This means that the elements of Group 2 have a smaller atomic radius and occurs. However, once started, the reactions with much higher ionization energy than those of Group 1. oxygen and chlorine are vigorous: Even though the Group 2 contains a much higher ioniza- ð Þþ ð Þ 0 ð Þþ tion energy, they still form ionic compoundsELSEVIER containing 2Mg solid O2 g 2MgO solid heat þ ð Þþ ð Þ 0 ð Þþ 2 cations. Beryllium, however, behaves differently. Ca solid Cl2 gas CaCl2 solid heat This is due to the fact that in order to remove two elec- trons from this particular atom, significantly more All the metals except beryllium form oxide layers in þ energy is required. It never forms the Be2 cation and air at room temperature that dulls the surface of the its bonds are polar covalent. metal. Barium is so reactive that it is stored under oil. Atomic and ionic radii increase smoothly down the All of the metals except beryllium reduce water and Group. The ionic radii are all much smaller than the cor- dilute acids to hydrogen: responding atomic radii. This arises because the atom ð Þþ þð Þ 0 ð Þþ ð Þ contains two electrons in an s level relatively far from Mg solid 2H aq Mg aq H2 gas the nucleus. It is these electrons that are removed to form the ion. Remaining electrons are thus in levels Magnesium reacts only slowly with water unless the closer to the nucleus, and in addition the increased effec- water is boiling, but calcium reacts rapidly even at room tive nuclear charge attracts the electrons toward the temperature, and forms a cloudy white suspension of nucleus and decreases the size of the ion. sparingly soluble calcium hydroxide. 1.1. GENERAL PROPERTIES 3

TABLE 1.3

Element Atomic number Relative atomic mass Melting point, C Density in kgm/m3

Be 4 9.012 1551 1847.7 Mg 12 24.31 922 1738 Ca 20 40.08 1112 1550 Sr 38 87.62 1042 2540 Ba 56 137.33 1002 3594

Ionization energies in kJ/mol 1st 2nd 3rd

Be 899.4 1757.1 14,848 Mg 737.7 1450.7 7732.6 Ca 589.7 1145 4910 Sr 549.5 1064.2 4210

Ba 502.8 965.1 3600 Standard electrode D Atomic radius/A˚ Ionic radius/A˚ (M2 ) potentials/V

Be 1.13 0.34 1.85 Mg 1.60 0.78 2.36 Ca 1.97 1.06 2.87 Sr 2.15 1.27 2.89

Ba 2.17 1.43 2.90

Calcium, strontium and barium can reduce hydrogen Calcium hydroxide is known as “slaked lime”. It is gas when heated, forming the hydride: sparingly soluble in water and the resulting mildly alka- line solution is known as “limewater” which is used to ð Þþ ð Þ 0 ð Þ Ca solid H2 gas CaH2 solid test for the acidic gas, carbon dioxide. The Group 2 halides are normally found in the The hot metals are also sufficiently strong reducing hydrated form. They are all ionic except beryllium chlo- agents to reduce nitrogen gas and form nitrides: ride. Anhydrous calcium chloride has such a strong ð Þþ ð Þ 0 ð Þ 3Mg solid N2 gas Mg3N2 solid affinity for water that it is used as a drying agent. Of the elements in this Group only magnesium is Magnesium can reduce, and burn, in carbon dioxide: produced on a large scale. It is extracted from seawater ð Þþ ð Þ 0 ð Þþ ð Þ by the addition of calcium hydroxide, which precipitates 2Mg solid CO2 gas 2MgOELSEVIERsolid C solid out the less soluble magnesium hydroxide. This This means that magnesium fires cannot be extin- hydroxide is then converted to the chloride with HCl, guished using carbon dioxide fire extinguishers. which is electrolyzed in a “Downs Cell” to extract The oxides of alkaline earth metals are normally magnesium metal. The metal is used in flares, tracer prepared by heating the hydroxide or carbonate to bullets and incendiary bombs as it burns with a brilliant release carbon dioxide gas. They have high lattice white light. It has also been alloyed with aluminum to enthalpies and melting points. Peroxides, MO2, are produce a low-density, strong material used in aircraft. known for all these elements except beryllium. It Magnesium oxide has such a high melting point that it þ appears that the Be2 cation is too small to accommodate is used to line furnaces. the peroxide anion. The alkaline earth elements are found in all living Calcium, strontium and barium oxides react with organisms. However, beryllium’s low aqueous solubility water to form hydroxides: means that it is rarely available to biological systems. That is, it has no known role in living organisms. It is ð Þþ ð Þ 0 ð Þ ð Þ CaO solid H2O liq Ca OH 2 solid generally highly toxic if encountered by them. 4 1. THE ALKALINE EARTHS AS METALS

In contrast, magnesium and calcium are ubiquitous TABLE 1.4 and essential to all known living organisms. These Location ppb by weight ppb by atoms elements are involved in more than one role. For example, Mg/Ca ion pumps play a pivotal role in Universe 1 0.1 some cellular processes, where magnesium functions Sun 0.1 0.01 as the active center in some enzymes, while calcium salts take a structural role (e.g. bones and teeth) in animals. Meteorite (carbonaceous) 30 70 Strontium and barium display a lower availability in Crustal rocks 4900 4300 the biosphere. Strontium plays an important role in Seawater 0.0006 0.00041 marine aquatic life, especially hard corals. They use stron- tium to build their exoskeleton. These elements also have Streams 0.1 0.01 some uses in medicine, for example “barium meals” in Humans 0.4 0.3 radio graphic imaging, while strontium compounds are employed in some toothpastes. Radium has a low avail- ability and is highly radioactive, making it toxic to life. susceptible persons. The author has had direct contact with such persons who present skeletal aspects of facial appearance and torso as the disease progresses. 1.2. PROPERTIES OF THE ALKALINE Beryllium is a relatively rare element in both the Earth EARTH METALS and the Universe because it is not formed in conven- tional stellar nucleosynthesis. It more accurately was Each of these metals display specific properties which formed during the “Big Bang”, and later from the action differ from the others but have some characteristics that of cosmic rays on interstellar dust. are nearly the same. The abundance of beryllium is shown in Table 1.4. The beryllium content of the earth’s surface rocks is 1.2.1. Beryllium about 4–6 ppm. Beryllium is a constituent in about 100 out of about 4000 known minerals, the most important The name beryllium comes from the Greek word for of which are “Bertrandite” (Be4Si2O7(OH)2), “Beryl” be´rullos, beryl, and from the Prakrit veruliya, in allusion (Al2Be3Si6O18), “Crysoberyl” (Al2BeO4), and “Phena- “to become pale”, in reference to the pale semiprecious kite” (Be2SiO4). Precious stone forms of beryl are “Aqua- gemstone “Beryl”. For about 160 years, beryllium was marine”, “Bixbite” and “Emerald”. also known as glucinium (with the accompanying chem- Beryllium has one of the highest melting points of any ical symbol Gl), the name coming from the Greek word of the light metals. It has exceptional elastic rigidity for “sweet”, due to the sweet taste of its salts. A bivalent (Young’s modulus ¼ 316 GPa). The modulus of elasticity element, beryllium is found in nature as a combination of beryllium is approximately 50% greater than that of with other elements in minerals. Notable gemstones steel. The combination of this modulus plus beryllium’s which contain beryllium include “Beryl” (Aquamarine, relatively low density gives it an unusually fast conduc- Emerald) and “Crysoberyl”. The free element is a steel- tion of sound at standard conditions (about 12.9 km/s). gray, strong, lightweight, brittle, alkaline earth metal Other significant properties are the high values for with an atomic weight of 9.01218 g/mol. It is primarily specific heat (1925 J/kg K) and thermal conductivity used as a hardening agent in alloys, notably beryllium– (216 W/m K). This makes beryllium the metal with the copper. Structurally, beryllium’s very low density best heat dissipation characteristics per unit weight of all ELSEVIER (1.85 times that of water), high melting point (1278 C), of the metals. In combination with the relatively low coef- high temperature stability, and low coefficient of thermal ficient of linear thermal expansion (11.4 10 6/K), these expansion, make it in many ways an ideal aerospace characteristics ensure that beryllium demonstrates material, and it has been used in rocket nozzles and is a unique degree of dimensional stability when heated. a significant component of future-planned space tele- At STP (standard temperature and pressure), beryllium scopes. Because of its relatively high transparency to resists oxidation when exposed to air (its ability to scratch X-rays and other ionizing radiation types, beryllium glass is due to the formation of a thin layer of the hard metal also has a number of uses as filters and windows oxide BeO). It also resists corrosion by concentrated HNO3. for radiation and particle physics experiments. Beryllium has a large scattering cross section for high- Commercial use of beryllium metal presents technical energy neutrons, thus effectively slowing the neutrons to challenges due to the toxicity (especially by inhalation) of the thermal energy range where the cross section is low beryllium-containing dusts. Beryllium produces a direct (0.008 b). The predominant beryllium isotope, 9Be, also corrosive effect to human tissue, and can cause a chronic undergoes a (n, 2n) neutron reaction to form 8Be, i.e. beryl- life-threatening allergic disease called “Berylliosis” in lium is a neutron multiplier, releasing more neutrons than 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 5

TABLE 1.5 triple-alpha process in helium-fueled stars where more synthesis time is available. 7Be decays by electron capture. Known isotopes of beryllium Therefore, its decay rate is dependent upon its electron Nuclide Z N Isotopic mass Half-life Decay mode configurationda rare occurrence in nuclear decay. 13 5Be 4 1 5.04079 No data available Proton emission The shortest-lived known isotope of beryllium is Be which decays through neutron emission. It has a half-life 6Be 4 2 6.019726 4.06848 10 21 s Alpha decay of 2.7 10 21 s. 6Be is also very short lived with a half- [0.092 MeV] life of 4.96 10 21 s. The exotic isotopes 11Be and 14Be 7 Be 4 3 7.01692983 53.22 days Electron capture are known to exhibit a “nuclear halo”. 2 8Be 4 4 8.00530510 6.72206 10 17 s Alpha decay Beryllium has the electronic configuration [He]2s [6.8 eV] and exhibits only the þ2 oxidation state. The only 9Be 4 5 9.0121822 Stable Stable evidence of a lower valence state of beryllium is in the fact that Be is soluble in BeCl2. The small atomic radius 10Be 4 6 10.0135338 1.51 106 years b-minus decay þ ensures that the Be2 ion is highly polarizing, a fact 11Be 4 7 11.021658 13.81 s b-minus decay leading to significant covalent character in beryllium’s 12Be 4 8 12.026921 21.31 ms b-minus decay bonding within various compounds. Beryllium is 4 coor- 2þ 12 21 dinate in complexes e.g. [Be(H2O)4] and tetrahalober- Be 4 8 12.026921 2.71 10 s Neutron 2 emission yllates, BeX4 . This characteristic is used in analytical techniques for determining Be using EDTA as a ligand 14 Be 4 10 14.04289 4.84 ms b-minus decay which preferentially forms octahedral complexes, thus 3þ 15Be 4 11 15.05346 <200 ns No data absorbing other cations such as Al which might inter-

16 fere in the solvent extraction of a complex formed Be 4 12 16.06192 <200 ns No data þ between Be2 and acetylacetone. Beryllium metal lies above aluminum in the electro- it absorbs. Beryllium is highly permeable to X-rays and chemical series and would be expected to be a reactive neutrons are liberated when it is struck by alpha particles. metal. However it is passivated by an oxide layer and Of beryllium’s isotopes, only 9Be is stable and the others does not react with air or water even at red heat. Once are relatively unstable or rare. It is thus a “mono-nuclide” ignited however, beryllium burns brilliantly in air form- 10 element. “Cosmogenic” Be is produced in the atmo- ing a mixture of BeO and Be3N2. sphere by cosmic ray spallation of oxygen and nitrogen. Beryllium dissolves readily in nonoxidizing acids, 10 Cosmogenic Be accumulates at the soil surface, where such as HCl and H2SO4, but not in nitric acid as this its relatively long half-life (1.51 million years) permits forms the oxide on the surface of the metal. This a long residence time before decaying to 9Be. Thus, 10Be behavior is similar to that of aluminum metal. Another and its daughter products have been used to examine strange feature is that Be is amphoteric. This means soil erosion and soil formation from “regolith” (which is that it has the properties of both an acid and a base. soil formed by material originating through rock weath- The following two reactions show this factor: ering or plant growth), the development of lateritic soils BeðOHÞ ðsolidÞþH SO ðaqÞ 0 BeSO ðsolidÞ as well as variations in solar activity, and the age of ice 2 2 4 4 þ ð Þ cores. Solar activity is inversely correlated with 10Be 2H2O liq production, because the solar wind decreases the flux of ð Þ ð Þþ ð Þ 0 ð Þ ð Þ Be OH solid 2 NaOH aq Na2Be OH aq galactic cosmic rays which reach the Earth (Table 1.5). 2 4 ELSEVIER þ 2 5 ð Þþ ð Þ ð Þ Beryllium-10 is also formed in nuclear explosions by 2Na aq Be OH 4 aq a reaction of fast neutrons with 13C in the carbon dioxide in air, and is one of the historical indicators of past Beryllium, again similarly to aluminum, dissolves in 2 activity at nuclear test sites. warm alkali to form the berylliate anion, Be(OH)4 The fact that 7Be and 8Be are unstable has profound and hydrogen gas. The solutions of salts, e.g. beryllium cosmological consequences as it means that elements sulfate and beryllium nitrate are acidic because of 2þ heavier than beryllium could not have been produced hydrolysis of the [Be(H2O)4] ion. For example: by nuclear fusion in the “Big Bang” since there was insuf- ð Þþ ð Þ 0 ½ ð Þ 2þð Þ ficient time during the nucleosynthesis phase of the Big BeSO4 solid 4H2O liq Be H2O 4 aq 4 Bang expansion to produce carbon by fusion of He þ SO2 ðaqÞ nuclei. The other factor was the relatively low concent- 4 8 ½ ð Þ 2þð Þþ 0 ½ ð Þ ð Þþð Þ rations of Be available because of its short half-life. Be H2O 4 aq H2O Be H2O 3 OH aq Astronomer Fred Hoyle first showed that the energy þ þ H O ðaqÞ levels of 8Be and 12C allow carbon production by a 3 6 1. THE ALKALINE EARTHS AS METALS

The hydrolytic reactions of beryllium(II) ions have Although emeralds and beryl were known to ancient been calorimetrically studied at 25 C in aqueous solu- civilizations, they were first recognized as the same tion and dioxane–water mixtures, both containing mineral (Be3Al2(SiO3)6) by Abbe´ Hau¨ y in 1798. Later 3 3.0 mol/dm Li2ClO4 as a constant ionic medium. On that year, Louis-Nicholas Vauquelin, a French chemist, the basis of the formation constants determined, the discovered that an unknown element was present in enthalpy and entropy changes for the reaction: emeralds and beryl. Attempts to isolate the new element finally succeeded in 1828 when two chemists, 2þ þ 0 ð ð Þ Þ2ðxyÞþ þ þ; xBe yH2O Bex OH y yH Friedrich Wo¨lhler of Germany and A. Bussy of France, independently produced beryllium by reducing beryl- 3þ 3þ were estimated for the Be2OH and Be3(OH)3 lium chloride (BeCl2) with potassium metal in a plat- complexes in aqueous solution and 0.1 mol fraction inum crucible. Today, beryllium is primarily obtained 3þ 3þ dioxane–water mixture and for Be2OH ,Be3(OH)3 , þ from the minerals Beryl (Be3Al2(SiO3)6) and Bertrandite 2 $ $ and Be2(OH)2 complexes in 0.2 mol fraction dioxane– (4BeO 2SiO2 H2O) through a chemical process or water mixture. The enthalpy and entropy changes of þ through the electrolysis of a mixture of molten beryllium formation of the Be (OH) )2(x y) complex in solutions x y chloride (BeCl2) and sodium chloride (NaCl). of various mole fractions of dioxane were obtained and Beryllium metal did not become readily available shown to abide by the following reaction: until 1957. Currently, the metal is produced by reducing þ þ þ BeF with Mg metal. The price on the US market for 2Be2 þ H O 0 Be OH3 þ H 2 2 2 vacuum-cast beryllium ingots was $338 per pound þ Thus, it is clear that the Be2 cation in aqueous solu- ($745/kg) in 2001. tion never appears but is an “aquo-complex”. This illus- The metal, beryllium, has had many uses and applica- trates the amphoteric nature of beryllium salts. tions in Industry. Among these are the following: Beryllium differs from its brothers (or sisters) in • Group 2 in that it usually forms covalent bonds. But, Because of its low atomic number and very low unlike other covalent molecules, it is soluble in organic absorption for X-rays, the oldest and still one of the solvents and is a poor conductor when molten. most important applications of beryllium is in Because of its high affinity for oxygen at elevated radiation windows for X-ray tubes. Extreme demands temperatures and its ability to reduce water when its are placed on purity and cleanliness of Be to avoid oxide film is removed, the extraction of beryllium from artifacts in the X-ray images. Thin beryllium foils are its compounds is very difficult. Although electrolysis used as radiation windows for X-ray detectors, and of a fused mixture of beryllium and sodium fluorides the extremely low absorption minimizes the heating was used to isolate the element in the nineteenth effects caused by high intensity, low energy X-rays century, the metal’s high melting point makes this typical of synchrotron radiation. • process more energy intensive than the corresponding Vacuum-tight windows and beam tubes for radiation production of alkali metals by the Down’s Process. Early experiments on synchrotrons are manufactured in the twentieth century, the production of beryllium by exclusively from beryllium. In scientific setups for the thermal decomposition of BeI2 was investigated various X-ray emission studies, the sample holder is following the success of a similar process for the produc- usually made of beryllium because its emitted X-rays tion of zirconium, but this proved to be uneconomic for have much lower energies (~100 eV) than the X-rays volume production. Beryllium metal is available from most studied materials. • commercially and is never normallyELSEVIER made in the labora- Because of its low atomic number, beryllium is almost tory. Its extraction from ores is complex. transparent to energetic particles. Therefore it is used The mineral beryl, [Be3Al2(SiO3)6] is the most impor- to build the “beam pipe” around the collision region tant source of beryllium. It is roasted with sodium hexa- in “Collider Particle Physics” experiments. Notably fluorosilicate, Na2SiF6, at 700 C to form beryllium all four main detector experiments at the Large fluoride. This salt is water soluble and beryllium may Hadron Collider Accelerator in Berne, Switzerland be precipitated as the hydroxide Be(OH) by adjustment use a beryllium beam pipe. 2 • of the pH to 12. Beryllium’s low density allows collision products to Pure beryllium may be obtained by electrolysis of reach the surrounding detectors without a significant molten BeCl2 containing some NaCl. Salt is added since interaction. Its stiffness allows a powerful vacuum to molten BeCl2 conducts very poorly. Another method be produced within the pipe to minimize interaction involves the reduction of beryllium fluoride with with gases. Its thermal stability allows it to function magnesium at 1300 C: correctly at temperatures of only a few degrees above the absolute zero, and its diamagnetic nature þ 0 þ BeF2 Mg MgF2 Be keeps it from interfering with the complex multipole 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 7

magnet systems used to steer and focus the particle Environmental and toxicity considerations have since beams. led to substitution by other materials. • Due to its stiffness, lightweight, and dimensional • Cross-rolled beryllium sheet is an excellent structural stability over a wide temperature range, beryllium support for printed circuit boards in surface-mounted metal is used in the defense and aerospace industries technology. In critical electronic applications, for lightweight structural components in high-speed beryllium is both a structural support and a heat sink. aircraft, missiles, space vehicles and communication The application also requires a coefficient of thermal satellites. Several liquid-fueled rockets use nozzles of expansion that is well matched to that of alumina and pure beryllium. polyimide glass substrates. The beryllium–beryllium • Beryllium is used as an alloying agent in the oxide composite “E-Materials” have been specially production of beryllium–copper, which contains up to designed for these electronic applications and have 2.5% beryllium. Beryllium–copper alloys are used in the additional advantage that the thermal expansion many applications because of their combination of coefficient can be tailored to match diverse substrate high electrical and thermal conductivity, high strength materials. and hardness, nonmagnetic properties, along with • Due to their non-magnetic properties, beryllium-based good corrosion and fatigue resistance. These tools are often used by military naval personnel when applications include the making of spot-welding working on or around sea-mines, as these devices electrodes, springs, non-sparking tools and electrical often have fuses that detonate on direct magnetic contacts. contact or when influenced by a magnetic field. • The excellent elastic rigidity of beryllium has led to its • Beryllium-based tools are used for maintenance and extensive use in precision instrumentation, e.g. in construction near MRI scanners. Magnetic tools gyroscope inertial guidance systems, and in support would be pulled by the scanner’s strong magnetic structures for optical systems. field. Apart from being difficult to remove once • Beryllium mirrors are a field of particular interest in magnetic items are stuck in the scanner, the induced astronomical applications. Large-area mirrors, missile effect can have dangerous consequences. frequently with a honeycomb support structure, are • In the telecommunications industry, tools made of used. For example, in meteorological satellites where beryllium are used to tune the highly magnetic low weight and long-term dimensional stability are klystrons used for high-power microwave critical, the use of beryllium is essential. Smaller applications. beryllium mirrors are used in optical guidance • Beryllium is used in nuclear weapon designs as the systems and in fire control systems, as in the German outer layer of the “pit” (the core of an implosion “main-battle” tanks of World War II. In these systems, weapondthe fissile material and any reflector or very rapid movement of the mirror is required tamper bonded to it) of the primary stage, surrounding which again dictates low mass and high rigidity. the fissile material. It is a good implosion pusher and Usually the beryllium mirror is coated with hard a very good neutron reflector, and is used in certain electroless nickel that can be more easily polished to moderated reactors such as “Molten-salt Reactors”. a finer optical finish than beryllium. In some • Beryllium is sometimes used as a neutron source in applications, though, the beryllium blank is polished which the beryllium is mixed with an alpha emitter without any coating. This is particularly applicable to such as 210Po, 226Ra, 239Pu or 241Am. a cryogenic operation where any thermal expansion • Beryllium has also been proposed as a cladding mismatch can cause the coatingELSEVIER to buckle. material for nuclear fuel, due to its combination of • The James Webb Space Telescope (JWST), a planned mechanical, chemical, and nuclear properties. infrared space observatory (which will replace, in • Beryllium’s characteristics (low weight and high part, the Hubble Space Observatory), will have 18 rigidity) make it useful as a material for high- hexagonal beryllium sections for its mirrors. Because frequency audio-drivers in audio applications. Until JWST will face a temperature of 33 K, the mirror is recently, most beryllium tweeters used an alloy of made of beryllium, because it is capable of handling beryllium and other metals. This was due to extreme cold better than glass. Beryllium contracts beryllium’s high cost and difficult formability. These and deforms less than glass and remains more challenges, coupled with the high performance of uniform at such temperatures. For the same reason, beryllium, caused some manufacturers to falsely the optics of the Spitzer Space Telescope, launched by claim that they used pure beryllium. Some high-end NASA in 2003, are entirely built of beryllium metal. audio companies now manufacture pure beryllium • An earlier major application of beryllium was in tweeters or speakers using such tweeters. Because brakes for military aircraft because of its hardness, beryllium is many times more expensive than high melting point and exceptional heat dissipation. titanium, hard to shape due to its brittleness, 8 1. THE ALKALINE EARTHS AS METALS

and toxicity if mishandled, these tweeters are TABLE 1.6 limited to high-end and “public address” applications. Physical properties of beryllium metal • Beryllium is an effective p-type dopant in III–V CAS number 7440-41-7 semiconductors. It is widely used in materials such as Phase Solid

GaAs, AlGaAs, InGaAs, and InAlAs grown by 3 “Molecular Beam Epitaxy”. Density at rt 1.85 g/cm • BeO is useful for many applications that require the Liquid density at MP 1.690 g/cm3 combined properties of an electrical insulator, an Melting point 1560 K, 1287 C, 2349 F excellent heat conductor with high strength and hardness, with a very high melting point. Beryllium Heat of fusion 7.895 kJ/mol oxide is frequently used as an insulator base plate in Heat of vaporization 297 kJ/mol high-power transistors in RF transmitters for Specific heat capacity 16.443 J/mol K telecommunications. Beryllium oxide is also being studied for use in increasing the thermal conductivity Electronegativity 1.57 (Pauling scale) of UO2 nuclear fuel pellets. Ionization energies 1st: 899.5 kJ/mol • Beryllium compounds were once used in fluorescent 2nd: 1757 kJ/mol lamps as zinc beryllium silicate phosphor, but this use was discontinued because of berylliosis induced in 3rd: 14,848 kJ/mol the workers manufacturing the phosphor and Atomic radius 1.12 A˚ fluorescent lamps. Although the use of beryllium Covalent radius 0.96 A˚ compounds in fluorescent lighting was discontinued ˚ in 1949, potential for exposure to beryllium still exists Van der Waals radius 1.53 A in the nuclear and aerospace industries and in the Crystal structure Hexagonal refining of beryllium metal and melting of beryllium- Magnetic ordering Diamagnetic containing alloys, the manufacturing of electronic devices, and the handling of other beryllium- Thermal conductivity 200 W/mol K (at 300 K) containing material. Thermal expansion 11.3 mm/mol K (25 C) Early researchers tasted beryllium and its various Speed of sound (thin rod) 12,870 ms (at rt) compounds for sweetness in order to verify its presence. Young’s modulus 287 Gpa Modern diagnostic equipment no longer necessitates this Shear modulus 132 Gpa highly risky procedure and no attempt should be made to ingest this highly toxic substance. Beryllium and its Bulk modulus 130 Gpa compounds should be handled with great care and Poisson ratio 0.032 special precautions must be taken when carrying out Mohs hardness 5.5 any activity that could result in the release of beryllium dust (causing lung cancer is a possible result of pro- Vickers hardness 1670 Mpa longed exposure to beryllium laden dust). This substance Brinell hardness 600 MPa can be handled safely if certain procedures are followed. No attempt should be made to work with beryllium before familiarization with correctELSEVIER handling procedures. 1.2.2. Magnesium A successful test for beryllium in air and on surfaces was developed and published (2006) as an international Magnesium is a Group 2 element (Group IIA in older voluntary consensus standard (ASTM D7202; www. labeling schemes). This element has the symbol Mg, astm.org). The procedure uses dilute ammonium atomic number 12, atomic weight of 24.305 g/mol and bifluoride for dissolution and fluorescence detection common oxidation number þ2. It is the eighth most with beryllium bound to sulfonated hydroxybenzoqui- abundant element in the earth’s crust by mass, although noline, allowing detection up to 100 times lower than ninth in the Universe as a whole. This preponderance of the recommended limit for beryllium concentration in magnesium in the Universe is related to the fact that it is the workplace. Fluorescence increases with increasing easily built up in supernova stars from a sequential beryllium concentration. This procedure has been addition of three helium nuclei to carbon (which in successfully tested on a variety of surfaces and is effec- turn is made from three helium nuclei). Magnesium tive for the dissolution and ultratrace detection of refrac- constitutes about 2% of the Earth’s crust by mass, which tory beryllium oxide and siliceous beryllium (Table 1.6). makes it the eighth most abundant element in the crust. 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 9

TABLE 1.7 TABLE 1.8

Location ppb by weight ppb by atoms Nuclide Z N Isotopic mass Half-life Nuclear spin

Universe 600,000 30,000 19Mg 12 7 19.03547 Not known 1/2 Sun 700,000 30,000 20Mg 12 8 20.018863 90.8 ms 0þ Meteorite (carbonaceous) 120,000,000 100,000,000 21Mg 12 9 21.011713 122 ms (5/2, 3/2)þ Crustal rocks 29,000,000 25,000,000 22Mg 12 10 21.9995738 3.8755 s 0þ Seawater 1,326,000 337,000 23Mg 12 11 22.9941237 11.317 s 3/2þ

Stream 4,100 170 24Mg 12 12 23.98504170 Stable (79%) 0þ Humans 270,000 70,000 25Mg 12 13 24.98583692 Stable 5/2þ 26Mg 12 14 25.982592929 Stable 0þ Magnesium ion’s high solubility in water helps to 27Mg 12 15 26.98434059 9.458 min 1/2þ ensure that it is the third most abundant element dis- 28Mg 12 16 27.9838768 20.91 h 0þ solved in seawater. 29 The name originates from the Greek word for Mg 12 17 28.988600 1.30 s 3/2þ a district in Thessaly called “Magnesia”. It is related to 29Mg 12 17 28.988600 1.30 s 3/2þ the terms “magnetite” and “manganese”, which also 31Mg 12 19 30.996546 230 ms 3/2þ originated from this area, and required differentiation 32 as separate substances. Magnesium is the seventh most Mg 12 20 31.998975 86 ms 0þ abundant element in the Earth’s crust by mass and 33Mg 12 21 33.005254 90.5 ms 7/2 eighth by molarity. It is found in large deposits of 34Mg 12 22 34.00946 20 ms 0þ Magnesite, Dolomite and other minerals, and in mineral 35 waters, where the magnesium ion is soluble. The abun- Mg 12 23 35.01734 70 ms (7/2) dance of magnesium is shown in Table 1.7. 36Mg 12 24 36.02300 3.9 ms 0þ In 1618 a farmer at Epsom in England attempted to 37Mg 12 25 37.03140 40 ms 7/2 give his cows water from a well. They refused to drink 38 because of the water’s bitter taste. However the farmer Mg 12 26 38.03757 1.0 ms 0þ noticed that the water seemed to heal scratches and 39Mg 12 27 39.04677 <260 ns 7/2 rashes. The fame of “Epsom Salts” spread. Eventually 40Mg 12 28 40.05393 1.0 ms 0þ the compound was recognized to be hydrated magne- sium sulfate, MgSO4. The first person to propose that magnesium was an element was Joseph Black of Edinburgh in 1755. In radioactive and in the 1950s to 1970s was made commer- 1792, an impure form of metallic magnesium was cially by several nuclear power plants for use in scien- produced by Anton Rupprecht who heated magnesia tific experiments. This isotope has a relatively short (magnesium oxide, MgO) with charcoal. He named the half-life (21 h) and so its use was limited by shipping element “Austrium” after his native Austria. In 1808, times. 26Mg has found application in isotopic geology, a small sample of the pure metal was isolated by similar to that of aluminum. 26Mg is a radiogenic Humphry Davy by the electrolysisELSEVIER of moist MgO. He daughter product of 26Al, which has a half-life of proposed the name “magnium” based on the mineral 717,000 years. Large enrichments of stable 26Mg have Magnesite (MgCO3) that came from Magnesia in Greece. been observed in the Ca–Al-rich inclusions of some Neither name survived and eventually the metal was carbonaceous chrondrite meteorites. The anomalous called magnesium. The metal itself was first produced abundance of 26Mg is attributed to the decay of its parent in quantity in England by Davy in 1808 using then the 26Al in the inclusions. Therefore, the meteorite must new method of electrolysis of a mixture of molten have formed in the solar nebula before the 26Al had magnesia and mercuric oxide. Antoine Bussy prepared decayed. Hence, these fragments are among the oldest it in a consistent form in 1831. objects in the solar system and have preserved informa- The known isotopes of magnesium are listed in tion about its early history. Table 1.8. It is conventional to plot 26Mg/24Mg against an Al/ Magnesium has three stable isotopes: 24Mg, 25Mg and Mg ratio. In an isochronic dating plot, the Al/Mg ratio 26Mg. All are present in significant amounts (see Table plotted is 27Al/24Mg. The slope of the isochron has no 1.8). About 79% of Mg is 24Mg. The isotope 28Mg is age significance, but indicates the initial 26Al/27Al ratio 10 1. THE ALKALINE EARTHS AS METALS in the sample at the time when the systems were sepa- The metal is now mainly obtained by electrolysis of rated from a common reservoir. magnesium salts obtained from brine. Commercially, Magnesium is a rather tough metal. Elemental magne- the chief use for the metal is as an alloying agent to sium is a moderately strong, silvery-white, lightweight make Al–Mg alloys, sometimes called “magnalium” or metal (two-thirds the density of aluminum). Magnesium “magnelium”. Since magnesium is less dense than tarnishes slightly in air, and finely divided magnesium aluminum, these alloys are valued for their relative readily ignites upon heating in air and burns with lightness and strength. a dazzling white flame, making it a useful ingredient in Magnesium is an important element for plant and flares. Normally, magnesium is coated with a layer of animal life. Chlorophylls are porphyrins (a class of oxide, MgO, that protects magnesium from air and water. pigments including heme and chlorophyll) whose mole- Like its neighbor, Ca, magnesium reacts with water at cules contain a flat ring of four-linked heterocyclic room temperature, though it reacts much more slowly groups, based upon magnesium. The adult human daily than calcium. When it is submerged in water, hydrogen requirement of magnesium is about 0.3 g/day. Magne- bubbles will almost unnoticeably begin to form on the sium is the 11th most abundant element by mass in the surface of the metal. If powdered, it will react much human body. Its ions are essential to all living cells, more rapidly. The reaction occurs much faster at higher where they play a major role in manipulating important temperatures. Magnesium also reacts exothermically biological polyphosphate compounds like ATP, DNA with most acids, such as hydrochloric acid (HCl). As and RNA. Hundreds of enzymes thus require magne- with aluminum, zinc and many other metals, the reac- sium ions in order to function. Magnesium, being tion with hydrochloric acid produces the chloride of the metallic ion at the center of chlorophyll, is thus the metal and releases hydrogen gas. a common additive to fertilizers. Magnesium Magnesium is a highly flammable metal, but, while it compounds are used medicinally as common laxatives, is easy to ignite when powdered or shaved into thin antacids (i.e. “Milk of Magnesia”), and in a number of strips, it is difficult to ignite in mass or bulk. Once situations where stabilization of abnormal nerve excita- ignited, it is difficult to extinguish, being able to burn tion and blood vessel spasm is required (i.e. to treat in both nitrogen (forming magnesium nitride), and eclampsia). Magnesium ions are sour to the taste, and also in CO2 (forming magnesium oxide and carbon). in low concentrations help to impart a natural tartness This property was used in incendiary weapons used in to fresh mineral waters. the “fire bombing” of cities in World War II, the only Magnesium metal can be made commercially by practical civil defense being used to smother a burning several processes and would not normally be made in flare under dry sand to exclude the atmosphere. On the laboratory because of its ready availability. There burning in air, magnesium produces a brilliant white are massive amounts of magnesium in seawater. This light. Thus, magnesium powder (as “flash powder”) can be recovered as magnesium chloride, MgCl2 was used as a source of illumination in the early days through reaction with calcium oxide, CaO: of photography. Later, magnesium ribbon was used in electrically ignited flash bulbs. þ 0 2þ þ CaO H2O Ca 2OH Magnesium powder is used in the manufacture of 2þ þ 0 ð Þ fireworks and marine flares where a brilliant white light Mg 2OH Mg OH 2 is required. Flame temperatures of magnesium and MgðOHÞ þ 2HCl 0 MgCl þ 2H O magnesium alloys can reach 1371 C (2500 F), although 2 2 2 flame height above the burning metal is usually less ELSEVIERElectrolysis of hot molten MgCl2 produces magne- than 300 mm (12 in). Magnesium may be used as an sium as a liquid. This is poured off and chlorine gas is ignition source for “thermite”, or otherwise difficult to recovered: ignite mixture of aluminum and iron oxide powder. þ Magnesium compounds are typically white crystals. cathode: Mg2 ðliqÞþ2e 0 MgðsolidÞ Most are soluble in water, providing the sour-tasting 2þ : ð Þ 0 = ð Þþ magnesium ion, Mg . Small amounts of dissolved anode Cl liq 1 2 Cl2 gas e magnesium ion contribute to the tartness and taste of natural waters. Magnesium ion in large amounts is an The other method used to produce magnesium ionic laxative, and magnesium sulfate (known as involves Dolomite, i.e. MgCa(CO3)2, an important “Epsom Salts”) is sometimes used for this purpose. So- magnesium mineral, and involves a non-electrolytic method. In this process, Dolomite is “calcined” by heat- called “milk of magnesia” is a water suspension of one ¼ $ of the few insoluble magnesium compounds, ing to form the oxide, “calcined dolomite” MgO CaO, and this product is then reacted with ferrosilicon alloy: Mg(OH)2. The undissolved particles give rise to its appearance and name. Milk of magnesia is a mild base ½ $ þ 0 þ þ commonly used as an antacid. 2 MgO CaO FeSi 2Mg Ca2SiO4 Fe 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 11

Magnesium, as the metal, may be distilled out from renewed interest in magnesium engine blocks, as featured this mixture of solid products (BP ¼ 1091 C). Although in the 2006 BMW 325i and 330i models. The application of magnesium is found in over 60 minerals, only a magnesium alloy in the 2006 Corvette engine cradle has Dolomite-(MgCa(CO3)2), Magnesite-(MgCO3), Brucite- advanced the technology of designing robust automotive $ (Mg(OH)2), Carnallite-(KMgCl3 6H2O), Talc-(Mg3Si4O10 partsusingmagnesium.Thesealloys,recentdevelopments (OH)2) and Olivine-((Mg,Fe)2SiO4) are of commercial in high-temperature low-creep metals, are becoming importance. competitivetoaluminumbecauseoflowercosts. þ The Mg2 cation is the second most abundant in The second application field of magnesium is in elec- seawater (occurring at about 12% of the mass of sodium tronic devices. Due to low weight, good mechanical and cations), which makes seawater and sea-salt an attrac- electrical properties, magnesium is widely used for tive commercial source of Mg. To extract the magne- manufacturing of mobile phones, laptop computers, sium, Ca(OH)2 is added to seawater to form Mg(OH)2 cameras, and other electronic components. as a precipitate: Historically,magnesium was one of the main aerospace construction metals and was used for German military ð Þþ ð Þ ð Þ 0 ð Þ ð Þ MgCl2 aq Ca OH 2 solid Mg OH 2 solid aircraft as early as World War I and extensively for þ ð Þ CaCl2 aq German aircraft in World War II. The Germans coined the name “Elektron” for the magnesium alloy that is still Magnesium hydroxide is insoluble in water so it can used today. Due to perceived hazards with magnesium be filtered out and then reacted with hydrochloric acid parts in the event of fire, the application of magnesium to obtain concentrated magnesium chloride: in the commercial aerospace industry was generally ð Þ ð Þþ ð Þ 0 ð Þþ restricted to engine-related components. Currently the Mg OH 2 solid 2 HCl aq MgCl aq 2H2O 2 use of magnesium alloys in aerospace is increasing, mostly driven by the increasing importance of fuel economy and From molten magnesium chloride, an electrolysis the need to reduce weight. The development and testing of process produces magnesium. new magnesium alloys continues, notably Elektron-21, In the United States, magnesium is principally which has successfully undergone extensive aerospace obtained by electrolysis of fused magnesium chloride testing for suitability in engine, internal and airframe from brines, wells, and seawater. The United States has components. The European Community currently runs traditionally been the major world supplier of this metal, three R&D magnesium projects in its Aerospace Priority supplying 45% of world production even as recently as agenda called “Six Framework Program”. 1995. Today, the US market share is at 7%, with a single Magnesium is flammable, burning at a temperature domestic producer left, “US Magnesium”, a company of approximately 1371 C (1644 K; 2500 F). The auto- born from the now-defunct “Magcorp”. As of 2005, ignition temperature of magnesium ribbon is China has taken over as the dominant supplier, pegged approximately 510 C (783 K; 950 F) in air. The high at 60% world market share, which increased from 4% in temperature at which magnesium burns makes it 1995. Unlike the above-described electrolytic process, a handy tool for starting emergency fires during outdoor China is almost completely reliant on a different method recreation. Other related uses include flashlight photog- of obtaining the metal from its ores, the so-called “Silico- raphy, flares, pyrotechnics and fireworks sparklers. thermic Pidgeon” process which involves the reduction Magnesium is also used: of the oxide at high temperatures with silicon. Magnesium is the third most commonly used struc- • To remove sulfur from iron and steel. ELSEVIER• To refine titanium in the “Kroll” process. tural metal, following steel and aluminum. Magnesium, in its purest form, can be compared with aluminum, and • To photoengrave plates in the printing industry. is strong and light so that it is used in several high- • To combine in alloys, where this metal is essential for volume part manufacturing applications, including airplane and missile construction. automotive and truck components. • In the form of turnings or ribbons, to prepare “Grignard Specialty, high-grade car wheels of magnesium alloy Reagents”, which are useful in organic synthesis. are called “mag wheels”. In 1957 a Corvette SS, designed • As an alloying agent, improving the mechanical, for racing, was constructed with magnesium body panels. fabrication and welding characteristics of aluminum. An earlier Mercedes-Benz race car model had a body • As an additive agent in conventional propellants and made from “Elektron”, a magnesium alloy; these cars the production of “nodular graphite” in cast iron. ran (with successes) at Le Mans and other world-class • As a reducing agent for the production of uranium race events in 1955. Volkswagen Group has used magne- and other metals from their salts. sium in its engine components for many years. For a long • As a desiccant, since it easily reacts with water. time, Porsche AG has used a magnesium alloy for its • As a sacrificial (galvanic) anode to protect underground engine blocks due to the weight advantage. There is tanks, pipelines, buried structures, and water heaters. 12 1. THE ALKALINE EARTHS AS METALS

TABLE 1.9 Physical constants of Mg metal are listed in Table 1.9.

Physical constants of magnesium CAS number 7439-95-4 1.2.3. Calcium Phase Solid Calcium is the chemical element with the symbol Ca Density 1.738 g/cm3 and atomic number 20. It has an atomic weight of Liquid density at MP 1.584 g/cm3 40.078 g/mol. Calcium is the fifth most abundant dis- Melting point 923 K, 650 C, 1202 F solved ion in seawater by both molarity and mass, after þ 2þ 2 Na ,Cl,Mg and SO4 . Calcium metal is quite Boiling point 1363 K, 1091 C, 1994 F reactive. It readily forms a white coating of calcium Heat of fusion 8.48 kJ/mol nitride (Ca3N2) in air at room temperature. It reacts Heat of vaporization 128 kJ/mol with water and the metal burns in air with an orange- red flame, forming largely the nitride, but some oxide. Specific heat capacity 24.869 J/mol/K In the visible portion of the spectrum of many stars, Electronegativity 1.31 (Pauling scale) including the Sun, strong absorption lines of singly Atomic radius 1.60 A˚ ionized calcium ions are evident. Prominent among ˚ ˚ these are the H line at 3968.5 A and the K line at Covalent radius 1.41 A 3933.7 A˚ of singly ionized calcium, or Ca II. For the Van der Waals radius 1.73 A˚ Sun, and stars with low temperatures, the prominence Crystal structure Hexagonal of the H and K lines is an indication of strong magnetic activity in the chromosphere. Measurement of periodic Magnetic ordering Paramagnetic variations of these active regions can also be used to Electrical resistivity 43 nU m deduce the rotational periods of these stars. Thermal conductivity 156 W/m K (at 300 K) Calcium as the element is a gray silvery metal. The metal is relatively hard. Calcium is an essential constit- Speed of sound (thin rod) 4940 ms uent of leaves, bones, teeth, and shells in the Earth’s Young’s modulus 45 Gpa environment. Calcium is the fifth most abundant Shear modulus 17 Gpa element in the earth’s crust and makes up more than 3% of the crust by weight. Calcium does not occur as Bulk modulus 46 Gpa the metal itself in nature and instead is found in various Poisson ratio 0.29 minerals including Limestone and Fluorite. Stalagmites Mohs hardness 2.5 and stalactites contain calcium carbonate (CaCO3). Calcium occurs most commonly in sedimentary rocks Brinell hardness 260 Mpa in the minerals “Calcite”, “Dolomite” and “Gypsum”. Stable isotopes It also occurs in igneous and metamorphic rocks

24 24 chiefly in the silicate minerals: “Plagioclase”-(CaAl2- Mg 78.99% Mg is stable with 12 Si O ), “Amphiboles”-(Ca Mg Si O (OH) ), “Pyrox- neutrons 2 8 2 5 8 22 2 enes”-(CaMg(Si,Al)2O6) and “Garnets”-(Ca3Al2(SiO4)3. 25 25 Mg 10% Mg is stable with 13 The abundance of calcium is shown in Table 1.10. ELSEVIERneutrons Calcium (Latin word calcis meaning “lime”) was 26Mg 11.01% 26Mg is stable with 14 known as early as the first century when the Ancient neutrons Romans prepared lime as calcium oxide from limestone and used “slaked lime” as a “whitewash” on various homes and buildings. Literature dating back to 975 AD notes that “Plaster-of-Paris” (calcium sulfate), is useful Magnesium metal and its alloys are explosive for setting broken bones. The metal was not isolated hazards. They are highly flammable in their pure form until 1808 when Sir Humphrey Davy of England electro- when molten or in powder or in ribbon form. Burning lyzed a mixture of lime and mercuric oxide, using then or molten magnesium metal reacts violently with water. the new “Voltaic Cell” as an energy source. Davy was When working with powdered magnesium, safety trying to isolate calcium. When he heard that Swedish glasses for eye protection are employed, because the chemist Jo¨ns Berzelius and his colleague, Pontin, had bright white light produced by burning magnesium prepared calcium amalgam by electrolyzing lime in contains UV light that can permanently damage human mercury, he set up a similar system and was successful. eye retinas. He worked with electrolysis throughout his life and also 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 13

TABLE 1.10 TABLE 1.11

The abundance of calcium Nuclear Nuclide Z N Isotopic mass Half-life spin Location ppb by weight ppb by atoms 34Ca 20 14 34.01412 <35 ns 0þ Universe 70,000 2000 35Ca 20 15 35.00494 25.7 ms 1/2þ Sun 70,000 2000 36Ca 20 16 35.99309 102 ms 0þ Meteorite (carbonaceous) 11,000,000 5,200,000 37Ca 20 17 36.985870 181.1 ms (3/2þ) Crustal rocks 50,000,000 26,000,000 38Ca 20 18 37.976318 440 ms 0þ Seawater 4220 650 39Ca 20 19 38.9707197 859.6 ms 3/2þ Stream 1500 38 40Ca 20 20 39.96259098 Stable >5.9 1021 0þ Humans 14,000,000 2,200,000 years (96.941%) 41Ca 20 21 40.96227806 1.03 105 years 7/2

42Ca 20 22 41.95861801 Stable (0.467%) 0þ discovered/isolated boron, sodium, potassium, magne- 43Ca 20 23 42.9587666 Stable (0.135%) 7/2 sium, calcium, strontium and barium. Calcium metal was not available on a large scale until the beginning 44Ca 20 24 43.9554818 Stable (2.026%) 0þ of the twentieth century. 45Ca 20 25 44.9561866 162.67 days 7/2 Calcium, with a density of 1.55 g/cm3, is the lightest 46Ca 20 26 45.9536926 Stable [>100 1015 0þ of the alkali earth metals. Magnesium (1.74) and beryl- years] (0.004%) lium (1.84) are heavier although they are lighter in 47 atomic mass. Both strontium and barium metals get Ca 20 27 46.9545460 4.536 day 7/2 heavier along with the heavier atomic mass. Calcium 48Ca 20 28 47.952534 43.1 1018 years 0þ has a higher resistivity than copper or aluminum. Yet, (0.187%) weight for weight, allowing for its much lower density, 49Ca 20 29 48.955674 8.718 min 3/2 it is a better conductor than either. However, its use in 50Ca 20 30 49.957519 13.9 s 0þ terrestrial applications is usually limited because of its high reactivity when exposed to air. 51Ca 20 31 50.96151 10.0 s 3/2 Chemically, calcium is reactive and soft for a metal 52Ca 20 32 51.96510 4.6 s 0þ (though harder than lead, it can be cut with a knife 53Ca 20 33 52.97005 90 ms 3/2 with some difficulty). It is a silvery metallic element that must be extracted by electrolysis from a molten 54Ca 20 34 53.97435 50 ms 0þ salt like CaCl2. Once produced, it is not stable when 55Ca 20 35 54.98055 30 ms exposed to air. It is somewhat difficult to ignite, unlike 56Ca 20 36 55.98557 10 ms 0þ magnesium, but when lit, the metal burns in air with a brilliant high-intensity reddish light. 57Ca 20 37 56.99236 5 ms 5/2 Calcium metal reacts with water, evolving hydrogen gas at a rate rapid enough to be noticeable, but not fast enough at room temperature to generateELSEVIER much heat. In produced in the atmosphere, 41Ca is produced by powdered form, however, the reaction with water is neutron activation of 40Ca. Most of its production is in extremely rapid, as the increased surface area of the the upper meter or so of the soil column, where the powder accelerates the reaction with the water. Part of cosmogenic neutron flux is still sufficiently strong. the slowness of the calcium–water reaction results 41Ca has received much attention in stellar studies from the metal being partly protected by an insoluble because it decays to 41K, a critical indicator of solar white Ca(OH)2 layer. In acid solutions where the salt system anomalies. product is water soluble, calcium reacts vigorously. A major part (97%) of naturally occurring calcium is Calcium has 24 known isotopes, as shown in Table 1.11. in the form of 40Ca. This isotope is one of the daughter Calcium has four stable isotopes (40Ca and 42Ca products of 40K decay, along with 40Ar. While K/Ar through 44Ca), plus two more isotopes (46Ca and 48Ca) dating has been used extensively in the geological that have such a long half-lives that for all practical sciences, the prevalence of 40Ca in nature has impeded purposes they can be considered stable. It also has its use in dating. Techniques using mass spectrometry a cosmogenic isotope, radioactive 41Ca, with a half-life and a “double-spike” isotope dilution have been used of 103,000 years. Unlike cosmogenic isotopes that are for K/Ca age dating. 14 1. THE ALKALINE EARTHS AS METALS

The most abundant isotope, 40Ca, has a nucleus of 20 Long-term calcium deficiency can lead to rickets and protons and 20 neutrons. This is the heaviest stable poor blood clotting. In case of a menopausal woman, it isotope of any element known that has equal numbers can lead to “Osteoporosis”, in which the bone deterio- of protons and neutrons. In supernova explosions, rates and an increased risk of fractures develops. While calcium is formed from the reaction of carbon with a lifelong deficit can affect bone and tooth formation, various numbers of alpha particles (helium nuclei), until over-retention can cause “hypercalcemia” (elevated this most common calcium isotope (containing 10 levels of calcium in the blood), impaired kidney function helium nuclei) has been formed. and decreased absorption of other minerals. High Calcium salts are colorless (unless the anion is calcium intakes or high calcium absorption was previ- þ colored) and ionic solutions of calcium (Ca2 ) are color- ously thought to contribute to the development of less as well. Many calcium salts are not soluble in water. kidney stones. However, recent research has indicated When in solution, the calcium ion to the human taste that a high calcium intake is associated with a lower varies remarkably, being reported as mildly salty, sour, risk for kidney stones. Vitamin D is needed to absorb “mineral like” or even “soothing”. It is apparent that calcium. many animals can taste, or develop a taste, for calcium, Calcium metal is readily available commercially. and use this sense to detect the mineral in “salt licks” or Commercially it is made by the electrolysis of molten other sources. In human nutrition, soluble calcium salts calcium chloride, CaCl2: may be added to tart juices without much effect to the : 2þð Þþ 0 ð Þ average palate. cathode Ca liq 2e Ca solid Calcium is essential for living organisms, particu- : ð Þ 0 ð Þþ anode Cl liq 1=2 Cl2 gas e larly in cell physiology, where the movement of the þ calcium ion, Ca2 , into and out of the cytoplasm oper- in the same way that magnesium is produced. The ates as a signal for many cellular processes. As a major calcium chloride is made by the action of hydrochloric material used in mineralization of bones and shells, acid upon calcium carbonate. Calcium chloride is also calcium is the most abundant metal by mass in a by-product in the Solvay process used to make sodium many animals. Calcium is the fifth most abundant carbonate: element by mass in the human body, where it is þ 0 þ þ CaCO3 2 HCl CaCl2 H2O CO2 a common cellular ionic messenger for many bodily functions, and serves also as a structural element. Alternatively, and on a small scale, calcium can be The relatively high atomic number of calcium in the made through the reduction of CaO with aluminum or skeleton causes bone to be radioopaque. Of the human of CaCl2 with sodium metal: body’s solid components after drying (as for example, þ 0 þ after cremation), about a third of the total mass is 6CaO 2Al 3Ca Ca3Al2O6 approximately 1 kg of calcium that composes the þ 0 þ average skeleton (the remainder being mostly phos- CaCl2 2Na Ca 2NaCl phates and oxides). Some important uses of calcium metal are: Calcium is an important component of a healthy diet and a necessary mineral for maintaining life. Calcium • Use as a reducing agent in the extraction of plays an important role in building stronger, denser other metals, such as uranium, zirconium and bones early in life and keeping bones strong and healthy thorium. later in life. Approximately 99% ofELSEVIER the body’s calcium is • Use as a deoxidizer, desulfurizer, or decarbonizer for stored in the bones and teeth. The rest of the calcium in various ferrous and nonferrous alloys. the body has other important uses, such as “neurotrans- • Use in the making of cements and mortars to be used mitter release”, muscle contraction and “exocytosis” in construction. (a durable process by which a cell directs the contents • Calcium is also used to remove oxygen, sulfur and of secretory vesicles out of the cell membrane). Espe- carbon from certain alloys. cially important are “neurotransmitter release”, and • Calcium can be alloyed with aluminum, beryllium, muscle contraction. In the electrical conduction system copper, lead and magnesium. of the heart, calcium replaces sodium as the mineral • Calcium is also used in vacuum tubes as a getter, that depolarizes the cell, thereby proliferating the action a material that combines with and removes trace gases potential. In cardiac muscle, sodium influx commences from vacuum tubes. an action potential, but during potassium efflux, the • Use in dehydrating oils, decarburization and cardiac myocyte experiences calcium influx, prolonging desulfurization of iron and its alloys. the action potential and creating a plateau phase of • Also used in fertilizer, concrete and plaster of paris dynamic equilibrium. (Table 1.12). 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 15

TABLE 1.12 TABLE 1.12 (cont’d)

Physical constants of calcium Physical constants of calcium 6 CAS number 7440-70-2 Conductivity Electrical: 0.298 10 /cm Phase Solid Thermal: 2.01 W/cm K Atomic mass 40.078 Enthalpy of atomization: 184 kJ/mol at 25 C Density 1.55 g/cm3 Enthalpy of vaporization: 153.6 kJ/mol Liquid density at MP 1.378 g/cm3 Enthalpy of fusion: 8.54 kJ/mol Melting point 1115 K; 842 C; 1548 F Vapor pressure 254 Pa at 839 C Boiling point 1757 K; 1448 C; 2703 F Abundance of calcium Earth’s crust: 41.000 ppm Heat of fusion 8.54 kJ/mol Seawater: 390 ppm ¼ 6 Heat of vaporization 154.7 kJ/mol Sun 2.24 10 (relative to H ¼ 1012) Specific heat capacity 25.929 J/mol K Main isotopes of calcium 40Ca ¼ 96.941% 41Ca ¼ trace Electronegativity 1.00 (Pauling scale) 42Ca ¼ 0.647% 43Ca ¼ 0.135% Ionization energies 1st: 589.8 kJ/mol 44Ca ¼ 2.086% 48Ca ¼ 0.187% 2nd: 1145.4 kJ/mol 3rd: 4912.4 kJ/mol Due to its high reactivity with common materials like Atomic radius 2.23 A˚ water, there is very little demand for metallic calcium. Covalent radius 1.76 A˚ Van der Waals radius 2.31 A˚ 1.2.4. Strontium Atomic volume 29.9 cm3/mol Strontium has the symbol Sr, the atomic number 38 Crystal structure Face-centered cubic and an atomic weight of 87.623 g/mol. As an alkaline Magnetic ordering Diamagnetic earth metal, strontium is a soft silver-white or yellowish Electrical resistivity 33.6 nU m (20 C) metallic element that is highly reactive chemically. Due to its extreme reactivity with oxygen and water, this Thermal conductivity 204 W/m K element occurs naturally only in compounds with other Thermal expansion 22.3 mm/m K elements. The metal turns yellow when exposed to air. It occurs naturally in the minerals “Celestine” (SrSO4) and Speed of sound (thin rod) 3810 m/sex (20 C) 90 “strontianite” (SrCO3). The isotope, Sr, is present in Young’s modulus 20 Gpa “radioactive fallout” and has a half-life of 28.90 years. Shear modulus 7.4 Gpa The following table presents the abundance of stron- Bulk modulus 17 Gpa tium. Strontium commonly occurs in nature, the 15th most abundant element on earth, averaging 0.034% in Poisson ratio 0.31 all igneous rock. It is found chiefly as the form of the Mohs hardness 1.75 ELSEVIERsulfate mineral “Celestite” (SrSO4) and the carbonate Brinell hardness 167 Gpa “Strontianite” (SrCO3)(Table 1.13). Of the two, Celestite occurs much more frequently in Cross section (thermal 0.432 b sedimentary deposits of sufficient size to make the neutron capture) development of mining facilities attractive. Strontianite Electrochemical 0.7477 g/amp h is more useful of the two common minerals because equivalent strontium is used most often in the carbonate form, Heat of fusion 8.54 kJ/mol but few deposits have been discovered that are suitable Ionization potential Ist: 6.113 eV for development. The metal can be prepared by electrol- ysis of melted SrCl2 containing a small amount of KCl 2nd: 11.871 eV (to assist conductivity): 3rd: 50.908 eV þ Sr2 þ 2e 0 Sr Coefficient of linear K ¼ 22 10-6 expansion 0 ð Þþ 2Cl Cl2 gas 2e 16 1. THE ALKALINE EARTHS AS METALS

TABLE 1.13 as the oxide forms. Finely powdered strontium metal will ignite spontaneously in air at room temperature. Abundance of strontium Volatile strontium salts impart a crimson color to flames, Location ppb by weight ppb by atoms and these salts are used in pyrotechnics and in the Universe 40 0.6 production of flares. Strontium has four naturally occurring isotopes 84Sr Sun 50 0.7 (0.56%), 86Sr (9.86%), 87Sr (7.0%) and 88Sr (82.58%), but Meteorite (carbonaceous) 8900 2000 there are 33 known isotopes (Tables 1.14 and 1.15). Crustal rocks 360,000 85,000 This element (Sr) has four stable, naturally occurring isotopes: 84Sr (0.56%), 86Sr (9.86%), 87Sr (7.0%) and 88Sr Seawater 8100 570 (82.58%). Only 87Sr is radiogenic since it is produced Streams 60 0.7 by decay from the radioactive alkali metal 87Rb, which 10 Humans 4600 330 has a half-life of 4.88 10 years. Thus, in any material, there are two sources of 87Sr. That formed during primordial nucleosynthesis along with 84Sr, 86Sr and 88Sr, and that formed by radioactive decay of 87Rb. The This is similar to those methods used for the other ratio 87Sr/86Sr is the parameter typically reported in alkaline earth metals. Alternatively it is made by geologic investigations. The ratios reported in minerals reducing strontium oxide with Al metal in vacuum at and rocks have values ranging from 0.7 to greater than a temperature at which strontium distills off. Three allo- 4.0. Because strontium has an electronic configuration tropes of the metal exist, with transition points at 235 C similar to that of calcium, it readily substitutes for Ca (a 0 b) and 540 C(b 0 g). The largest commercially in minerals. exploited deposits are found in England. Sixteen unstable isotopes are known to exist. Of great- Both strontium and “Strontianite” are named after est importance are strontium-89 (89Sr) with a half-life of Strontian, a village in Scotland near which the mineral 50.57 days, and strontium-90 (90Sr) with a half-life of was first discovered in the ores taken from the lead 28.78 years. They decay by emitting an electron and an mines. In 1787, an intriguing mineral came to Edin- anti-neutrino (n ) in beta-minus decay (b decay) to burgh from a Lead mine in a small village on the shores e become yttrium, 90Y (half-life ¼ 64 h). 89Sr is an artificial of Loch Sunart, Argyll, in the western highlands of radioisotope that is used in the treatment of bone cancer. Scotland. At that time, the substance was thought to In circumstances where cancer patients have wide- be some sort of Barium compound. It was 3 years later spread and painful bony metastases, the administration that Scott’s Irish chemist, Adair Crawford, published of 89Sr results in the delivery of b-particles directly to the a paper claiming that the mineral held a new species area of the bony problem, where calcium turnover is including a new chemical element. Other chemists later greatest. 90Sr is a by-product of nuclear fission found prepared a number of compounds with the element, in “nuclear fallout” and presents a health problem since noting that it caused the candle’s flame to burn red, it substitutes for calcium in bone, preventing its expul- while barium compounds gave a green color. The sion from the body. Significant absorption usually new mineral was named “Strontite” in 1793 by Thomas results in death. Hope, another professor of medicine at the University Because it is a long-lived high-energy beta-emitter, of Glasgow. 90Sr is used in SNAP (Systems for Nuclear Auxiliary This element was eventually isolated by Humphrey Power) devices. These devices hold promise for use in Davy in 1808 during his studiesELSEVIER of the electrolysis of spacecraft, remote weather stations, navigational buoys, various “alkaline earths” containing molten chloride etc., where a lightweight, long-lived, nuclear-electric such as SrCl and mercuric oxide. He announced his 2 power source is required. work in a lecture to the Royal Society on 30 June 1808. According to the latest records, China was the top In keeping with the naming of the other alkaline earths, producer of strontium in 2007, with over two-thirds he changed the name to Strontium. world share, followed by Spain and Mexico. Annual Strontium is a gray/silvery metal that is softer than worldwide production is around 137,000 t. Primary Ca. It is even more reactive in water, and reacts on mining areas are UK, Tunisia, Russia, Germany, Mexico contact to produce Sr(OH) and hydrogen gas. It burns 2 and USA. in air to produce both SrO and Sr3N2. But since it does Several uses for radioactive strontium have emerged: not react with nitrogen below 380 C, it only forms the oxide spontaneously at room temperature. It should be • The species, 89Sr, is the active ingredient in Metastron, kept under kerosene to prevent oxidation. Freshly a radiopharmaceutical used for bone pain secondary exposed strontium metal rapidly turns a yellowish color to metastatic bone cancer. The strontium acts like 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 17

TABLE 1.14 TABLE 1.15

Nuclear Physical constants of strontium Nuclide Z N Isotopic mass Half-life spin CAS number 7440-24-6 73Sr 38 35 72.96597 >25 ms 1/2 Phase Solid 74Sr 38 36 73.95631 50 ms 0þ Atomic mass 87.621 g/mol 75Sr 38 37 74.94995 88(3) ms (3/2) Density 2.64 g/cm3 76Sr 38 38 75.94177 7.89 s 0þ Liquid Density at MP 2.375 g/cm3 77Sr 38 39 76.937945 9.0 s 5/2þ Melting point 1050 K; 777 C; 1431 F 78Sr 38 40 77.932180 159 s 0þ Boiling point 1655 K; 1382 C; 2520 F 79Sr 38 41 78.929708 2.25 min 3/2() Vapor pressure at 769 C 253 Pa 80Sr 38 42 79.924521 106.3 min 0þ Heat of fusion 7.43 kJ/mol 81Sr 38 43 80.923212 22.3 min 1/2 Heat of vaporization 136 kJ/mol 82Sr 38 44 81.918402 25.36 days 0þ Specific heat capacity 26.4 J/mol K 83Sr 38 45 82.917557 32.41 h 7/2þ Electronegativity 0.95 (Pauling scale) 84Sr 38 46 83.913425 Stable (0.56%) 0þ Atomic radius 2.45 A˚ 85Sr 38 47 84.912933 64.853 days 9/2þ Covalent radius 1.91 A˚ 86Sr 38 48 85.9092602 Stable (9.86%) 0þ Thermal neutron capture cross section Barns ¼ 1.28 87Sr 38 49 86.9088771 Stable (7.00%) 9/2þ Atomic volume: 33.7 cm3/mol Covalent radius: 1.91 A˚ 88Sr 38 50 87.9056121 Stable (82.58%) 0þ Crystal structure: Cubic face centered Electrochemical 89Sr 38 51 88.9074507 50.57 days 5/2þ equivalent: 1.635 g/amp h 90Sr 38 52 89.907738 28.90 years 0þ Electron work function ¼ 2.59 eV Heat of fusion: 8.3 kJ/mol

91Sr 38 53 90.91020 9.63 h 5/2þ Valence electron potential in eV 25.7 92Sr 38 54 91.911038 2.66 h 0þ Ionization potential Ist: 5.695 93Sr 38 55 92.914026 7.423 min 5/2þ 2nd: 11.03 94Sr 38 56 93.915361 75.3 s 0þ 3rd: 43.6 95Sr 38 57 94.919359 23.90 s 1/2þ Coefficient of linear expansion 23 10 6 (per K) 96Sr 38 58 95.921697 1.07 s 0þ Magnetic ordering Paramagnetic

97Sr 38 59 96.926153 429 ms 1/2þ Electrical resistivity (20 C) 132 nU m 98Sr 38 60 97.928453 0.653 s 0þ Thermal conductivity 35.4 W/m K 99Sr 38 61 98.93324 0.269 s 3/2þ Thermal expansion 22.5 mm/m K 100Sr 38 62 99.93535ELSEVIER 202 ms 0þ Elastic modulus Rigidity: 6.1 Gpa 101Sr 38 63 100.94052 118 ms (5/2) Bulk: 12 Gpa 102Sr 38 64 101.94302 69 ms 0þ Young’s: 15.7 Gpa

103Sr 38 65 102.94895 50 ms Poisson ratio 0.28 104Sr 38 66 103.95233 30 ms 0þ Conductivity Electrical: 0.0795 106/cm 105Sr 38 67 104.95858 20 ms Thermal: 0.353 W/cm K Hardness scale 1.5 Mohs calcium and is preferentially incorporated into bone at Abundance of Sr Earth’s Crust ¼ 370 ppm sites of increased “Osteogenesis”. This localization Seawater ¼ 7.6 ppm focuses the radiation exposure on the cancerous lesion. Sun (relative to H at 12 ¼ • 90Sr has been used as a power source for radioisotope 10 790) thermoelectric generators (RTGs). 90Sr produces about 18 1. THE ALKALINE EARTHS AS METALS

0.93 W of heat per gram (it is lower for the form of 90Sr because it reacts with water and carbon dioxide, it is 90 used in RTGs, which is SrF2). But, Sr has a lifetime not found as a mineral. The most common naturally approximately three times shorter and has a lower occurring minerals are the very insoluble barium 238 density than Pu, another RTG fuel. The main sulfate, BaSO4 (Barite), and barium carbonate, BaCO3 advantage of 90Sr is that it is cheaper than 238Pu and (Witherite). Barium’s name originates from the Greek can be recovered from nuclear waste. word “bary”, meaning “heavy”, describing the high • 90Sr is also used in cancer therapy. Its beta emission density of some common barium-containing ores. and long half-life are ideal for superficial Alchemists in the early Middle Ages knew about some radiotherapy. barium minerals. Smooth pebble-like stones of mineral • 87Sr/86Sr ratios are commonly used to determine the Baryte found in Bolona, Italy were known as “Bologna likely source in areas of sediment in natural systems, Stones”. After exposed to light, they would glow for especially in marine and fluvial environments. Dasch years (probably because they contained some barium (1969) showed that surface sediments of Atlantic sulfide (BaS) formed during the calcination of the displayed 87Sr/86Sr ratios that could be regarded as “stone” with charcoal carbon). It was this quality that bulk averages of the 87Sr/86Sr ratios of geological attracted them to witches and alchemists. terranes from adjacent landmasses. A good example Though barium minerals are dense, barium metal of a fluvial marine system to which Sr isotope itself is comparatively light. Its cosmic abundance is esti- provenance studies have been successfully employed mated as 3.7 atoms (on the same basis, Si ¼ 106 atoms). is the River Nile–Mediterranean Sea system. Barium constitutes about 0.03% of the Earth’s crust, • Due to the differing ages of the rocks that constitute chiefly as the minerals “Barite” (also called barytes or the majority of the Blue and White Nile catchment heavy spar) and “Witherite”. The abundance of barium areas reaching the River Nile delta and East is 0.0425% in the Earth’s crust and 13 mg/l in seawater. Mediterranean Sea, the changing provenance of A rare gem containing barium is known, called sediment can be discerned through Sr isotopic “Benitoite” (BaTiSi3O9). Large deposits of Barite are studies. Such changes have been climatically found in China, Germany, India, Morocco, and in the controlled in the Late Quaternary eras. US. Because barium quickly oxidizes in air, it is difficult • More recently, 87Sr/86Sr ratios have also been used to to obtain the free metal and it is never found free in determine the source of ancient archaeological nature. The following table lists the abundance of materials such as timbers and corn in Chaco Canyon, barium as found in nature (Table 1.16). New Mexico. Barium is a soft and ductile metal. Its simple • 87Sr/86Sr ratios in teeth may also be used to track compounds are notable for their relatively high specific animal migrations or in criminal forensics. gravity (as compared to the other alkaline earth • A recent in vitro study conducted in the NY College of elements). Barium, which is slightly harder than lead, Dental Sciences using strontium to stimulate has a silvery white luster when freshly cut. “osteoblasts” showed marked improvement of The Swedish chemist Carl Wilhelm Scheele discov- bone-building by osteoblasts. ered (1774) a new base (baryta, or barium oxide) as a minor constituent in “Pyrolusite”, but could not isolate As a pure metal, Sr is used in strontium 90%– barium as the metal. From this base, he prepared some aluminum 10% alloys of an eutectic composition for crystals of barium sulfate, which he sent to Johan the modification of aluminum–silicon casting alloys. The AJ62 alloy, a durable creep-resistant magnesium alloy used in car and motorcycle enginesELSEVIER by BMW Motor Car company contains 2% by weight of strontium metal. TABLE 1.16 Other than its usage as radioactive tracers and radio- Abundance of Barium active sources for human body treatment, the only other usage has been in fireworks where the crimson red color Location ppb by weight ppb by atoms is due to strontium. Universe 10 0.09 Sun 10 0.1 1.2.5. Barium Meteorite (carbonaceous) 2800 410 Crustal rocks 340,000 51,000 Barium has the symbol Ba, atomic number 56, and is the fifth element in Group 2. Its atomic weight is Seawater 30 1.4 137.332 g/mol. Barium is a soft silvery metal. It is never Streams 25 0.2 found in nature in its pure form due to its reactivity with Humans 300 14 air. Its oxide is historically known as “Baryta” but 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 19

Gottlieb Gahn, the discoverer of manganese. A month Ba3N2 and BaH2, respectively. Barium reduces oxides, later Gahn found that the mineral Barite is composed chlorides and sulfides of less reactive metals. For of barium sulfate. Only after the electric battery became example: available could Sir Humphry Davy finally isolate (1808) Ba þ CdO 0 BaO þ Cd the element itself by electrolysis. Oxidized barium was þ 0 þ at first called “barote”, by Guyton-de Morveau, a name Ba ZnCl2 BaCl2 Zn that was changed by Lavoisier to baryta. Barium was first 3Baþ Al S 0 3 BaS þ 2Al isolated by electrolysis of molten barium salts in 1808, by 2 3 Davy, who, by analogy with calcium named “barium” When heated with nitrogen and carbon, it forms the after baryta, with the “-ium” ending signifying a metallic cyanide: element. þ þ 0 ð Þ Because barium quickly oxidizes in air, it is difficult to Ba N2 2C Ba CN 2 obtain the free metal and it is never found free in nature. The metal is primarily found in, and extracted from, Barium combines with several metals, including Barite. Because Barite is so insoluble, it cannot be used aluminum, zinc, lead and tin, forming intermetallic directly for the preparation of other barium compounds, compounds and alloys. or barium metal. Instead, the ore is heated with carbon About 40 isotopes of barium have been isolated as to reduce it to barium sulfide: shown in Table 1.17. Naturally occurring barium is a mixture of seven þ 0 þ BaSO4 2C BaS 2CO2 stable isotopes: barium-138 (71.66%), barium-137 (11.32%), barium-136 (7.81%), barium-135 (6.59%), The barium sulfide is then hydrolyzed or treated with barium-134 (2.42%), barium-130 (0.101%), and barium- acids to form other barium compounds, such as the chlo- 132 (0.097%). About six times this many radioactive ride, nitrate or carbonate. isotopes have been prepared with mass numbers ranging Barium is commercially produced through the elec- from 114 to 153. Of the 40 isotopes known, most are trolysis of molten barium chloride (BaCl ): 2 highly radioactive and have half-lives in the several ð Þ 2þð Þþ 0 ð Þ milliseconds to a few days range. The only notable excep- Cathode reaction Ba liq 2e Ba solid tions are 133Ba with a half-life of 10.51 years, 128Ba (2.43 ð Þ 0 ð Þþ 141 140 Anode reaction 2Cl Cl2 gas 2e days), Ba (11.50 days) and Ba (12.75 days). The element is used in metallurgy, and its compounds This process is similar to that of the other alkaline in pyrotechnics, petroleum mining, and radiology. earth metals. Barium metal is also produced by the Metallic barium has few industrial uses. It has been reduction of barium oxide with finely divided historically used to scavenge air in vacuum tubes. There, aluminum at temperatures between 1100 and 1200 C: the metal is used as a getter in electron tubes to perfect the vacuum by combining with final traces of gases. It þ 0 $ þ ð Þ 4 BaO 2Al BaO Al2O3 3Ba gas is also used as a deoxidizer in copper refining, and as a constituent in certain alloys. The alloy with nickel The barium vapor is cooled by means of a water readily emits electrons when heated and, for this reason, jacket and condensed into the solid metal. The solid is used in electron tubes and in spark plug electrodes. block may be cast into rods or extruded into wires. The presence of barium (atomic number 56), observed This is the most effective methodELSEVIER i.e. the reduction of after uranium (atomic number 92) had been bombarded the oxide by heating with aluminum or silicon in by neutrons, was the clue that led to the recognition of a high vacuum, to produce the metal. A mixture of nuclear fission (1939). barium monoxide and peroxide can also be used in the The most important use of elemental barium is as reduction. Being a flammable solid, it is packaged under a scavenger removing last traces of oxygen and other Argon gas in steel containers or plastic bags. Only a few gases in television and other electronic tubes. Addition- tons of the metal are produced each year. ally, an isotope of barium, 133Ba, is routinely used as Barium reacts exothermically with oxygen at room a standard source in the calibration of gamma ray detec- temperature to form both BaO and BaO2. The reaction tors in nuclear physics studies. is violent if the barium is powdered. It also reacts Physical properties of Ba metal shown in Table 1.18. violently with dilute acids, alcohol and water: þ 0 ð Þ þ ð Þ Ba 2H2O Ba OH 2 H2 gas 1.2.6. Radium At elevated temperatures, barium combines with Radium has the symbol Ra and atomic number 88. Its chlorine, nitrogen and hydrogen to produce BaCl2, atomic weight is 226.0254 g/mol. Radium is an alkaline 20 1. THE ALKALINE EARTHS AS METALS

TABLE 1.17 TABLE 1.17 (cont’d)

Nuclide Z N Mass Decay time Spin Nuclide Z N Mass Decay time Spin 148 þ 114Ba 56 58 113.95068 530 ms 0þ Ba 56 92 147.93772 0.612 s 0 149 115Ba 56 59 114.94737 0.45 s (5/2þ) Ba 56 93 148.94258 344 ms 3/2 150 þ 116Ba 56 60 115.94138 1.3 s 0þ Ba 56 94 149.94568 300 ms 0 151 117Ba 56 61 116.93850 1.75 s (3/2)þ Ba 56 95 150.95081 200 ms 3/2 152 þ 118Ba 56 62 117.93304 5.2 s 0þ Ba 56 96 151.95427 100 ms 0 153 119Ba 56 63 118.93066 5.4 s (5/2þ) Ba 56 97 152.95961 80 ms 5/2 120Ba 56 64 119.92604 24 s 0þ 121Ba 56 65 120.92405 29.7 s 5/2(þ) TABLE 1.18 122Ba 56 66 121.91990 1.95 min 0þ Physical properties of barium metal 123Ba 56 67 122.918781 2.7 min 5/2(þ) Name, symbol and atomic Barium, Ba, 56 124 þ Ba 56 68 123.915094 11.0 min 0 number 125 þ Ba 56 69 124.914473 3.5 min 1/2 Atomic weight 137.331 g/mol 126 þ Ba 56 70 125.911250 100 min 0 Phase Solid 127 þ Ba 56 71 126.911094 12.7 min 1/2 Density 3.51 g/mol (20 C) 80.33(12) keV Liquid density at MP 3.338 g/cm3 128 þ Ba 56 72 127.908318 2.43 days 0 Melting point 1000 K; 727 C: 1341 F 129 þ Ba 56 73 128.908679 2.23 h 1/2 Boiling point 2170 K; 1897 C; 3447 F 8.42(6) keV Heat of fusion 7.12 kJ/mol 130 þ Ba 56 74 129.9063208 Stable (0.106%) 0 Heat of vaporization 140.3 kJ/mol [>4.0 1021 years] Specific heat capacity 8.07 J/mol K 131Ba 56 75 130.90694 11.50 days 1/2þ Electronegativity 0.89 (Pauling scale) 187.14(12) keV Ionization energies 1st: 502.9 kJ/mol 132Ba 56 76 131.9050613 Stable (0.101%) 0þ [>300 1018 years] 2nd: 965.1 kJ/mol 133Ba 56 77 132.9060075 10.51 years 1/2þ 3rd: 3600 kJ/mol 134Ba 56 78 133.9045084 Stable (2.417%) 0þ Atom radii Atomic- 2.22 A˚ 135Ba 56 79 134.9056886 Stable (6.592%) 3/2þ Covalent- 2.15 A˚ 136Ba 56 80 135.9045759 Stable (7.854%) 0þ Van der Waals- 2.68 A˚

137Ba 56 81 136.9058274 Stable (11.232%) 3/2 Magnetic ordering Paramagnetic 138 ELSEVIERþ Electrical resistivity 332 num Ba 56 82 137.9052472 Stable (71.698%) 0 139Ba 56 83 138.9088413 83.06 min 7/2 Thermal conductivity 18.5 W/m K 140Ba 56 84 139.910605 12.752 day 0þ Thermal expansion 20.6 mm/m K 141Ba 56 85 140.914411 18.27 min 3/2 Crystal structure Body-centered cubic

142Ba 56 86 141.916453 10.6 min 0þ Speed of sound (thin rod) 1620 ms 143Ba 56 87 142.920627 14.5 s 5/2 Modulus Young’s- 13 GPa 144Ba 56 88 143.922953 11.5 s 0þ Shear- 4.9 GPa 145Ba 56 89 144.92763 4.31 s 5/2 Bulk- 9.6 GPa 146Ba 56 90 145.93022 2.22 s 0þ Mohs hardness 1.25 147Ba 56 91 146.93495 0.893 s (3/2þ) CAS number 7440-39-3 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 21 earth metal that is found in trace amounts in uranium (Ontario), the United States (New Mexico, Utah and ores. Its most stable isotope, 226Ra, has a half-life of Virginia), Australia, and in other places. 1602 years and decays into radon gas. Radium (Ra) has no stable isotopes. A standard The heaviest of the alkaline earth elements, radium is atomic mass cannot be given (but is usually given as intensely radioactive and resembles barium in its chem- 226.0 g/mol). The longest lived, and most common, ical behavior. This metal is found in tiny quantities in the isotope of radium is 226Ra that occurs in the disintegra- uranium ore “Pitchblende”, and various other uranium tion chain of 238U (often referred to as the radium series). minerals. Radium preparations are remarkable for main- Radium (Ra) has 33 different known isotopes, four of taining themselves at a higher temperature than their which are found in nature, with 226Ra being the most surroundings, and for their radiations, which are of common. 223Ra, 224Ra, 226Ra and 228Ra are all generated three kinds: alpha particles, beta particles and gamma naturally in the decay of either Uranium (U) or Thorium rays. (Th). 226Ra is a product of 238U decay, and is the longest- When freshly prepared, pure radium metal is almost lived isotope of radium with a half-life of 1602 years. The pure white, but blackens when exposed to air (probably next longest is 228Ra, a product of 232Th breakdown, with due to nitride formation). Radium is luminescent when a half-life of 5.75 years (Table 1.20). struck by electromagnetic radiation of the proper wave- Radium is over 1 million times more radioactive than length (giving a faint blue color). It reacts violently with the same mass of uranium. Its decay occurs in at least water to form radium hydroxide and is slightly more seven stages. The successive main products have been volatile than barium. The normal phase of radium is studied and were called “radium emanation” or “exra- a solid. Since all the isotopes of radium are radioactive dio” (now identified as radon), radium A (polonium), and short-lived on the geological time scale, any radium B (lead), radium C (bismuth), etc. Radon is primeval radium would have disappeared long ago. a heavy gas in contrast to the others (which are solids). Therefore, radium occurs naturally only as a disintegra- These solid products are themselves radioactive tion product in the three natural radioactive decay series elements, each with an atomic weight a little lower (Thorium, Uranium, and Actinium series). Radium-226 than its predecessor. is a member of the uranium decay series. Its parent is The chemistry of radium is what would be expected Thorium-230 and its daughter Radon-222. The following of the heaviest of the alkaline earths, but the intense lists the known abundance of radium (Table 1.19). radioactivity is its most characteristic property. One Radium is a decay product of uranium and is there- gram of radium-226 undergoes 3.7 1010 disintegra- fore found in all uranium-bearing ores. (One ton of tions per second, producing energy equivalent to Pitchblende yields one seventh of a gram of radium). 6.8 10 3 calories, sufficient to raise the temperature Radium was originally acquired from pitchblende of a well-insulated sample at the rate of 1 C every ore from the Czech Republic. Carnotite 10 s. The practical energy release is even greater than $ (K2(UO2)2(VO4)2 3H2O) sands in Colorado provide this due to the production of a large number of short- some of the element, but richer ores are found in the lived radioactive decay products. The alpha particles Democratic Republic of Congo and the Great Lakes emitted by radium may be used to initiate nuclear reac- area of Canada. Radium can also be extracted from tions. Radium loses about 1% of its activity in 25 years, uranium processing waste. Large radium-containing being transformed into elements of lower atomic weight uranium deposits have been located in Canada with lead being the final product of disintegration. The SI unit of radioactivity is the “Becquerel” (Bq), ELSEVIERequal to one disintegration per second. The “Curie” is TABLE 1.19 a non-SI unit defined as the amount of radioactivity which has the same disintegration rate as 1 g of Ra-226 Abundance of Radium (3.7 1010 disintegrations per second, or 37 GBq). Location ppb by weight ppb by atoms Radium (Latin radius, ray) was discovered by Pierre Curie, Marie Curie, and an assistant, G. Be´mont. This Universe 0 No data occurred after Marie Curie had observed that the radio- Sun No data No data activity of pitchblende was four or five times greater Meteorite (carbonaceous) No data No data than that of the uranium it contained and was not fully explained on the basis of radioactive polonium, which Crustal rocks 0.00010 0.00001 she had just discovered in pitchblende residues origi- Seawater 0.00000001 0.0000000003 nating from North Bohemia, in the Czech Republic. Streams 0.0000004 0.000000002 While studying pitchblende the Curies removed uranium from it and found that the remaining material Humans 0.0000011 0.00000003 was still radioactive. They then separated out 22 1. THE ALKALINE EARTHS AS METALS

TABLE 1.20 TABLE 1.21

Nuclide Z N Isotopic mass Decay time Spin Radium emanation 222Rn 202Ra 88 114 202.00989 2.6 ms 0þ Radium A 218Po 203Ra 88 115 203.00927 4.0 ms (3/2) Radium B 214Pb 204Ra 88 116 204.006500 60 ms 0þ Radium C 214Bi 205 214 Ra 88 117 205.0062 220 ms (3/2 ) Radium C1 Po 206 þ 210 Ra 88 118 206.003827 0.24 s 0 Radium C2 Tl 207Ra 88 119 207.00380 1.3 s (5/2, 3/2) Radium D 210Pb 208Ra 88 120 208.001840 1.3 s 0þ Radium E 210Bi 209Ra 88 121 209.00199 4.6 s 5/2 Radium F 210Po 210Ra 88 122 210.000495 3.7 s 0þ 211Ra 88 123 211.000898 13 s 5/2() 212Ra 88 124 211.999794 13.0 s 0þ In 1910, radium was isolated as a pure metal by Curie 213 and Debierne through the electrolysis of a pure radium Ra 88 125 213.000384 2.74 min 1/2 chloride solution by using a mercury cathode and 214Ra 88 126 214.000108 2.46 s 0þ distilling it in an atmosphere of hydrogen gas. The sepa- 215Ra 88 127 215.002720 1.55 ms (9/2þ)# ration was followed by the increase in intensity of the new lines in the ultraviolet spectrum and by a steady 216Ra 88 128 216.003533 182 ns 0þ increase in the apparent atomic weight of the material 217Ra 88 129 217.006320 1.63 ms (9/2þ) until a value of 225.18 was obtained, remarkably close 218Ra 88 130 218.007140 25.2 ms0þ to the accepted value of 226.03. By 1902, 0.1 g of pure radium chloride was prepared by refining several tons 219Ra 88 131 219.010085 10 ms (7/2)þ of pitchblende residues, and by 1910 Marie Curie and 220Ra 88 132 220.011028 17.9 ms 0þ Andre´-Louis Debierne had isolated the metal itself. 221Ra 88 133 221.013917 28 s 5/2þ Radium was first industrially produced in the begin- ning of the twentieth century by Birac, a subsidiary 222Ra 88 134 222.015375 38.0 s 0þ company of UMHK in its Olen plant in Belgium. This 223Ra 88 135 223.0185022 11.43 days 3/2þ company offered to Marie Curie her first gram of 224Ra 88 136 224.0202118 3.6319 days 0þ radium. Historically the decay products of radium were known as radium A, B, C, etc. These are now 225Ra 88 137 225.023612 14.9 days 1/2þ known to be isotopes of other elements as shown in 226Ra 88 138 226.0254098 1600 years 0þ Table 1.21. 227Ra 88 139 227.0291778 42.2 min 3/2þ On February 4, 1936 radium E became the first radio- active element to be made synthetically. Since all the 228Ra 88 140 228.0310703 5.75 years 0þ isotopes of radium are radioactive and short-lived on 229Ra 88 141 229.034958 4.0(2) min 5/2(þ) the geological time scale, any primeval radium would 230Ra 88 142 230.037056ELSEVIER 93(2) min 0þ have disappeared long ago. Therefore, radium occurs naturally only as a disintegration product in the three 231Ra 88 143 231.041221 103 s (5/2þ) natural radioactive decay series (thorium, uranium, 232Ra 88 144 232.04364 250 s 0þ and actinium series). Radium-226 is a member of the 233Ra 88 145 233.04806 30 s 1/2þ uranium decay series. Its parent is thorium-230 and its daughter radon-222. 234Ra 88 146 234.05070 30 s 0þ Radium was formerly used in self-luminous paints for watches, nuclear panels, aircraft switches, clocks, and instrument dials. More than 100 former watch-dial painters who used their lips to shape the paintbrush a radioactive mixture consisting mostly of barium that died from the radiation from the radium that had produced a brilliant green flame color and crimson- become stored in their bones. Soon afterward, the carmine spectral lines that had never been documented adverse effects of radioactivity became widely known. before. The Curies announced their discovery to the Nevertheless, radium was still used in dials as late as French Academy of Sciences on 26 December 1898. the 1950s. Although the beta-radiation from tritium is 1.2. PROPERTIES OF THE ALKALINE EARTH METALS 23

TABLE 1.22 TABLE 1.23

Physical constants of radium Decay Decay energy Decay CAS number 7440-14-4 Isotope Abundance Half-life mode (MeV) product

Atomic weight 226.0 g/mol 223Ra Trace 11.43 days Alpha 5.99 219Rn 2 Electronic configuration (Rn) 7s 224Ra Trace 3.6319 days Alpha 5.789 220Rn

Phase Solid 226 222 Ra ~100% 1602 years Alpha 4.871 Rn 3 Density at 20 C 5.51 fm/cm 228 228 Ra Trace 5.75 years Beta 0.046 Ac Melting point 973 K: 700 C: 1292 F Boiling point 2010 K; 1727 C; 3159 F children to prevent middle ear problems or enlarged Heat of fusion 8.5 kJ/mol tonsils from the late 1940s through early 1970s. In 1909, the famous Rutherford experiment used Heat of vaporization 113 kJ/mol radium as an alpha source to probe the atomic structure Electronegativity 0.9 (Pauling scale) of gold. This experiment led to the Rutherford model of Ionization energies 1st: 509.3 kJ/mol the atom and revolutionized the field of nuclear physics. Radium (usually in the form of RaCl2) was used in 2nd: 979.3 kJ/mol medicine to produce radon gas which in turn was Covalent radius 2.21 A˚ used as a cancer treatment. For example, several radon Van der Waals radius 2.83 A˚ sources were used in Canada in the 1920s and 1930s. The isotope 223Ra is currently under investigation for Magnetic ordering Non-magnetic its use in cancer treatment of bone metastasis. Electrical resistivity 1.0 mU m Some of the few practical uses of radium are derived Crystal structure Body-centered cubic from its radioactive properties. More recently discov- ered radioisotopes, such as 60Co and 137Cs are replacing Thermal conductivity 18.6 W/m K radium in even these limited uses because several of these isotopes are more powerful emitters, safer to handle, and available in more concentrated form. The potentially dangerous if ingested, it has replaced radium current price for radium metal is ~$40 million per lb in these applications. (Table 1.22). Radium was also put in some foods for taste and as The major isotopes of Radium as a metal are listed in a preservative, but this also exposed many people to Table 1.23. radiation. Radium was once an additive in products In the next chapters, we will survey the properties of like toothpaste, hair creams, and even food items due the alkaline earths as they form compounds. We will to its supposed curative powers. Such products soon begin with the Halides of Group 17 since they are the fell out of vogue and were prohibited by authorities in most electronegative elements in the Periodic Table. many countries, after it was discovered they could This will be followed by a description of the compounds have serious adverse health effects. In the United States, formed with succeeding Groups in the Periodic Chart nasal radium irradiation wasELSEVIER also administered to encompassing Groups 16, 15, 14, and 13.