Alkaline earth metals by Pramod Kothari

Alkaline Earth Metals Lulu Press, Raleigh, N.C. USA

Dr. Pramod Kothari Assistant Professor, Department Of Chemistry Government Post Graduate College, Berinag, District â&#x20AC;&#x201C; Pithoragarh Uttarakhand (India)

Copyright ÂŠ Creative Commons Attribution-Share Alike 3.0 //creativecommons.org/licenses/by-sa/3.0/ Disclaimer All the material contained in this book is provided for educational and informational purposes only. No responsibility can be taken for any results or outcomes resulting from the use of this material. While every attempt has been made to provide information that is both accurate and effective, the author does not assume any responsibility for the accuracy or use/misuse of this information.

Dr. Pramod Kothari / Alkaline Earth Metals, ISBN/EAN13: 978-1-304-87535-8

Preface The alkaline earth metals are a group of chemical elements in the periodic table with very similar properties. They are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure and readily lose their two outermost electrons to form cations with charge 2+ and an oxidation state, or oxidation number of +2. In the modern IUPAC nomenclature, the alkaline earth metals comprise the group 2 elements. The alkaline earth metals are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). This group lies in the s-block of the periodic table as all alkaline earth metals have their outermost electron in an s-orbital. All the discovered alkaline earth metals occur in nature. Experiments have been conducted to attempt the synthesis of element 120, which is likely to be the next member of the group, but they have all met with failure. However, element 120 may not be an alkaline earth metal due to relativistic effects, which are predicted to have a large influence on the chemical properties of superheavy elements.

Dr. Pramod Kothari Assistant Professor, Department Of Chemistry Government Post Graduate College, Berinag, District â&#x20AC;&#x201C; Pithoragarh Uttarakhand (India)

Dr. Pramod Kothari / Alkaline Earth Metals, ISBN/EAN13: 978-1-304-87535-8

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Table of Contents Alkaline earth metal ............................................................................... 1 Electron configuration ......................................................................... 10 Berillium ............................................................................................... 19 Magnesium. ........................................................................................ 35 Calcium ............................................................................................... 47 Strontium. ............................................................................................ 60 Barium ................................................................................................. 71 Radium……………………………………………………………………...81

Dr. Pramod Kothari / Alkaline Earth Metals, ISBN/EAN13: 978-1-304-87535-8

Chapter 1: Alkaline earth metal The alkaline earth metals are a group of chemical elements in the periodic table with very similar properties. They are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure and readily lose their two outermost electrons to form cations with charge 2+ and an oxidation state, or oxidation number of +2. In the modern IUPAC nomenclature, the alkaline earth metals comprise the group 2 elements. The alkaline earth metals are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). This group lies in the s-block of the periodic table as all alkaline earth metals have their outermost electron in an s-orbital. All the discovered alkaline earth metals occur in nature. Experiments have been conducted to attempt the synthesis of element 120, which is likely to be the next member of the group, but they have all met with failure. However, element 120 may not be an alkaline earth metal due to relativistic effects, which are predicted to have a large influence on the chemical properties of superheavy elements. Characteristics

Chemical As with other groups, the members of this family show patterns in their electronic configuration, especially the outermost shells, resulting in trends in chemical behavior: Z Element

No. of electrons/shell Electron configuration

2, 2

[He] 2s

12 magnesium 2, 8, 2

[Ne] 3s

20 calcium

2, 8, 8, 2

[Ar] 4s

38 strontium

2, 8, 18, 8, 2

[Kr] 5s

56 barium

2, 8, 18, 18, 8, 2

[Xe] 6s

88 radium

2, 8, 18, 32, 18, 8, 2

[Rn] 7s

beryllium

Most of the chemistry has been observed only for the first five members of the group. The chemistry of radium is not well-established due to its radioactivity; thus, the presentation of its properties here is limited. The alkaline earth metals are all silver-colored and soft, and have relatively low densities, melting points, and boiling points. In chemical terms, all of the alkaline metals react with the halogens to form the alkaline earth metal halides, all of which being ionic crystalline

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compounds (except for beryllium chloride, which is covalent). All the alkaline earth metals except beryllium also react with water to form strongly alkaline hydroxides and, thus, should be handled with great care. The heavier alkaline earth metals react more vigorously than the lighter ones. The alkaline metals have the second-lowest first ionization energies in their respective periods of the periodic table because of their somewhat low effective nuclear charges and the ability to attain a full outer shell configuration by losing just two electrons. The second ionization energy of all of the alkaline metals is also somewhat low. Beryllium is an exception: It does not react with water or steam, and its halides are covalent. If beryllium did form compounds with an ionization state of +2, it would polarize electron clouds that are near it very strongly and would cause extensive orbital overlap, since beryllium has a high charge density. All compounds that include beryllium have a covalent bond. Even the compound beryllium fluoride, which is the most ionic beryllium compound, has a low melting point and a low electrical conductivity when melted. All the alkaline earth metals have two electrons in their valence shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions. Compounds and reactions The alkaline earth metals all react with the halogens to form ionic halides, such as calcium chloride (CaCl 2), as well as reacting with oxygen to form oxides such as strontium oxide (SrO). Calcium, strontium, and barium react with water to produce hydrogen gas and their respective hydroxides, and also undergo transmetalation reactions to exchange ligands. Alkaline earth metals fluorides solubility-related constants Metal

"MF2"

MF2

unit

lattice

energies

Solubility

2,455

458 3,371

3,526

soluble

1,922

458 2,838

2,978

0.0012

1,577

458 2,493

2,651

0.0002

1,415

458 2,331

2,513

0.0008

1,361

458 2,277

2,373

0.006

Physical and atomic The table below is a summary of the key physical and atomic properties of the alkaline earth metals.

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Alkaline

Standard

Melting Melting Boiling Boiling Density Electronegativity First

earth

atomic

point

metal

weight

(K)

(°C)

(K)

(°C)

(g/cm ) (Pauling)

ionization radius energy

test

color

(pm)

(kJ·mol−1)

(u)

Beryllium

Covalent Flame

9.012182(3) 1560

Magnesium 24.3050(6) 923

1287

2742

2469

1.85

1.57

899.5

105

White

650

1363

1090

1.738

1.31

737.7

150

Brilliantwhite

Calcium

40.078(4)

1115

842

1757

1484

1.54

1.00

589.8

180

Brickred

Strontium

87.62(1)

1050

777

1655

1382

2.64

0.95

549.5

200

Crimson

Barium

137.327(7) 1000

727

2170

1897

3.594

0.89

502.9

215

Applegreen

Radium

[226]

973

700

2010

1737

5.5

0.9

509.3

221

Crimson red

Nuclear stability All of the alkaline earth metals except magnesium and strontium have at least one naturally occurring radioisotope: beryllium-7, beryllium-10, and calcium-41 are trace radioisotopes, calcium-48 and barium-130 have very long half-lives and, thus, occur naturally, and all isotopes of radium are radioactive. Calcium-48 is the lightest nuclide to undergo double beta decay. The natural radioisotope of calcium, calcium-48, makes up about 0.1874% of natural calcium, and, thus, natural calcium is weakly radioactive. Barium-130 makes up approximately 0.1062% of natural barium, and, thus, barium is weakly radioactive, as well.

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History

Etymology The alkaline earth metals are named after their oxides, the alkaline earths, whose oldfashioned names were beryllia, magnesia, lime, strontia, and baryta. These oxides are basic (alkaline) when combined with water. "Earth" is an old term applied by early chemists to nonmetallic substances that are insoluble in water and resistant to heating—properties shared by these oxides. The realization that these earths were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Traité Élémentaire de Chimie (Elements of Chemistry) of 1789 he called them salt-forming earth elements. Later, he suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. In 1808, acting on Lavoisier's idea, Humphry Davy became the first to obtain samples of the metals by electrolysis of their molten earths, thus supporting Lavoisier's hypothesis and causing the group to be named the alkaline earth metals. Discovery The calcium compounds calcite and lime have been known and used since prehistoric times. The same is true for the beryllium compounds beryl and emerald. The other compounds of the alkaline earth metals were discovered starting in the early 15th century. The magnesium compound magnesium sulfate was first discovered in 1618 by a farmer at Epsom in England. Strontium carbonate was discovered in minerals in the Scottish village of Strontian in 1790. The last element is the least abundant: radioactive radium, which was extracted from uraninite in 1898. All elements except beryllium were isolated by electrolysis of molten compounds. Magnesium, calcium, and strontium were first produced by Humphry Davy in 1808, whereas beryllium was independently isolated by Friedrich Wöhler and Antoine Bussy in 1828 by reacting beryllium compounds with potassium. In 1910, radium was isolated as a pure metal by Curie and André-Louis Debierne also by electrolysis. Beryllium

Emerald, a variety of beryl, the mineral that beryllium was first known to be in. Beryl, a mineral that contains beryllium, has been known since the time of the Ptolemaic dynasty in Egypt. Although it was originally thought that beryl was an aluminium silicate, beryl was later found to contain a then-unknown element when, in 1797, Louis-Nicolas Vauquelin dissolved aluminium hydroxide from beryl in an alkali. In 1828, Friedrich Wöhler and Antoine Bussy independently isolated this new element, beryllium, by the same method, which involved a reaction of beryllium chloride with metallic potassium; this reaction was not able to produce large ingots of beryllium. It was not until 1898, when Paul Lebeau performed

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an electrolysis of a mixture of beryllium fluoride and sodium fluoride, that large pure samples of beryllium were produced. Magnesium Magnesium was first produced by Sir Humphry Davy in England in 1808 using electrolysis of a mixture of magnesia and mercuric oxide.Antoine Bussy prepared it in coherent form in 1831. Davyâ&#x20AC;&#x2122;s first suggestion for a name was magnium, but the name magnesium is now used. Calcium Lime has been used as a material for building since 7000 to 14,000 BCE, and kilns used for lime have been dated to 2,500 BCE in Khafaja, Mesopotamia. Calcium as a material has been known since at least the first century, as the ancient Romans were known to have used calcium oxide by preparing it from lime. Calcium sulfate has been known to be able to set broken bones since the tenth century. Calcium itself, however, was not isolated until 1808, when Humphry Davy, in England, used electrolysis on a mixture of lime and mercuric oxide, after hearing that JĂśns Jakob Berzelius had prepared a calcium amalgam from the electrolysis of lime in mercury. Strontium In 1790, physician Adair Crawford, who had been working with barium, realized that Strontian ores showed different properties than other supposed ores of barium. Therefore, he concluded that these ores contained new minerals, which were named strontites in 1793 by Thomas Charles Hope, a chemistry professor at the University of Glasgow, who confirmed Crawford's discovery. Strontium was eventually isolated in 1808 by Sir Humphry Davy by electrolysis of a mixture of strontium chloride and mercuric oxide. The discovery was announced by Davy on 30 June 1808 at a lecture to the Royal Society. Barium

Barite, the material that was first found to contain barium. Barite, a mineral containing barium, was first recognized as containing a new element in 1774 by Carl Scheele, although he was able to isolate only barium oxide. Barium oxide was isolated again two years later by Johan Gottlieb Gahn. Later in the 18th century, William Withering noticed a heavy mineral in the Cumberland lead mines, which are now known to contain barium. Barium itself was finally isolated in 1808 when Sir Humphry Davy used electrolysis with molten salts, and Davy named the element barium, after baryta. Later, Robert Bunsen and Augustus Matthiessen isolated pure barium by electrolysis of a mixture of barium chloride and ammonium chloride.

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Radium While studying uraninite, on 21 December 1898, Marie and Pierre Curie discovered that, even after uranium had decayed, the material created was still radioactive. The material behaved somewhat similarly to barium compounds, although some properties, such as the color of the flame test and spectral lines, were much different. They announced the discovery of a new element on 26 December 1898 to the French Academy of Sciences. Radium was named in 1899 from the word radius, meaning ray, as radium emitted power in the form of rays. Occurrence

Series of alkaline earth metals. Beryllium occurs in the earth's crust at a concentration of two to six parts per million (ppm), much of which is in soils, where it has a concentration of six ppm. Beryllium is one of the rarest elements in seawater, even rarer than elements such as scandium, with a concentration of 0.2 parts per trillion. However, in freshwater, beryllium is somewhat more common, with a concentration of 0.1 parts per billion. Magnesium and calcium are very common in the earth's crust, with calcium the fifth-mostabundant element, and magnesium the eighth. None of the alkaline earth metals are found in their elemental state, but magnesium and calcium are found in many rocks and minerals: magnesium in carnellite, magnesite, and dolomite; and calcium in chalk, limestone, gypsum, and anhydrite. Strontium is the fifteenth-most-abundant element in the Earth's crust. Most strontium is found in the minerals celestite and strontianite. Barium is slightly less common, much of it in the mineral barite. Radium, being a decay product of uranium, is found in all uranium-bearing ores. Due to its relatively short half-life, radium from the Earth's early history has decayed, and present-day samples have all come from the much slower decay of uranium.

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Production

Emerald, a variety of beryl, is a naturally occurring compound of beryllium. Most beryllium is extracted from beryllium hydroxide. One production method is sintering, done by mixing beryl, sodium fluorosilicate, and soda at high temperatures to form sodium fluoroberyllate, aluminium oxide, and silicon dioxide. A solution of sodium fluoroberyllate and sodium hydroxide in water is then used to form beryllium hydroxide by precipitation. Alternatively, in the melt method, powdered beryl is heated to high temperature, cooled with water, then heated again slightly in sulfuric acid, eventually yielding beryllium hydroxide. The beryllium hydroxide from either method then produces beryllium fluoride and beryllium chloride through a somewhat long process. Electrolysis or heating of these compounds can then produce beryllium. In general, strontium carbonate is extracted from the mineral celestite through two methods: by leaching the celestite with sodium carbonate, or in a more complicated way involving coal. To produce barium, barite ore is separated from quartz, sometimes by froth flotation methods, resulting in relatively pure barite. Carbon is then used to reduce the baryte into barium sulfide, which is dissolved with other elements to form other compounds, such as barium nitrate. These in turn are thermally decompressed into barium oxide, which eventually yields pure barium after a reaction with aluminium. The most important supplier of barium is China, which produces more than 50% of world supply. Applications Beryllium is used mostly for military applications, but there are other uses of beryllium, as well. In electronics, beryllium is used as a p-type dopant in some semiconductors, and beryllium oxide is used as a high-strength electrical insulator and heat conductor. Due to its light weight and other properties, beryllium is also used in mechanics when stiffness, light weight, and dimensional stability are required at wide temperature ranges. Magnesium has many different uses. One of its most common uses was in industry, where it has many structural advantages over other materials such as aluminium, although this usage has fallen out of favor recently due to magnesium's flammability. Magnesium is also often alloyed with aluminium or zinc to form materials with more desirable properties than any pure metal. Magnesium has many other uses in industrial applications, such as having a role in the production of iron and steel, and the production of titanium. Calcium also has many uses. One of its uses is as a reducing agent in the separation of other metals from ore, such as uranium. It is also used in the production of the alloys of

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many metals, such as aluminium and copper alloys, and is also used to deoxidize alloys as well. Calcium also has a role in the making of cheese, mortars, and cement. Strontium and barium do not have as many applications as the lighter alkaline earth metals, but still have uses. Strontium carbonate is often used in the manufacturing of red fireworks, and pure strontium is used in the study of neurotransmitter release in neurons. Barium has some use in vacuum tubes to remove gases, and barium sulfate has many uses in the petroleum industry, as well as other industries. Due to its radioactivity, radium no longer has many applications, but it used to have many. Radium used to be used often in luminous paints, although this use was stopped after workers got sick. As people used to think that radioactivity was a good thing, radium used to be added to drinking water, toothpaste, and many other products, although they are also not used anymore due to their health effects. Radium is no longer even used for its radioactive properties, as there are more powerful and safer emitters than radium. Biological role and precautions Magnesium and calcium are ubiquitous and essential to all known living organisms. They are involved in more than one role, with, for example, magnesium or calcium ion pumps playing a role in some cellular processes, magnesium functioning as the active center in some enzymes, and calcium salts taking a structural role, most notably in bones. Strontium plays an important role in marine aquatic life, especially hard corals, which use strontium to build their exoskeletons. It and barium have some uses in medicine, for example "barium meals" in radiographic imaging, whilst strontium compounds are employed in some toothpastes. Excessive amounts of strontium-90 are toxic due to its radioactivity and strontium-90 mimics calcium and then can kill. Beryllium and radium, however, are toxic. Beryllium's low aqueous solubility means it is rarely available to biological systems; it has no known role in living organisms and, when encountered by them, is usually highly toxic. Radium has a low availability and is highly radioactive, making it toxic to life. Extensions The next alkaline earth metal after radium is thought to be element 120, although this may not be true due to relativistic effects. The synthesis of element 120 was first attempted in March 2007, when a team at the Flerov Laboratory of Nuclear Reactions in Dubna bombarded plutonium-244 with iron-58 ions; however, no atoms were produced, leading to a limit of 400 fb for the cross-section at the energy studied. In April 2007, a team at the GSI attempted to create element 120 by bombarding uranium-238 with nickel-64, although no atoms were detected, leading to a limit of 1.6 pb for the reaction. Synthesis was again attempted at higher sensitivities, although no atoms were detected. Other reactions have been tried, although all have been met with failure. The chemistry of element 120 is predicted to be closer to that of calcium or strontium instead of barium or radium. This is unusual as periodic trends would predict element 120 to be more reactive than barium and radium. This lowered reactivity is due to the expected energies of element 120's valence electrons, increasing element 120's ionization energy and decreasing the metallic and ionic radii.

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Bibliography ď&#x201A;ˇ

Weeks, Mary Elvira; Leichester, Henry M. (1968). Discovery of the Elements. Easton, PA: Journal of Chemical Education. LCCCN 68-15217.

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Chapter 2: Electron configuration

Electron atomic and molecular orbitals

A Bohr Diagram of lithium

In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals. 2 2 6 For example, the electron configuration of the neon atom is 1s 2s 2p . Electronic configurations describe electrons as each moving independently in an orbital, in an average field created by all other orbitals. Mathematically, configurations are described by Slater determinants or configuration state functions. According to the laws of quantum mechanics, for systems with only one electron, an energy is associated with each electron configuration and, upon certain conditions, electrons are able to move from one configuration to another by emission or absorption of a quantum of energy, in the form of a photon. For atoms or molecules with more than one electron, the motion of electrons are correlated and such a picture is no longer exact. A very large number of electronic configurations are

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needed to exactly describe any multi-electron system, and no energy can be associated with one single configuration. However, the electronic wave function is usually dominated by a very small number of configurations and therefore the notion of electronic configuration remains essential for multi-electron systems. Electronic configuration of polyatomic molecules can change without absorption or emission of photon through vibronic couplings. Knowledge of the electron configuration of different atoms is useful in understanding the structure of the periodic table of elements. The concept is also useful for describing the chemical bonds that hold atoms together. In bulk materials this same idea helps explain the peculiar properties of lasers and semiconductors. Shells and subshells See also: Electron shell s (ℓ=0) p (ℓ=1) m=0

m=0

m=±1

n=1

n=2

Electron configuration was first conceived of under the Bohr model of the atom, and it is still common to speak of shells and subshells despite the advances in understanding of the quantum-mechanical nature of electrons. An electron shell is the set of allowed states electrons may occupy which share the same principal quantum number, n (the number before the letter in the orbital label). An atom's nth 2 electron shell can accommodate 2n electrons, e.g. the first shell can accommodate 2 electrons, the second shell 8 electrons, and the third shell 18 electrons. The factor of two arises because the allowed states are doubled due to electron spin—each atomic orbital admits up to two otherwise identical electrons with opposite spin, one with a spin +1/2 (usually noted by an up-arrow) and one with a spin −1/2 (with a down-arrow). A subshell is the set of states defined by a common azimuthal quantum number, ℓ, within a shell. The values ℓ = 0, 1, 2, 3 correspond to the s, p, d, and f labels, respectively. The maximum number of electrons which can be placed in a subshell is given by 2(2ℓ + 1). This gives two electrons in an s subshell, six electrons in a p subshell, ten electrons in a d subshell and fourteen electrons in an f subshell. The numbers of electrons that can occupy each shell and each subshell arise from the equations of quantum mechanics, in particular the Pauli exclusion principle, which states that no two electrons in the same atom can have the same values of the four quantum numbers.

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Notation Physicists and chemists use a standard notation to indicate the electron configurations of atoms and molecules. For atoms, the notation consists of a sequence of atomic orbital labels (e.g. for phosphorus the sequence 1s, 2s, 2p, 3s, 3p) with the number of electrons assigned to each orbital (or set of orbitals sharing the same label) placed as a superscript. For example, hydrogen has one electron in the s-orbital of the first shell, so its configuration is 1 written 1s . Lithium has two electrons in the 1s-subshell and one in the (higher-energy) 2s2 1 subshell, so its configuration is written 1s 2s (pronounced "one-s-two, two-s-one"). 2 2 6 2 3 Phosphorus (atomic number 15) is as follows: 1s 2s 2p 3s 3p . For atoms with many electrons, this notation can become lengthy and so an abbreviated notation is used, since all but the last few subshells are identical to those of one or another of 2 2 6 the noble gases. Phosphorus, for instance, differs from neon (1s 2s 2p ) only by the presence of a third shell. Thus, the electron configuration of neon is pulled out, and 2 3 phosphorus is written as follows: [Ne] 3s 3p . This convention is useful as it is the electrons in the outermost shell which most determine the chemistry of the element. The order of writing the orbitals is not completely fixed: some sources group all orbitals with the same value of n together, while other sources (as here) follow the order given by 6 2 Madelung's rule. Hence the electron configuration of iron can be written as [Ar] 3d 4s (keeping the 3d-electrons with the 3s- and 3p-electrons which are implied by the 2 6 configuration of argon) or as [Ar] 4s 3d (following the Aufbau principle, see below). The superscript 1 for a singly occupied orbital is not compulsory. It is quite common to see the letters of the orbital labels (s, p, d, f) written in an italic or slanting typeface, although the International Union of Pure and Applied Chemistry (IUPAC) recommends a normal typeface (as used here). The choice of letters originates from a now-obsolete system of categorizing spectral lines as "sharp", "principal", "diffuse" and "fundamental" (or "fine"), based on their observed fine structure: their modern usage indicates orbitals with an azimuthal quantum number, l, of 0, 1, 2 or 3 respectively. After "f", the sequence continues alphabetically "g", "h", "i"... (l = 4, 5, 6...), skipping "j", although orbitals of these types are rarely required. The electron configurations of molecules are written in a similar way, except that molecular orbital labels are used instead of atomic orbital labels (see below). Energy â&#x20AC;&#x201D; ground state and excited states The energy associated to an electron is that of its orbital. The energy of a configuration is often approximated as the sum of the energy of each electron, neglecting the electronelectron interactions. The configuration that corresponds to the lowest electronic energy is called the ground state. Any other configuration is an excited state. 2

As an example, the ground state configuration of the sodium atom is 1s 2s 2p 3s, as deduced from the Aufbau principle (see below). The first excited state is obtained by 2 2 6 promoting a 3s electron to the 3p orbital, to obtain the 1s 2s 2p 3p configuration, abbreviated as the 3p level. Atoms can move from one configuration to another by absorbing or emitting energy. In a sodium-vapor lamp for example, sodium atoms are excited to the 3p level by an electrical discharge, and return to the ground state by emitting yellow light of wavelength 589 nm. Usually the excitation of valence electrons (such as 3s for sodium) involves energies corresponding to photons of visible or ultraviolet light. The excitation of core electrons is possible, but requires much higher energies generally corresponding to x-ray photons. This would be the case for example to excite a 2p electron to the 3s level and form the excited 2 2 5 2 1s 2s 2p 3s configuration.

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The remainder of this article deals only with the ground-state configuration, often referred to as "the" configuration of an atom or molecule. History Niels Bohr (1923) was the first to propose that the periodicity in the properties of the elements might be explained by the electronic structure of the atom. His proposals were based on the then current Bohr model of the atom, in which the electron shells were orbits at a fixed distance from the nucleus. Bohr's original configurations would seem strange to a 2 2 6 2 4 present-day chemist: sulfur was given as 2.4.4.6 instead of 1s 2s 2p 3s 3p (2.8.6). The following year, E. C. Stoner incorporated Sommerfeld's third quantum number into the description of electron shells, and correctly predicted the shell structure of sulfur to be 2.8.6. However neither Bohr's system nor Stoner's could correctly describe the changes in atomic spectra in a magnetic field (the Zeeman effect). Bohr was well aware of this shortcoming (and others), and had written to his friend Wolfgang Pauli to ask for his help in saving quantum theory (the system now known as "old quantum theory"). Pauli realized that the Zeeman effect must be due only to the outermost electrons of the atom, and was able to reproduce Stoner's shell structure, but with the correct structure of subshells, by his inclusion of a fourth quantum number and his exclusion principle (1925): It should be forbidden for more than one electron with the same value of the main quantum number n to have the same value for the other three quantum numbers k [l], j [ml] and m [ms]. The Schrรถdinger equation, published in 1926, gave three of the four quantum numbers as a direct consequence of its solution for the hydrogen atom: this solution yields the atomic orbitals which are shown today in textbooks of chemistry (and above). The examination of atomic spectra allowed the electron configurations of atoms to be determined experimentally, and led to an empirical rule (known as Madelung's rule (1936), see below) for the order in which atomic orbitals are filled with electrons. Aufbau principle and Madelung rule The Aufbau principle (from the German Aufbau, "building up, construction") was an important part of Bohr's original concept of electron configuration. It may be stated as: a maximum of two electrons are put into orbitals in the order of increasing orbital energy: the lowest-energy orbitals are filled before electrons are placed in higherenergy orbitals.

The approximate order of filling of atomic orbitals, following the arrows from 1s to 7p. (After 7p the order includes orbitals outside the range of the diagram, starting with 8s.)

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The principle works very well (for the ground states of the atoms) for the first 18 elements, then decreasingly well for the following 100 elements. The modern form of the Aufbau principle describes an order of orbital energies given by Madelung's rule (or Klechkowski's rule). This rule was first stated by Charles Janet in 1929, rediscovered by Erwin Madelung in 1936, and later given a theoretical justification by V.M. Klechkowski 1. Orbitals are filled in the order of increasing n+l; 2. Where two orbitals have the same value of n+l, they are filled in order of increasing n. This gives the following order for filling the orbitals: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, (8s, 5g, 6f, 7d, 8p, and 9s) In this list the orbitals in parentheses are not occupied in the ground state of the heaviest atom now known (Uuo, Z = 118). The Aufbau principle can be applied, in a modified form, to the protons and neutrons in the atomic nucleus, as in the shell model of nuclear physics and nuclear chemistry. Periodic table

Electron configuration table The form of the periodic table is closely related to the electron configuration of the atoms of the elements. For example, all the elements of group 2 have an electron configuration of 2 [E] ns (where [E] is an inert gas configuration), and have notable similarities in their chemical properties. In general, the periodicity of the periodic table in terms of periodic table blocks is clearly due to the number of electrons (2, 6, 10, 14...) needed to fill s, p, d, and f subshells. The outermost electron shell is often referred to as the "valence shell" and (to a first approximation) determines the chemical properties. It should be remembered that the similarities in the chemical properties were remarked more than a century before the idea of electron configuration. It is not clear how far Madelung's rule explains (rather than simply describes) the periodic table, although some properties (such as the common +2 oxidation state in the first row of the transition metals) would obviously be different with a different order of orbital filling.

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Shortcomings of the Aufbau principle The Aufbau principle rests on a fundamental postulate that the order of orbital energies is fixed, both for a given element and between different elements; neither of these is true (although they are approximately true enough for the principle to be useful). It considers atomic orbitals as "boxes" of fixed energy into which can be placed two electrons and no more. However the energy of an electron "in" an atomic orbital depends on the energies of all the other electrons of the atom (or ion, or molecule, etc.). There are no "one-electron solutions" for systems of more than one electron, only a set of many-electron solutions which cannot be calculated exactly (although there are mathematical approximations available, such as the Hartreeâ&#x20AC;&#x201C;Fock method). The fact that the Aufbau principle is based on an approximation can be seen from the fact that there is an almost-fixed filling order at all, that, within a given shell, the s-orbital is always filled before the p-orbitals. In a hydrogen-like atom, which only has one electron, the s-orbital and the p-orbitals of the same shell have exactly the same energy, to a very good approximation in the absence of external electromagnetic fields. (However, in a real hydrogen atom, the energy levels are slightly split by the magnetic field of the nucleus, and by the quantum electrodynamic effects of the Lamb shift). Ionization of the transition metals The naĂŻve application of the Aufbau principle leads to a well-known paradox (or apparent paradox) in the basic chemistry of the transition metals. Potassium and calcium appear in the 1 periodic table before the transition metals, and have electron configurations [Ar] 4s and 2 [Ar] 4s respectively, i.e. the 4s-orbital is filled before the 3d-orbital. This is in line with Madelung's rule, as the 4s-orbital has n+l = 4 (n = 4, l = 0) while the 3d-orbital has n+l = 5 5 1 (n = 3, l = 2). However, chromium and copper have electron configurations [Ar] 3d 4s and 10 1 [Ar] 3d 4s respectively, i.e. one electron has passed from the 4s-orbital to a 3d-orbital to generate a half-filled or filled subshell. In this case, the usual explanation is that "half-filled or completely filled subshells are particularly stable arrangements of electrons". The apparent paradox arises when electrons are removed from the transition metal atoms to form ions. The first electrons to be ionized come not from the 3d-orbital, as one would expect if it were "higher in energy", but from the 4s-orbital. The same is true when chemical compounds are formed. Chromium hexacarbonyl can be described as a chromium atom (not ion, it is in the oxidation state 0) surrounded by six carbon monoxide ligands: it is diamagnetic, and the electron configuration of the central chromium atom is described as 6 3d , i.e. the electron which was in the 4s-orbital in the free atom has passed into a 3d-orbital on forming the compound. This interchange of electrons between 4s and 3d is universal among the first series of the transition metals. The phenomenon is only paradoxical if it is assumed that the energies of atomic orbitals are fixed and unaffected by the presence of electrons in other orbitals. If that were the case, the 3d-orbital would have the same energy as the 3p-orbital, as it does in hydrogen, yet it clearly 2+ doesn't. There is no special reason why the Fe ion should have the same electron configuration as the chromium atom, given that iron has two more protons in its nucleus than chromium and that the chemistry of the two species is very different. When care is taken to compare "like with like", the paradox disappears. Other exceptions to Madelung's rule There are several more exceptions to Madelung's rule among the heavier elements, and it is more and more difficult to resort to simple explanations such as the stability of half-filled subshells. It is possible to predict most of the exceptions by Hartreeâ&#x20AC;&#x201C;Fock calculations,

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which are an approximate method for taking account of the effect of the other electrons on orbital energies. For the heavier elements, it is also necessary to take account of the effects of Special Relativity on the energies of the atomic orbitals, as the inner-shell electrons are moving at speeds approaching the speed of light. In general, these relativistic effects tend to decrease the energy of the s-orbitals in relation to the other atomic orbitals. Electron shells filled in violation of Madelung's rule (red) Period 4 Element

Period 5 Z Electron

Element

Configuration

Period 6 Z Electron

Element

Period 7 Z Electron

Configuration

Yttrium

39 [Kr] 5s2 4d1

57 [Xe] 6s25d1

Actinium

89 [Rn] 7s26d1

Cerium

58 [Xe] 6s24f1 5d1 Thorium

90 [Rn] 7s26d2

Protactinium 91 [Rn] 7s25f2 6d1

Neodymium

60 [Xe] 6s2 4f4

Uranium

92 [Rn] 7s25f3 6d1

Promethium

61 [Xe] 6s2 4f5

Neptunium

93 [Rn] 7s25f4 6d1

Samarium

62 [Xe] 6s2 4f6

Plutonium

94 [Rn] 7s2 5f6

Europium

63 [Xe] 6s2 4f7

Americium

95 [Rn] 7s2 5f7

Gadolinium

64 [Xe] 6s24f7 5d1 Curium

96 [Rn] 7s25f7 6d1

Terbium

65 [Xe] 6s2 4f9

97 [Rn] 7s2 5f9

Lutetium

71 [Xe] 6s2 4f14 Lawrencium

Berkelium

Zirconium

40 [Kr] 5s2 4d2

Hafnium

Niobium

41 [Kr] 5s1 4d4

Tantalum

5f 7p

72 [Xe] 6s2 4f14 Rutherfordium 104 [Rn] 7s2 5f14 5d2

Vanadium 23 [Ar] 4s2 3d3

7s2

103 [Rn] 14

5d 22 [Ar] 4s2 3d2

Electron Configuration

Titanium

Lanthanum

Praseodymium 59 [Xe] 6s2 4f3

Scandium 21 [Ar] 4s2 3d1

Element

Configuration

6d2

73 [Xe] 6s2 4f14 5d3

Chromium 24 [Ar] 4s1 3d5

Molybdenum 42 [Kr] 5s1 4d5

Tungsten

74 [Xe] 6s2 4f14 5d4

Manganese 25 [Ar] 4s2 3d5

Technetium 43 [Kr] 5s2 4d5

Rhenium

[Xe] 6s2 4f14 5d5

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Iron

26 [Ar] 4s2 3d6

Cobalt

27 [Ar] 4s2 3d7

Ruthenium 44 [Kr] 5s1 4d7

Rhodium

45 [Kr] 5s1 4d8

Osmium

Iridium

Platinum

Gold

Mercury

[Ar] 4s2 3d8 or Nickel

28 [Ar] 4s1 3d9

Palladium 46

[Kr] 4d10

(disputed)

Copper

29 [Ar] 4s1 3d10

Zinc

30 [Ar] 4s2 3d10

Silver

47 [Kr] 5s1 4d10

Cadmium 48 [Kr] 5s2 4d10

[Xe] 6s2 4f14 5d6 [Xe] 6s2 4f14 5d7

[Xe] 6s1 4f145d9

[Xe] 6s1 4f145d10 [Xe] 6s2 4f14 5d10

The electron-shell configuration of elements beyond rutherfordium has not yet been empirically verified, but they are expected to follow Madelung's rule without exceptions until element 120. Electron configuration in molecules In molecules, the situation becomes more complex, as each molecule has a different orbital structure. The molecular orbitals are labelled according to their symmetry, rather than the atomic orbital labels used for atoms and monatomic ions: hence, the electron configuration of 2 2 2 2 4 2 2 2 the dioxygen molecule, O2, is 1σg 1σu 2σg 2σu 1πu 3σg 1πg . The term 1πg represents the two electrons in the two degenerate π*-orbitals (antibonding). From Hund's rules, these electrons have parallel spins in the ground state, and so dioxygen has a net magnetic moment (it is paramagnetic). The explanation of the paramagnetism of dioxygen was a major success for molecular orbital theory. Electron configuration in solids In a solid, the electron states become very numerous. They cease to be discrete, and effectively blend into continuous ranges of possible states (an electron band). The notion of electron configuration ceases to be relevant, and yields to band theory. Applications The most widespread application of electron configurations is in the chemical properties, in both inorganic and organic chemistry. In configurations, along with some simplified form of molecular orbital theory, modern equivalent of the valence concept, describing the number and bonds that an atom can be expected to form.

rationalization of effect, electron have become the type of chemical

This approach is taken further in computational chemistry, which typically attempts to make quantitative estimates of chemical properties. For many years, most such calculations relied upon the "linear combination of atomic orbitals" (LCAO) approximation, using an ever larger and more complex basis set of atomic orbitals as the starting point. The last step in such a calculation is the assignment of electrons among the molecular orbitals according to the

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Aufbau principle. Not all methods in calculational chemistry rely on electron configuration: density functional theory (DFT) is an important example of a method which discards the model. A fundamental application of electron configurations is in the interpretation of atomic spectra. In this case, it is necessary to convert the electron configuration into one or more term symbols, which describe the different energy levels available to an atom. Term symbols can be calculated for any electron configuration, not just the ground-state configuration listed in tables, although not all the energy levels are observed in practice. It is through the analysis of atomic spectra that the ground-state electron configurations of the elements were experimentally determined. See also 

Born-Oppenheimer approximation



Atomic electron configuration table



Electron configurations of the elements (data page)



Periodic table (electron configurations)



Atomic orbital



Energy level



Term symbol



Molecular term symbol



HOMO/LUMO



Periodic Table Group



d electron count



Extension of the periodic table beyond the seventh period Discusses the limits of the periodic table

References 

Jolly, William L. (1991). Modern Inorganic Chemistry (2nd ed.). New York: McGrawHill. pp. 1–23. ISBN 0-07-112651-1.

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Chapter 3: Beryllium Beryllium is the chemical element with the symbol Be and atomic number 4. Because any beryllium synthesized in stars is short-lived, it is a relatively rare element in both the universe and in the crust of the Earth. It is a divalent element which occurs naturally only in combination with other elements in minerals. Notable gemstones which contain beryllium include beryl (aquamarine, emerald) and chrysoberyl. As a free element it is a steel-gray, strong, lightweight and brittle alkaline earth metal. Beryllium increases hardness and resistance to corrosion when alloyed with aluminium, cobalt, copper (notably beryllium copper), iron and nickel. In structural applications, high flexural rigidity, thermal stability, thermal conductivity and low density (1.85 times that of water) make beryllium a quality aerospace material for high-speed aircraft, missiles, spacecraft, and communication satellites. Because of its low density and atomic mass, beryllium is relatively transparent to X-rays and other forms of ionizing radiation; therefore, it is the most common window material for X-ray equipment and in particle physics experiments. The high thermal conductivities of beryllium and beryllium oxide have led to their use in heat transport and heat sinking applications. The commercial use of beryllium presents technical challenges because of the toxicity of inhaled beryllium-containing dusts. Beryllium is corrosive to tissue, and can cause a chronic life-threatening allergic disease called berylliosis in some people.

Characteristics

Physical properties Beryllium is a steel gray and hard metal that is brittle at room temperature and has a closepacked hexagonal crystal structure. It has exceptional flexural rigidity (Young's modulus 287 GPa) and a reasonably high melting point. The modulus of elasticity of beryllium is approximately 50% greater than that of steel. The combination of this modulus and a relatively low density results in an unusually fast sound conduction speed in beryllium – about 12.9 km/s at ambient conditions. Other significant properties are high specific heat −1

−1

(1925 J·kg ·K ) and thermal conductivity (216 W·m ·K ), which make beryllium the metal with the best heat dissipation characteristics per unit weight. In combination with the relatively low coefficient of linear thermal expansion (11.4×10

−6

−1

K ), these characteristics

result in a unique stability under conditions of thermal loading.

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Nuclear properties Natural beryllium, save for slight contamination by cosmogenic radioisotopes, is essentially beryllium-9, which has a nuclear spin of 3/2-. Beryllium has a large scattering cross section for high-energy neutrons, about 6 barns for energies above approximately 10 KeV. Therefore, it works as a neutron reflector and neutron moderator, effectively slowing the neutrons to the thermal energy range of below 0.03 eV, where the total cross section is at least an order of magnitude lower â&#x20AC;&#x201C; exact value strongly depends on the purity and size of the crystallites in the material. Beryllium also releases neutrons under bombardment by gamma rays. Thus, natural beryllium bombarded either by alphas or gammas from a suitable radioisotope is a key component of most radioisotope-powered nuclear reaction neutron sources for the laboratory production of free neutrons. As a metal, beryllium is transparent to most wavelengths of X-rays and gamma rays, making it useful for the output windows of X-ray tubes and other such apparatus.

Isotopes and nucleosynthesis Main articles: Isotopes of beryllium and beryllium-10 Both stable and unstable isotopes of beryllium are created in stars, but these do not last long. It is believed that most of the stable beryllium in the universe was originally created in the interstellar medium when cosmic rays induced fission in heavier elements found in 9

interstellar gas and dust. Primordial beryllium contains only one stable isotope, Be, and therefore beryllium is a monoisotopic element.

Plot showing variations in solar activity, including variation in sunspot number (red) and

concentration (blue). Note that the beryllium scale is inverted, so increases on this scale indicate lower

Be levels

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Radioactive cosmogenic

Be is produced in the atmosphere of the Earth by the cosmic ray

spallation of oxygen. Be accumulates at the soil surface, where its relatively long half-life (1.36 million years) permits a long residence time before decaying to boron-10. Thus,

and its daughter products are used to examine natural soil erosion, soil formation and the development of lateritic soils, and as a proxy for measurement of the variations in solar activity and the age of ice cores. The production of

Be is inversely proportional to solar

activity, because increased solar wind during periods of high solar activity decreases the flux of galactic cosmic rays that reach the Earth. Nuclear explosions also form reaction of fast neutrons with

Be by the

C in the carbon dioxide in air. This is one of the indicators of 7

past activity at nuclear weapon test sites. The isotope Be (half-life 53 days) is also cosmogenic, and shows an atmospheric abundance linked to sunspots, much like 8

Be has a very short half-life of about 7×10

−17

Be.

s that contributes to its significant cosmological

role, as elements heavier than beryllium could not have been produced by nuclear fusion in the Big Bang. This is due to the lack of sufficient time during the Big Bang's nucleosynthesis 4

phase to produce carbon by the fusion of He nuclei and the very low concentrations of available beryllium-8. The British astronomer Sir Fred Hoyle first showed that the energy 8

levels of Be and

C allow carbon production by the so-called triple-alpha process in helium-

fueled stars where more nucleosynthesis time is available. This process allows carbon to be produced in stars, but not in the Big Bang. Star-created carbon (the basis of carbon-based life) is thus a component in the elements in the gas and dust ejected by AGB stars and supernovae (see also Big Bang nucleosynthesis), as well as the creation of all other elements with atomic numbers larger than that of carbon. The innermost electrons of beryllium may contribute to chemical bonding. Therefore, when 7

Be decays by electron capture, it does so by taking electrons from atomic orbitals that may

participate in bonding. This makes its decay rate dependent to a measurable degree upon its electron configuration – a rare occurrence in nuclear decay. The shortest-lived known isotope of beryllium is −21

emission. It has a half-life of 2.7 × 10 −21

5.0 × 10

s. The exotic isotopes

Be which decays through neutron

s. Be is also very short-lived with a half-life of

Be and

Be are known to exhibit a nuclear halo. This

phenomenon can be understood as the nuclei of

Be and

Be have, respectively, 1 and 4

neutrons orbiting substantially outside the classical Fermi 'waterdrop' model of the nucleus.

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Occurrence

Beryllium ore

Emerald is a naturally occurring compound of beryllium. The Sun has a concentration of 0.1 parts per billion (ppb) of beryllium. Beryllium has a concentration of 2 to 6 parts per million (ppm) in the Earth's crust. It is most concentrated in 9

the soils, 6 ppm, and is found in 0.2 parts per trillion (ppt) of sea water. Trace amounts of Be are found in the Earth's atmosphere. In sea water, beryllium is exceedingly rare, comprising only 0.0006 ppb by weight. In stream water, however, beryllium is more abundant with 0.1 ppb by weight. Beryllium is found in over 100 minerals, but most are uncommon to rare. The more common beryllium containing minerals include: bertrandite (Be4Si2O7(OH)2), beryl (Al2Be3Si6O18), chrysoberyl (Al2BeO4) and phenakite (Be2SiO4). Precious forms of beryl are aquamarine, red beryl and emerald. The green color in gem-quality forms of beryl comes from varying amounts of chromium (about 2% for emerald). The two main ores of beryllium, beryl and bertrandite, are found in Argentina, Brazil, India, Madagascar, Russia and the United States. Total world reserves of beryllium ore are greater than 400,000 tonnes.

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Production The extraction of beryllium from its compounds is a difficult process due to its high affinity for oxygen at elevated temperatures, and its ability to reduce water when its oxide film is removed. The United States, China and Kazakhstan are the only three countries involved in the industrial scale extraction of beryllium. Beryllium is most commonly extracted from beryl, which is either sintered using an extraction agent or melted into a soluble mixture. The sintering process involves mixing beryl with sodium fluorosilicate and soda at 770 °C to form sodium fluoroberyllate, aluminium oxide and silicon dioxide.Beryllium hydroxide is precipitated from a solution of sodium fluoroberyllate and sodium hydroxide in water. Extraction of beryllium using the melt method involves grinding beryl into a powder and heating it to 1,650 °C. The melt is quickly cooled with water and then reheated 250 to 300 °C in concentrated sulfuric acid, mostly yielding beryllium sulfate and aluminium sulfate. Aqueous ammonia is then used to remove the aluminium and sulfur, leaving beryllium hydroxide. Beryllium hydroxide created using either the sinter or melt method is then converted into beryllium fluoride or beryllium chloride. To form the fluoride, aqueous ammonium hydrogen fluoride

added

beryllium

hydroxide

yield

precipitate

ammonium

tetrafluoroberyllate, which is heated to 1,000°C to form beryllium fluoride. Heating the fluoride to 900 °C with magnesium forms finely divided beryllium and additional heating to 1,300 °C creates the compact metal. Heating beryllium hydroxide forms the oxide which becomes beryllium chloride when mixed with carbon and chloride. Electrolysis of molten beryllium chloride is then used to obtain the metal.

Chemical properties See also category: Beryllium compounds Beryllium's chemical behavior is largely a result of its small atomic and ionic radii. It thus has very high ionization potentials and strong polarization while bonded to other atoms, which is why all of its compounds are covalent. It is more chemically similar to aluminium than its close neighbors in the periodic table due to having a similar charge-to-radius ratio. An oxide layer forms around beryllium that prevents further reactions with air unless heated above 1000 °C. Once ignited, beryllium burns brilliantly forming a mixture of beryllium oxide and beryllium nitride. Beryllium dissolves readily in non-oxidizing acids, such as HCl and diluted H2SO4, but not in nitric acid or water as this forms the oxide. This behavior is similar to that of aluminium metal. Beryllium also dissolves in alkali solutions.

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Beryllium

hydrolysis

function

Water molecules attached to Be are omitted 2

The beryllium atom has the electronic configuration [He] 2s . The two valence electrons give beryllium a +2 oxidation state and thus the ability to form two covalent bonds; the only evidence of lower valence of beryllium is in the solubility of the metal in BeCl2. Due to the octet rule, atoms tend to seek a valence of 8 in order to resemble a noble gas. Beryllium tries to achieve a coordination number of 4 because its two covalent bonds fill half of this octet. A coordination of 4 allows beryllium compounds, such as the fluoride or chloride, to form polymers. This characteristic is employed in analytical techniques using EDTA as a ligand. EDTA preferentially forms octahedral complexes â&#x20AC;&#x201C; thus absorbing other cations such as Al

which

might interfere â&#x20AC;&#x201C; for example, in the solvent extraction of a complex formed between Be

and acetylacetone. Beryllium(II) readily forms complexes with strong donating ligands such as phosphine oxides and arsine oxides. There have been extensive studies of these complexes which show the stability of the O-Be bond.

[citation needed]

Solutions of beryllium salts, e.g. beryllium sulfate and beryllium nitrate, are acidic because of hydrolysis of the [Be(H2O)4] 2+

[Be(H2O)4]

+ H2O

ion. +

[Be(H2O)3(OH)] + H3O

Other products of hydrolysis include the trimeric ion [Be3(OH)3(H2O)6] . Beryllium hydroxide, Be(OH)2, is insoluble even in acidic solutions with pH less than 6, that is at biological pH. It is amphoteric and dissolves in strongly alkaline solutions. Beryllium forms binary compounds with many non-metals. Anhydrous halides are known for F, Cl, Br and I. BeF2 has a silica-like structure with corner-shared BeF4 tetrahedra. BeCl2 and BeBr2 have chain structures with edge-shared tetrahedra. All beryllium halides have a linear

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monomeric molecular structure in the gas phase. Beryllium difluoride, BeF2, is different than the other difluorides. In general, beryllium has a tendency to bond covalently, much more so than the other alkaline earths and its fluoride is partially covalent (although still more ionic than its other halides). BeF 2 has many similarities to SiO2 (quartz) a mostly covalently bonded network solid. BeF2 has tetrahedrally coordinated metal and forms glasses (is difficult to crystallize). When crystalline, beryllium fluoride has the same room temperature crystal structure as quartz and shares many higher temperature structures also. Beryllium difluoride is very soluble in water, unlike the other alkaline earths. (Although they are strongly ionic, they do not dissolve because of the especially strong lattice energy of the fluorite structure.) However, BeF2 has much lower electrical conductivity when in solution or when molten than would be expected if it were fully ionic.

Order and disorder in difluorides

The strong and stable ionic fluorite

Disordered structure of beryllium

structure

glass (sketch, two dimensions)

adopted

calcium

difluoride and many other difluorides Beryllium oxide, BeO, is a white refractory solid, which has the wurtzite crystal structure and a thermal conductivity as high as in some metals. BeO is amphoteric. Salts of beryllium can be produced by treating Be(OH)2 with acid. Beryllium sulfide, selenide and telluride are known, all having the zincblende structure. Beryllium nitride, Be3N2 is a high-melting-point compound which is readily hydrolyzed. Beryllium azide, BeN6 is known and beryllium phosphide, Be3P2 has a similar structure to

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Be3N2. Basic beryllium nitrate and basic beryllium acetate have similar tetrahedral structures with four beryllium atoms coordinated to a central oxide ion. A number of beryllium borides are known, such as Be5B, Be4B, Be2B, BeB2, BeB6 and BeB12. Beryllium carbide, Be2C, is a refractory brick-red compound that reacts with water to give methane. No beryllium silicide has been identified.

History The mineral beryl, which contains beryllium, has been used at least since the Ptolemaic dynasty of Egypt. In the first century CE, Roman naturalist Pliny the Elder mentioned in his encyclopedia Natural History that beryl and emerald ("smaragdus") were similar. The Papyrus Graecus Holmiensis, written in the third or fourth century CE, contains notes on how to prepare artificial emerald and beryl.

Louis-Nicolas Vauquelin discovered beryllium Early analyses of emeralds and beryls by Martin Heinrich Klaproth, Torbern Olof Bergman, Franz Karl Achard, and Johann Jakob Bindheim always yielded similar elements, leading to the fallacious conclusion that both substances are aluminium silicates. Mineralogist René Just Haüy discovered that both crystals are geometrically identical, and he asked chemist Louis-Nicolas Vauquelin for a chemical analysis. In a 1798 paper read before the Institut de France, Vauquelin reported that he found a new "earth" by dissolving aluminium hydroxide from emerald and beryl in an additional alkali. The editors of the journal Annales de Chimie et de Physique named the new earth "glucine" for the sweet taste of some of its compounds. Klaproth preferred the name "beryllina" due to fact that yttria also formed sweet salts. The name "beryllium" was first used by Wöhler in 1828.

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Friedrich Wöhler was one of the men who independently isolated beryllium Friedrich Wöhler and Antoine Bussy independently isolated beryllium in 1828 by the chemical reaction of metallic potassium with beryllium chloride, as follows: BeCl2 + 2 K → 2 KCl + Be Using an alcohol lamp, Wöhler heated alternating layers of beryllium chloride and potassium in a wired-shut platinum crucible. The above reaction immediately took place and caused the crucible to become white hot. Upon cooling and washing the resulting gray-black powder he saw that it was made of fine particles with a dark metallic luster. The highly reactive potassium had been produced by the electrolysis of its compounds, a process discovered 21 years before. The chemical method using potassium yielded only small grains of beryllium from which no ingot of metal could be cast or hammered. The direct electrolysis of a molten mixture of beryllium fluoride and sodium fluoride by Paul Lebeau in 1898 resulted in the first pure (99.5 to 99.8%) samples of beryllium. The first commercially-successful process for producing beryllium was developed in 1932 by Alfred Stock and Hans Goldschmidt. Their process involves the electrolysation of a mixture of beryllium fluorides and barium, which causes molten beryllium to collect on a water-cooled iron cathode. A sample of beryllium was bombarded with alpha rays from the decay of radium in a 1932 experiment by James Chadwick that uncovered the existence of the neutron. This same method is used in one class of radioisotope-based laboratory neutron sources that produce 30 neutrons for every million α particles. Beryllium production saw a rapid increase during World War II, due to the rising demand for hard beryllium-copper alloys and phosphors for fluorescent lights. Most early fluorescent lamps used zinc orthosilicate with varying content of beryllium to emit greenish light. Small

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additions of magnesium tungstate improved the blue part of the spectrum to yield an acceptable

white

light.

Halophosphate-based

phosphors

replaced

beryllium-based

phosphors after beryllium was found to be toxic. Electrolysis of a mixture of beryllium fluoride and sodium fluoride was used to isolate beryllium during the 19th century. The metal's high melting point makes this process more energy-consuming than corresponding processes used for the alkali metals. Early in the 20th century, the production of beryllium by the thermal decomposition of beryllium iodide was investigated following the success of a similar process for the production of zirconium, but this process proved to be uneconomical for volume production. Pure beryllium metal did not become readily available until 1957, even though it had been used as an alloying metal to harden and toughen copper much earlier. Beryllium could be produced by reducing beryllium compounds such as beryllium chloride with metallic potassium or sodium. Currently most beryllium is produced by reducing beryllium fluoride with purified magnesium. The price on the American market for vacuum-cast beryllium ingots was about $338 per pound ($745 per kilogram) in 2001. Between 1998 and 2008, the world's production of beryllium had decreased from 343 to about 200 tonnes, of which 176 tonnes (88%) came from the United States.

Etymology Early usage of the word beryllium can be traced to many languages, including Latin Beryllus; French Béry; Greek βήρσλλος, bērullos, beryl; Prakrit veruliya (वॆरुलिय); Pāli veḷuriya (वेिुरिय), veḷiru (भेलिरु) or viḷar (लभिि्) – "to become pale", in reference to the pale semiprecious gemstone beryl. The original source is probably the Sanskrit word वैडूयय vaidurya-, which is of Dravidian origin and could be derived from the name of the modern city of Belur. For about 160 years, beryllium was also known as glucinum or glucinium (with the accompanying chemical symbol "Gl",), the name coming from the Greek word for sweet: γλσκσς, due to the sweet taste of beryllium salts.

Applications It is estimated that most beryllium is used for military applications, so information is not readily available.

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Radiation windows

Beryllium target which "converts" a proton beam into a neutron beam

A square beryllium foil mounted in a steel case to be used as a window between a vacuum chamber and an X-ray microscope. Beryllium is highly transparent to X-rays owing to its low atomic number. Because of its low atomic number and very low absorption for X-rays, the oldest and still one of the most important applications of beryllium is in radiation windows for X-ray tubes. Extreme demands are placed on purity and cleanliness of beryllium to avoid artifacts in the X-ray images. Thin beryllium foils are used as radiation windows for X-ray detectors, and the extremely low absorption minimizes the heating effects caused by high intensity, low energy X-rays typical of synchrotron radiation. Vacuum-tight windows and beam-tubes for radiation experiments on synchrotrons are manufactured exclusively from beryllium. In scientific setups for various X-ray emission studies (e.g., energy-dispersive X-ray spectroscopy) the sample holder is usually made of beryllium because its emitted X-rays have much lower energies (~100 eV) than X-rays from most studied materials. Low atomic number also makes beryllium relatively transparent to energetic particles. Therefore it is used to build the beam pipe around the collision region in particle physics setups, such as all four main detector experiments at the Large Hadron Collider (ALICE, ATLAS, CMS, LHCb), the Tevatron and the SLAC. The low density of beryllium allows collision products to reach the surrounding detectors without significant interaction, its

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stiffness allows a powerful vacuum to be produced within the pipe to minimize interaction with gases, its thermal stability allows it to function correctly at temperatures of only a few degrees above absolute zero, and its diamagnetic nature keeps it from interfering with the complex multipole magnet systems used to steer and focus the particle beams.

Mechanical applications Because of its stiffness, light weight and dimensional stability over a wide temperature range, beryllium metal is used for lightweight structural components in the defense and aerospace industries in high-speed aircraft, guided missiles, spacecraft, and satellites. Several liquidfuel rockets have used rocket nozzles made of pure beryllium. Beryllium powder was itself studied as a rocket fuel, but this use has never materialized. A small number of bicycle frames were built with beryllium, at "astonishing" prices. From 1998 to 2000, the McLaren Formula One team used Mercedes-Benz engines with beryllium-aluminium-alloy pistons. The use of beryllium engine components was banned following a protest by Scuderia Ferrari. Mixing about 2.0% beryllium into copper forms an alloy called beryllium copper that is six times stronger than copper alone. Beryllium alloys are used in many applications because of their combination of elasticity, high electrical conductivity and thermal conductivity, high strength and hardness, nonmagnetic properties, as well as good corrosion and fatigue resistance. These applications include non-sparking tools that are used near flammable gases (beryllium nickel), in springs and membranes (beryllium nickel and beryllium iron) used in surgical instruments and high temperature devices. As little as 50 parts per million of beryllium alloyed with liquid magnesium leads to a significant increase in oxidation resistance and decrease in flammability.

Beryllium Copper Adjustable Wrench The high elastic stiffness of beryllium has led to its extensive use in precision instrumentation, e.g. in inertial guidance systems and in the support mechanisms for optical systems. Beryllium-copper alloys were also applied as a hardening agent in "Jason pistols", which were used to strip the paint from the hulls of ships. An earlier major application of beryllium was in brakes for military airplanes because of its hardness, high melting point, and exceptional ability to dissipate heat. Environmental considerations have led to substitution by other materials.

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To reduce costs, beryllium can be alloyed with significant amounts of aluminium, resulting in the AlBeMet alloy (a trade name). This blend is cheaper than pure beryllium, while still retaining many desirable properties.

Mirrors Beryllium mirrors are of particular interest. Large-area mirrors, frequently with a honeycomb support structure, are used, for example, in meteorological satellites where low weight and long-term dimensional stability are critical. Smaller beryllium mirrors are used in optical guidance systems and in fire-control systems, e.g. in the German-made Leopard 1 and Leopard 2 main battle tanks. In these systems, very rapid movement of the mirror is required which again dictates low mass and high rigidity. Usually the beryllium mirror is coated with hard electroless nickel plating which can be more easily polished to a finer optical finish than beryllium. In some applications, though, the beryllium blank is polished without any coating. This is particularly applicable to cryogenic operation where thermal expansion mismatch can cause the coating to buckle. The James Webb Space Telescope will have 18 hexagonal beryllium sections for its mirrors. Because JWST will face a temperature of 33 K, the mirror is made of beryllium, capable of handling extreme cold better than glass. Beryllium contracts and deforms less than glass â&#x20AC;&#x201C; and remains more uniform â&#x20AC;&#x201C; in such temperatures. For the same reason, the optics of the Spitzer Space Telescope are entirely built of beryllium metal.

Magnetic applications Beryllium is non-magnetic. Therefore, tools fabricated out of beryllium are used by naval or military explosive ordnance disposal teams for work on or near naval mines, since these mines commonly have magnetic fuzes. They are also found in maintenance and construction materials near magnetic resonance imaging (MRI) machines because of the high magnetic fields generated by them. In the fields of radio communications and powerful (usually military) radars, hand tools made of beryllium are used to tune the highly magnetic klystrons, magnetrons, traveling wave tubes, etc., that are used for generating high levels of microwave power in the transmitters.

Nuclear applications Thin plates or foils of beryllium are sometimes used in nuclear weapon designs as the very outer layer of the plutonium pits in the primary stages of thermonuclear bombs, placed to surround the fissile material. These layers of beryllium are good "pushers" for the implosion of the plutonium-239, and they are also good neutron reflectors, just as they are in beryllium-

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moderated nuclear reactors. Beryllium is also commonly used in some neutron sources in laboratory devices in which relatively few neutrons are needed (rather than having to use a nuclear reactor, or a particle accelerator-powered neutron generator). For this purpose, a target of beryllium-9 is bombarded with energetic alpha particles from a radioisotope such as polonium-210, radium226, plutonium-239, or americium-241. In the nuclear reaction that occurs, a beryllium nucleus is transmuted into carbon-12, and one free neutron is emitted, traveling in about the same direction as the alpha particle was heading. Such alpha decay driven beryllium neutron sources, named "urchin" neutron initiators, were used some in early atomic bombs. Neutron sources in which beryllium is bombarded with gamma rays from a gamma decay radioisotope, are also used to produce laboratory neutrons.

Two CANDU fuel bundles: Each about 50 cm in length and 10 cm in diameter. Notice the small appendages on the fuel clad surfaces Beryllium is also used in fuel fabrication for CANDU reactors. The fuel elements have small appendages that are resistance brazed to the fuel cladding using an induction brazing process with Be as the braze filler material. Bearing pads are brazed on to prevent fuel bundle to pressure tube contact, and inter-element spacer pads are brazed on to prevent element to element contact. Beryllium is also used at the Joint European Torus nuclear-fusion research laboratory, and it will be used in the more advanced ITER to condition the components which face the plasma. Beryllium has also been proposed as a cladding material for nuclear fuel rods, because of its good combination of mechanical, chemical, and nuclear properties.Beryllium fluoride is one of the constituent salts of the eutectic salt mixture FLiBe, which is used as a solvent, moderator and coolant in many hypothetical molten salt reactor designs.

Acoustics Low weight and high rigidity of beryllium make it useful as a material for high-frequency speaker drivers. Because beryllium is expensive (many times more than titanium), hard to

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shape due to its brittleness, and toxic if mishandled, beryllium tweeters are limited to highend home,pro audio, and public address applications. Due to the high performance of beryllium in acoustics, for marketing purposes some products are claimed to be made of the material when they are not.

Electronic Beryllium is a p-type dopant in III-V compound semiconductors. It is widely used in materials such as GaAs, AlGaAs, InGaAs and InAlAs grown by molecular beam epitaxy (MBE). Crossrolled beryllium sheet is an excellent structural support for printed circuit boards in surfacemount technology. In critical electronic applications, beryllium is both a structural support and heat sink. The application also requires a coefficient of thermal expansion that is well matched to the alumina and polyimide-glass substrates. The beryllium-beryllium oxide composite "E-Materials" have been specially designed for these electronic applications and have the additional advantage that the thermal expansion coefficient can be tailored to match diverse substrate materials. Beryllium oxide is useful for many applications that require the combined properties of an electrical insulator and an excellent heat conductor, with high strength and hardness, and a very high melting point. Beryllium oxide is frequently used as an insulator base plate in highpower transistors in radio frequency transmitters for telecommunications. Beryllium oxide is also being studied for use in increasing the thermal conductivity of uranium dioxide nuclear fuel pellets. Beryllium compounds were used in fluorescent lighting tubes, but this use was discontinued because of the disease berylliosis which developed in the workers who were making the tubes.

Precautions Approximately 35 micrograms of beryllium is found in the human body, but this amount is not considered harmful. Beryllium is chemically similar to magnesium and therefore can displace it from enzymes, which causes them to malfunction. Chronic berylliosis is a pulmonary and systemic granulomatous disease caused by inhalation of dust or fumes contaminated with beryllium; either large amounts over a short time or small amounts over a long time can lead to this ailment. Symptoms of the disease can take up to five years to develop; about a third of patients with it die and the survivors are left disabled. The International Agency for Research on Cancer (IARC) lists beryllium and beryllium compounds as Category 1 carcinogens. Acute beryllium disease in the form of chemical pneumonitis was first reported in Europe in 1933 and in the United States in 1943. A survey found that about 5% of workers in plants

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manufacturing fluorescent lamps in 1949 in the United States had beryllium-related lung diseases. Chronic berylliosis resembles sarcoidosis in many respects, and the differential diagnosis is often difficult. It killed some early workers in nuclear weapons design, such as Herbert L. Anderson. Beryllium may be found in coal slag. When the slag is formulated into an abrasive agent for blasting paint and rust from hard surfaces, the beryllium can become airborne and become a source of exposure. Early researchers tasted beryllium and its various compounds for sweetness in order to verify its presence. Modern diagnostic equipment no longer necessitates this highly risky procedure and no attempt should be made to ingest this highly toxic substance. Beryllium and its compounds should be handled with great care and special precautions must be taken when carrying out any activity which could result in the release of beryllium dust (lung cancer is a possible result of prolonged exposure to beryllium laden dust). Although the use of beryllium compounds in fluorescent lighting tubes was discontinued in 1949, potential for exposure to beryllium exists in the nuclear and aerospace industries and in the refining of beryllium metal and melting of beryllium-containing alloys, the manufacturing of electronic devices, and the handling of other beryllium-containing material. A successful test for beryllium in air and on surfaces has been recently developed and published as an international voluntary consensus standard ASTM D7202. The procedure uses dilute ammonium bifluoride for dissolution and fluorescence detection with beryllium bound to sulfonated hydroxybenzoquinoline, allowing up to 100 times more sensitive detection than the recommended limit for beryllium concentration in the workplace. Fluorescence increases with increasing beryllium concentration. The new procedure has been successfully tested on a variety of surfaces and is effective for the dissolution and ultratrace detection of refractory beryllium oxide and siliceous beryllium (ASTM D7458).

See also ď&#x20AC;

Sucker Bait, a novella by Isaac Asimov in which the health hazard of beryllium dust is an important plot point

References 1. Emsley, John (2001). Nature's Building Blocks: An Aâ&#x20AC;&#x201C;Z Guide to the Elements. Oxford, England, UK: Oxford University Press. ISBN 0-19-850340-7.

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Chapter 4: Magnesium Magnesium is a chemical element with the symbol Mg and atomic number 12. Its common oxidation number is +2. It is an alkaline earth metal and the eighth-most-abundant element in the Earth's crust and ninth in the known universe as a whole. Magnesium is the fourth-mostcommon element in the Earth as a whole (behind iron, oxygen and silicon), making up 13% of the planet's mass and a large fraction of the planet's mantle. The relative abundance of magnesium is related to the fact that it easily builds up in supernova stars from a sequential addition of three helium nuclei to carbon (which in turn is made from three helium nuclei).[citation needed] Due to magnesium ion's high solubility in water, it is the third-most-abundant element dissolved in seawater. Magnesium is produced in stars larger than 3 solar masses by fusing helium and neon in the alpha process at temperatures above 600 megakelvins.[citation needed] The free element (metal) is not found naturally on Earth, as it is highly reactive (though once produced, it is coated in a thin layer of oxide (see passivation), which partly masks this reactivity). The free metal burns with a characteristic brilliant-white light, making it a useful ingredient in flares. The metal is now obtained mainly by electrolysis of magnesium salts obtained from brine. In commerce, the chief use for the metal is as an alloying agent to make aluminium-magnesium

alloys,

sometimes

called

magnalium

magnelium.

Since

magnesium is less dense than aluminium, these alloys are prized for their relative lightness and strength. In human biology, magnesium is the eleventh-most-abundant element by mass in the human body. Its ions are essential to all living cells, where they play a major role in manipulating important biological polyphosphate compounds like ATP, DNA, and RNA. Hundreds of enzymes, thus, require magnesium ions to function. Magnesium compounds are used medicinally as common laxatives, antacids (e.g., milk of magnesia), and in a number of situations where stabilization of abnormal nerve excitation and blood vessel spasm is required (e.g., to treat eclampsia). Magnesium ions are sour to the taste, and in low concentrations they help to impart a natural tartness to fresh mineral waters. In vegetation, magnesium is the metallic ion at the center of chlorophyll, and is, thus, a common additive to fertilizers. Characteristics Physical properties Elemental magnesium is a rather strong, silvery-white, light-weight metal (two-thirds the density of aluminium). It tarnishes slightly when exposed to air, although, unlike the alkali metals, an oxygen-free environment is unnecessary for storage because magnesium is protected by a thin layer of oxide that is fairly impermeable and difficult to remove. Like its lower periodic table group neighbor calcium, magnesium reacts with water at room

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temperature, though it reacts much more slowly than calcium. When submerged in water, hydrogen bubbles almost unnoticeably begin to form on the surface of the metalâ&#x20AC;&#x201D;though, if powdered, it reacts much more rapidly. The reaction occurs faster with higher temperatures (see precautions). Magnesium's ability to react with water can be harnessed to produce energy and run a magnesium-based engine. Magnesium also reacts exothermically with most acids, such as hydrochloric acid (HCl). As with aluminium, zinc, and many other metals, the reaction with HCl produces the chloride of the metal and releases hydrogen gas. Chemical properties Magnesium is a highly flammable metal, but, while it is easy to ignite when powdered or shaved into thin strips, it is difficult to ignite in mass or bulk. Once ignited, it is difficult to extinguish, being able to burn in nitrogen (forming magnesium nitride), carbon dioxide (forming magnesium oxide, and carbon) and water (forming magnesium oxide and hydrogen). This property was used in incendiary weapons used in the firebombing of cities in World War II, the only practical civil defense being to smother a burning flare under dry sand to exclude the atmosphere. On burning in air, magnesium produces a brilliant-white light that includes strong ultraviolet. Thus, magnesium powder (flash powder) was used as a source of illumination in the early days of photography. Later, magnesium ribbon was used in electrically ignited flashbulbs. Magnesium powder is used in the manufacture of fireworks and marine flares where a brilliant white light is required. Flame temperatures of magnesium and magnesium alloys can reach 3,100 Â°C (3,370 K; 5,610 Â°F), although flame height above the burning metal is usually less than 300 mm (12 in). Magnesium may be used as an ignition source for thermite, a mixture of aluminium and iron oxide powder that is otherwise difficult to ignite. Those properties are due to magnesium's high specific heat, the fourthhighest specific heat among the metals. Magnesium compounds are typically white crystals. Most are soluble in water, providing the 2+

sour-tasting magnesium ion Mg . Small amounts of dissolved magnesium ion contribute to the tartness and taste of natural waters. Magnesium ion in large amounts is an ionic laxative, and magnesium sulfate (common name: Epsom salt) is sometimes used for this purpose. So-called "milk of magnesia" is a water suspension of one of the few insoluble magnesium compounds, magnesium hydroxide. The undissolved particles give rise to its appearance and name. Milk of magnesia is a mild base commonly used as an antacid, which has some laxative side-effect. Occurrence Magnesium is the eighth-most-abundant element in the Earth's crust by mass and tied in seventh place with iron in terms of molarity. It is found in large deposits of magnesite, dolomite, and other minerals, and in mineral waters, where magnesium ion is soluble.

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Although magnesium is found in over 60 minerals, only dolomite, magnesite, brucite, carnallite, talc, and olivine are of commercial importance. Forms Alloy As of 2013, magnesium alloy consumption is less than a million tons per year, compared with 50 million tons of aluminum alloys. Its use has been historically limited by its tendency to corrode, high-temperature creep, and flammability. Research and development eliminated magnesium's tendency toward high-temperature creep by inclusion of scandium and gadolinium. Flammability was greatly reduced by introducing a small amount of calcium into the mix. The presence of iron, nickel, copper, and cobalt strongly activates corrosion. This is due to their low solid solubility limits (above a very small percentage, they precipitate out as intermetallic compounds) and because they behave as active cathodic sites that reduce water and cause the loss of magnesium. Reducing the quantity of these metals improves corrosion resistance. Sufficient manganese overcomes the corrosive effects of iron. This requires precise control over composition, increasing costs. Adding a cathodic poison captures atomic hydrogen within the structure of a metal. This prevents the formation of free hydrogen gas, which is required for corrosive chemical processes. The addition of about one-third of a percent of arsenic reduces its corrosion rate in a salt solution by a factor of nearly ten. Compounds Magnesium forms a variety of industrially and biologically important compounds, including magnesium oxide, various salts, and others. Isotopes Magnesium has three stable isotopes:

Mg,

Mg and

Mg. All are present in significant

amounts (see table of isotopes above). About 79% of Mg is

Mg. The isotope

Mg is

radioactive and in the 1950s to 1970s was made commercially by several nuclear power plants for use in scientific experiments. This isotope has a relatively short half-life (21 hours) and so its use was limited by shipping times. 26

Mg has found application in isotopic geology, similar to that of aluminium.

radiogenic daughter product of

Mg is a

Al, which has a half-life of 717,000 years. Large

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enrichments of stable

Mg have been observed in the Ca-Al-rich inclusions of some

carbonaceous chondrite meteorites. The anomalous abundance of decay of its parent

Mg is attributed to the

Al in the inclusions. Therefore, the meteorite must have formed in the

solar nebula before the

Al had decayed. Hence, these fragments are among the oldest

objects in the solar system and have preserved information about its early history. It is conventional to plot 27

Mg/ Mg against an Al/Mg ratio. In an isochron dating plot, the

Al/Mg ratio plotted is Al/ Mg. The slope of the isochron has no age significance, but indicates the initial

Al/ Al ratio in the sample at the time when the systems were separated

from a common reservoir. Production 2011

Country

production

(tonnes)

China

661,000

U.S.

(Capacity.

Production

figures

withheld

avoid

disclosing

63,500

company proprietary data.) Russia

37,000

Israel

30,000

Kazakhstan

21,000

Brazil

16,000

Ukraine

2,000

Serbia

1,500

China is the dominant supplier of magnesium, with approximately 80% of the world market share. China is almost completely reliant on the silicothermic Pidgeon process (the reduction of the oxide at high temperatures with silicon) to obtain the metal. In the United States, magnesium is obtained principally by electrolysis of fused magnesium chloride

from

brines,

wells,

and

sea

water.

the

cathode,

the

ion is reduced by two electrons to magnesium metal: Mg2+ + 2 eâ&#x2C6;&#x2019; â&#x2020;&#x2019; Mg

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At the anode, each pair of Cl−ions is oxidized to chlorine gas, releasing two electrons to complete the circuit: −

2 Cl → Cl2 (g) + 2 e− A new process, solid oxide membrane technology, involves the electrolytic reduction of MgO. 2+

At the cathode, Mg ion is reduced by two electrons to magnesium metal. The electrolyte is Yttria-stabilized zirconia(YSZ). The anode is a liquid metal. At the YSZ/liquid metal anode O

2−

is oxidized. A layer of graphite borders the liquid metal anode, and at this interface carbon and oxygen react to form carbon monoxide. When silver is used as the liquid metal anode, there is no reductant carbon or hydrogen needed, and only oxygen gas is evolved at the anode. It has been reported that this method provides a 40% reduction in cost per pound over the electrolytic reduction method. This method is more environmentally sound than others because there is much less carbon dioxide emitted. The United States has traditionally been the major world supplier of this metal, supplying 45% of world production even as recently as 1995. Today, the US market share is at 7%, with a single domestic producer left, US Magnesium, a Renco Group company in Utah born from now-defunct Magcorp. History The name magnesium originates from the Greek word for a district in Thessaly called Magnesia. It is related to magnetite and manganese, which also originated from this area, and required differentiation as separate substances. See manganese for this history. In 1618, a farmer at Epsom in England attempted to give his cows water from a well there. The cows refused to drink because of the water's bitter taste, but the farmer noticed that the water seemed to heal scratches and rashes. The substance became known as Epsom salts and its fame spread. It was eventually recognized as hydrated magnesium sulfate, MgSO4·7 H2O. The metal itself was first produced by Sir Humphry Davy in England in 1808. He used electrolysis on a mixture of magnesia and mercuric oxide.Antoine Bussy prepared it in coherent form in 1831. Davy's first suggestion for a name was magnium, but the name magnesium is now used. Applications As metal Magnesium is the third-most-commonly-used structural metal, following iron and aluminium. It has been called the lightest useful metal by The Periodic Table of Videos.

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The main applications of magnesium are, in order: component of aluminium alloys, in diecasting (alloyed with zinc), to remove sulfur in the production of iron and steel, the production of titanium in the Kroll process. Magnesium, in its purest form, can be compared with aluminium, and is strong and light, so it is used in several high-volume part manufacturing applications, including automotive and truck components. Specialty, high-grade car wheels of magnesium alloy are called "mag wheels", although the term is often more broadly misapplied to include aluminium wheels. The high magnesium content used for the crankcase of the late-World War II Wright Duplex Cyclone eighteen-cylinder radial aviation engine was a serious problem for the earliest examples of the Boeing B-29 heavy bomber, as engine fires in flight could ignite the engine crankcases, literally "torching" the wing spar apart in under a minute, causing wing failure. An earlier Mercedes-Benz race car model, the Mercedes-Benz 300 SLR, had a body made from Elektron, a magnesium alloy; these cars ran (with successes) at Le Mans, the Mille Miglia, and other world-class race events in 1955 (though one was involved in the single worst accident in auto racing history, in terms of human casualties, at the Le Mans race.) Porsche's all-out quest to decrease the weight of their race cars led to the use of magnesium alloy frames in the famous 917/053 that won Le Mans in 1971. Volkswagen Group has used magnesium in its engine components for many years. For a long time, Porsche used magnesium alloy for its engine blocks due to the weight advantage. There is renewed interest in magnesium alloy engine blocks, as featured in the 2006 BMW 325i and 330i models. The BMW engine uses an aluminium alloy insert for the cylinder walls and cooling jackets surrounded by a high-temperature magnesium alloy AJ62A. The application of magnesium AE44 alloy in the 2006 Corvette Z06 engine cradle has advanced the technology of designing robust automotive parts in magnesium. Both these alloys are recent developments in high-temperature low creep magnesium alloys. Mitsubishi Motors also uses magnesium (branded magnesium alloy) for its paddle shifters. The general strategy for such alloys is to form intermetallic precipitates at the grain boundaries, for example by adding mischmetal or calcium. New alloy development and lower costs that make magnesium competitive with aluminium will increase the number of automotive applications.[citation needed] The second application field of magnesium is electronic devices. Because of low weight, and good mechanical and electrical properties, magnesium is widely used for manufacturing of

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mobile phones, laptop and tablet computers, cameras, and other electronic components. Historically, magnesium was one of the main aerospace construction metals and was used for German military aircraft as early as World War I and extensively for German aircraft in World War II. The Germans coined the name 'Elektron' for magnesium alloy. The term is still used today. The application of magnesium in the commercial aerospace industry was generally restricted to engine related components, due either to perceived hazards with magnesium parts in the event of fire or to corrosion. Currently, the use of magnesium alloys in aerospace is increasing, mostly driven by the increasing importance of fuel economy and the need to reduce weight. The development and testing of new magnesium alloys continues, notably Elektron 21, which has successfully undergone extensive aerospace testing for suitability in engine and internal and airframe components. The European Community runs three R&D magnesium projects in the Aerospace priority of Six Framework Program. Niche uses of the metal Magnesium, being readily available and relatively nontoxic, has a variety of uses: 

Magnesium is flammable, burning at a temperature of approximately 3,100 °C (3,370 K; 5,610 °F), and the autoignition temperature of magnesium ribbon is approximately 473 °C (746 K; 883 °F). It produces intense, bright, white light when it burns. Magnesium's high combustion temperature makes it a useful tool for starting emergency fires. Other uses include flash photography, flares, pyrotechnics, and fireworks sparklers. Magnesium is also often used to ignite thermite, or other materials that require a high ignition temperature. In the form of turnings or ribbons, to prepare Grignard reagents, which are useful in organic synthesis.



As an additive agent in conventional propellants and the production of nodular graphite in cast iron.



As a reducing agent for the production of uranium and other metals from their salts.



As a sacrificial (galvanic) anode to protect underground tanks, pipelines, buried structures, and water heaters.



Alloyed with zinc to produce the zinc sheet used in photoengraving plates in the printing industry, dry-cell battery walls, and roofing.

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

As a metal, this element's principal use is as an alloying additive to aluminium with these aluminium-magnesium alloys being used mainly for beverage cans, sports equipment such as golf clubs, fishing reels, and archery bows and arrows.

In compounds Magnesium compounds, primarily magnesium oxide (MgO), are used as a refractory material in furnace linings for producing iron, steel, nonferrous metals, glass, and cement. Magnesium oxide and other magnesium compounds are also used in the agricultural, chemical, and construction industries. Magnesium oxide from calcination is used as an electrical insulator in fire-resistant cables. Magnesium reacted with an alkyl halide gives a Grignard reagent, which is a very useful tool for preparing alcohols. Magnesium salts are frequently included in various foods, fertilizers (magnesium is a component of chlorophyll), and culture media. Magnesium sulfite is used in the manufacture of paper (sulfite process). Magnesium phosphate is used to fireproof wood used in construction. Magnesium hexafluorosilicate is used in mothproofing of textiles. In the form of turnings or ribbons, Mg is useful in purification of solvents, for example the preparation of super-dry ethanol. Biological Pharmaceutical preparations of magnesium are used to treat magnesium deficiency and hypomagnesemia, as well as eclampsia. Usually in lower dosages, magnesium is commonly included in dietary mineral preparations, including many multivitamin preparations. Sorted by type of magnesium salt, biological applications of magnesium include: 

Magnesium sulfate, as the heptahydrate called Epsom salts, is used as bath salts, as a laxative, and as a highly soluble fertilizer.



Magnesium hydroxide, suspended in water, is used in milk of magnesia antacids and laxatives.



Magnesium chloride, oxide, gluconate, malate, orotate, glycinate and citrate are all used as oral magnesium supplements. Some have claimed that oral magnesium supplements are therapeutic for Restless Leg Syndrome (RLS) in some individuals.

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

Magnesium borate, magnesium salicylate, and magnesium sulfate are used as antiseptics.



Magnesium bromide is used as a mild sedative (this action is due to the bromide, not the magnesium).



Magnesium stearate is a slightly flammable white powder with lubricating properties. In pharmaceutical technology, it is used in the manufacturing of numerous kinds of tablets to prevent the tablets from sticking to the equipment during the tablet compression process (i.e., when the tablet's substance is pressed into tablet form).



Magnesium carbonate powder is used by athletes such as gymnasts, weightlifters, and climbers to eliminate moisture and improving the grip on a gymnastic apparatus, lifting bar, and climbing rocks.

Biological roles Main article: Magnesium in biology Because of the important interaction between phosphate and magnesium ions, magnesium ions are essential to the basic nucleic acid chemistry of life, and thus are essential to all cells of all known living organisms. Over 300 enzymes require the presence of magnesium ions for their catalytic action, including all enzymes utilizing or synthesizing ATP, or those that use other nucleotides to synthesize DNA and RNA. ATP exists in cells normally as a chelate of ATP and a magnesium ion. Plants have an additional use for magnesium in that chlorophylls are magnesium-centered porphyrins. Magnesium deficiency in plants causes late-season yellowing between leaf veins, especially in older leaves, and can be corrected by applying Epsom salts (which is rapidly leached), or else crushed dolomitic limestone to the soil. Magnesium is a vital component of a healthy human diet. Human magnesium deficiency (including conditions that show few overt symptoms) is relatively rare although only 32% of people in the United States meet the RDA-DRI; low levels of magnesium in the body have been associated with the development of a number of human illnesses such as asthma, diabetes, and osteoporosis. Taken in the proper amount, magnesium plays a role in preventing both stroke and heart attack. The symptoms of people with fibromyalgia, migraines, and premenstrual syndrome are less severe, and magnesium can shorten the length of the migraine symptoms. Adult human bodies contain about 24 grams of magnesium, with 60% in the skeleton, 39% intracellular (20% in skeletal muscle), and 1% extracellular. Serum levels are typically 0.7–

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1.0 mmol/L or 1.8â&#x20AC;&#x201C;2.4 mEq/L. Serum magnesium levels may appear normal even in cases of underlying intracellular deficiency, although no known mechanism maintains a homeostatic level in the blood other than renal excretion of high blood levels. Intracellular magnesium is correlated with intracellular potassium. Magnesium is absorbed in the gastrointestinal tract, with more absorbed when status is lower. Magnesium competes with calcium in the human body, in this way it actually keeps calcium in check. However, this can cause a calcium deficiency if calcium levels are already low. Low and high protein intake inhibit magnesium absorption, and other factors such as phosphate, phytate, and fat affect absorption. Excess dietary magnesium is excreted in feces, urine, and sweat. Magnesium status may be assessed roughly through serum and erythrocyte Mg concentrations and urinary and fecal excretion, but intravenous magnesium loading tests are likely the most accurate and practical in most people. In these tests, magnesium is injected intravenously; a retention of 20% or more indicates deficiency. Other nutrient deficiencies are identified through biomarkers, but none are established for magnesium. The UK recommended daily values for magnesium is 300 mg for men and 270 mg for women. Spices, nuts, cereals, coffee, cocoa, tea, and vegetables are rich sources of magnesium. Green leafy vegetables such as spinach are also rich in magnesium as they contain chlorophyll. Observations of reduced dietary magnesium intake in modern Western countries compared to earlier generations may be related to food refining and modern fertilizers that contain no magnesium. Numerous pharmaceutical preparations of magnesium, as well as magnesium dietary supplements are available. Magnesium oxide, one of the most common forms in magnesium dietary supplements because it has high magnesium content per weight, has been reported the least bioavailable.Magnesium citrate has been reported as more bioavailable than oxide or amino-acid chelate (glycinate) forms. Excess magnesium in the blood is freely filtered at the kidneys, and for this reason it is difficult to overdose on magnesium from dietary sources alone. With supplements, overdose is possible, in particular in people with poor renal function; occasionally, with use of high cathartic doses of magnesium salts, severe hypermagnesemia has been reported to occur even without renal dysfunction. Alcoholism can produce a magnesium deficiency, which is reversible by oral or parenteral administration, depending on the degree of deficiency. Detection in biological fluids Magnesium concentrations in plasma or serum may be measured to monitor for efficacy and safety in those receiving the drug therapeutically, to confirm the diagnosis in potential poisoning victims or to assist in the forensic investigation in a case of fatal overdosage. The

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newborn children of mothers having received parenteral magnesium sulfate during labor may exhibit toxicity at serum magnesium levels that were considered appropriate for the mothers. Disease Results from a meta-analysis of randomized clinical trials demonstrated that a magnesium supplement can lower high blood pressure in a dose-dependent manner. Low serum magnesium levels are associated with metabolic syndrome, diabetes mellitus type 2 and hypertension. Low serum magnesium levels have been associated with a higher risk of developing metabolic syndrome. Magnesium therapy is recommended by the ACC/AHA/ESC 2006 Guidelines for Management of Patients With Ventricular Arrhythmias and the Prevention of Sudden Cardiac Death for patients with ventricular arrhythmia associated with torsade de pointes who present with long QT syndrome as well as for the treatment of patients with digoxin intoxication-induced arrhythmias. Magnesium is also the drug of choice in the management of pre-eclampsia and eclampsia. In addition to its therapeutic role, magnesium improves calcification. Patients with chronic kidney disease have a high prevalence of vascular calcification, and cardiovascular disease is the leading cause of death in this population. Several in vitro and animal studies point toward a protective role of magnesium through multiple molecular mechanisms. Magnesium is a natural calcium antagonist, and both human and animal studies have shown that low circulating magnesium levels are associated with vascular calcification. Results from an observational study conducted in the general Japanese population demonstrated that lower serum magnesium levels were significantly and independently associated with a greater average intima-media thickness and the risk of at least two carotid plaques. Magnesium supplementation might be useful in reducing the progression of atherosclerosis in chronic dialysis patients. Low serum magnesium may be an independent risk factor for death in patients with chronic kidney disease, and patients with mildly elevated serum magnesium levels could have a survival advantage over those with lower magnesium levels. Magnesium overdose Since the kidneys are responsible for the excretion of magnesium, anyone with a heart or kidney disorder should not take any extra magnesium except under their doctor's supervision. It is very rare to overdose on magnesium from food, however, people that ingest large amounts of milk of magnesia (as a laxative or antacid), epsom salts (as a laxative or tonic), or magnesium supplements may overdose, especially if they suffer from kidney problems. Too much magnesium can cause several serious health problems, including nausea, vomiting, severely lowered blood pressure, confusion, slowed heart rate, respiratory paralysis, deficiencies of other minerals, coma, cardiac arrhythmia, cardiac arrest, and eventually death. The most common side effects of magnesium toxicity are stomach upset

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and diarrhea. Safety precautions for magnesium metal Magnesium metal and its alloys are explosive hazards; they are highly flammable in their pure form when molten or in powder or ribbon form. Burning or molten magnesium metal reacts violently with water. When working with powdered magnesium, safety glasses with welding eye protection are employed, because the bright-white light produced by burning magnesium contains ultraviolet light that can permanently damage the retinas of the eyes. See also External links 

Magnesium at The Periodic Table of Videos (University of Nottingham)



Chemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World: Magnesium



Magnesium – a versatile and often overlooked element: new perspectives with a focus on chronic kidney disease

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Chapter 5: Calcium Calcium is the chemical element with symbol Ca and atomic number 20. Calcium is a soft gray alkaline earth metal, and is the fifth-most-abundant element by mass in the Earth's crust. Calcium is also the fifth-most-abundant dissolved ion in seawater by both molarity and mass, after sodium, chloride, magnesium, and sulfate. Calcium is essential for living organisms, in particular in cell physiology, where movement of the calcium ion Ca

into and out of the cytoplasm functions as a signal for many cellular

processes. As a major material used in mineralization of bone, teeth and shells, calcium is the most abundant metal by mass in many animals. Notable characteristics

Flame test. Brick-red color originates from calcium. In chemical terms, calcium is reactive and soft for a metal (though harder than lead, it can be

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cut with a knife with difficulty). It is a silvery metallic element that must be extracted by electrolysis from a fused salt like calcium chloride. Once produced, it rapidly forms a graywhite oxide and nitride coating when exposed to air. In bulk form (typically as chips or "turnings"), the metal is somewhat difficult to ignite, more so even than magnesium chips; but, when lit, the metal burns in air with a brilliant high-intensity orange-red light. Calcium metal reacts with water, generating hydrogen gas at a rate rapid enough to be noticeable, but not fast enough at room temperature to generate much heat, making it useful for generating hydrogen. In powdered form, however, the reaction with water is extremely rapid, as the increased surface area of the powder accelerates the reaction with the water. Part of the slowness of the calciumâ&#x20AC;&#x201C;water reaction results from the metal being partly protected by insoluble white calcium hydroxide. In water solutions of acids, where this salt is soluble, calcium reacts vigorously. 3

Calcium, with a density of 1.55 g/cm , is the lightest of the alkaline earth metals; magnesium (specific gravity 1.74) and beryllium (1.84) are more dense, although lighter in atomic mass. From strontium onward, the alkali earth metals become more dense with increasing atomic mass. It has two allotropes. Calcium has a higher electrical resistivity than copper or aluminium, yet weight-for-weight, due to its much lower density, it is a rather better conductor than either. However, its use in terrestrial applications is usually limited by its high reactivity with air. Calcium salts are colorless from any contribution of the calcium, and ionic solutions of 2+

calcium (Ca ) are colorless as well. As with magnesium salts and other alkaline earth metal salts, calcium salts are often quite soluble in water. Notable exceptions include the hydroxide, the sulfate (unusual for sulfate salts), the carbonate and the phosphates. With the exception of the sulfate, even the insoluble ones listed are in general more soluble than its transition metal counterparts. When in solution, the calcium ion to the human taste varies remarkably, being reported as mildly salty, sour, "mineral like" or even "soothing." It is apparent that many animals can taste, or develop a taste, for calcium, and use this sense to detect the mineral in salt licks or other sources. In human nutrition, soluble calcium salts may be added to tart juices without much effect to the average palate. Calcium is the fifth-most-abundant element by mass in the human body, where it is a common cellular ionic messenger with many functions, and serves also as a structural element in bone. It is the relatively high-atomic-number calcium in the skeleton that causes bone to be radio-opaque. Of the human body's solid components after drying and burning of organics (as for example, after cremation), about a third of the total "mineral" mass

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remaining, is the approximately one kilogram of calcium that composes the average skeleton (the remainder being mostly phosphorus and oxygen). H and K lines Visible spectra of many stars, including the Sun, exhibit strong absorption lines of singly ionized calcium. Prominent among these are the H-line at 3968.5 Ă&#x2026; and the K line at 3933.7 Ă&#x2026; of singly ionized calcium, or Ca II. For the Sun and stars with low temperatures, the prominence of the H and K lines can be an indication of strong magnetic activity in the chromosphere. Measurement of periodic variations of these active regions can also be used to deduce the rotation periods of these stars. Compounds Calcium, combined with phosphate to form hydroxylapatite, is the mineral portion of human and animal bones and teeth. The mineral portion of some corals can also be transformed into hydroxylapatite. Calcium hydroxide (slaked lime) is used in many chemical refinery processes and is made by heating limestone at high temperature (above 825 Â°C) and then carefully adding water to it. When lime is mixed with sand, it hardens into a mortar and is turned into plaster by carbon dioxide uptake. Mixed with other compounds, lime forms an important part of Portland cement. Calcium carbonate (CaCO3) is one of the common compounds of calcium. It is heated to form quicklime (CaO), which is then added to water (H 2O). This forms another material known as slaked lime (Ca(OH)2), which is an inexpensive base material used throughout the chemical industry. Chalk, marble, and limestone are all forms of calcium carbonate. When water percolates through limestone or other soluble carbonate rocks, it partially dissolves the rock and causes cave formation and characteristic stalactites and stalagmites and also forms hard water. Other important calcium compounds are calcium nitrate, calcium sulfide, calcium chloride, calcium carbide, calcium cyanamide and calcium hypochlorite. A few calcium compounds where calcium is in the oxidation state +1 have also been investigated recently. Isotopes Main article: Isotopes of calcium 40

Calcium has five stable isotopes ( Ca,

Ca,

Ca and

Ca), plus one more isotope

( Ca) that has such a long half-live that for all practical purposes it can also be considered

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stable. The 20% range in relative mass among naturally occurring calcium isotopes is greater than for any other element except hydrogen and helium. Calcium also has a cosmogenic isotope, radioactive

Ca, which has a half-life of 103,000 years. Unlike

cosmogenic isotopes that are produced in the atmosphere, activation of

Ca is produced by neutron

Ca. Most of its production is in the upper metre or so of the soil column, where

the cosmogenic neutron flux is still sufficiently strong. stellar studies because it decays to

Ca has received much attention in

K, a critical indicator of solar-system anomalies.

Ninety-seven percent of naturally occurring calcium is in the form of daughter products of

K decay, along with

Ca.

Ca is one of the

Ar. While K-Ar dating has been used extensively

in the geological sciences, the prevalence of

Ca in nature has impeded its use in dating.

Techniques using mass spectrometry and a double spike isotope dilution have been used for K-Ca age dating. The most abundant isotope,

Ca, has a nucleus of 20 protons and 20 neutrons. This is the

heaviest stable isotope of any element that has equal numbers of protons and neutrons. In supernova explosions, calcium is formed from the reaction of carbon with various numbers of alpha particles (helium nuclei), until the most common calcium isotope (containing 10 helium nuclei) has been synthesized.[citation needed] Isotope fractionation As with the isotopes of other elements, a variety of processes fractionate, or alter the relative abundance of, calcium isotopes. The best studied of these processes is the mass-dependent fractionation of calcium isotopes that accompanies the precipitation of calcium minerals, such as calcite, aragonite and apatite, from solution. Isotopically light calcium is preferentially incorporated into minerals, leaving the solution from which the mineral precipitated enriched in isotopically heavy calcium. At room temperature the magnitude of this fractionation is roughly 0.25â&#x20AC;° (0.025%) per atomic mass unit (AMU). Mass-dependent differences in calcium isotope composition conventionally are expressed the ratio of two isotopes (usually 44

Ca/ Ca) in a sample compared to the same ratio in a standard reference material. Ca/ Ca varies by about 1% among common earth materials.

Calcium isotope fractionation during mineral formation has led to several applications of calcium isotopes. In particular, the 1997 observation by Skulan and DePaolo that calcium minerals are isotopically lighter than the solutions from which the minerals precipitate is the basis of analogous applications in medicine and in paleooceanography. In animals with skeletons mineralized with calcium the calcium isotopic composition of soft tissues reflects the relative rate of formation and dissolution of skeletal mineral. In humans changes in the calcium isotopic composition of urine have been shown to be related to changes in bone mineral balance. When the rate of bone formation exceeds the rate of bone resorption, soft

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tissue

Ca/ Ca rises. Soft tissue

Ca/ Ca falls when bone resorption exceeds bone

formation. Because of this relationship, calcium isotopic measurements of urine or blood may be useful in the early detection of metabolic bone diseases like osteoporosis. A similar system exists in the ocean, where seawater removal of Ca

Ca/ Ca tends to rise when the rate of

from seawater by mineral precipitation exceeds the input of new calcium into

the ocean, and fall when calcium input exceeds mineral precipitation. It follows that rising 44

Ca/ Ca corresponds to falling seawater Ca

corresponds to rising seawater Ca

concentration, and falling

Ca/ Ca

concentration. In 1997 Skulan and DePaolo presented

the first evidence of change in seawater

Ca/ Ca over geologic time, along with a

theoretical explanation of these changes. More recent papers have confirmed this observation, demonstrating that seawater Ca

concentration is not constant, and that the

ocean probably never is in ―steady state‖ with respect to its calcium input and output. This has important climatological implications, as the marine calcium cycle is closely tied to the carbon cycle (see below). Geochemical cycling Calcium provides an important link between tectonics, climate and the carbon cycle. In the simplest terms, uplift of mountains exposes Ca-bearing rocks to chemical weathering and releases Ca

into surface water. This Ca

eventually is transported to the ocean where it

reacts with dissolved CO2 to form limestone. Some of this limestone settles to the sea floor where it is incorporated into new rocks. Dissolved CO 2, along with carbonate and bicarbonate ions, are referred to as dissolved inorganic carbon (DIC).

Travertine terraces Pamukkale, Turkey −

The actual reaction is more complicated and involves the bicarbonate ion (HCO 3 ) that forms when CO2 reacts with water at seawater pH: The result is that each Ca

ion released by chemical weathering ultimately removes one

CO2 molecule from the surficial system (atmosphere, ocean, soils and living organisms), storing it in carbonate rocks where it is likely to stay for hundreds of millions of years. The

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weathering of calcium from rocks thus scrubs CO 2 from the ocean and atmosphere, exerting a strong long-term effect on climate. Analogous cycles involving magnesium, and to a much smaller extent strontium and barium, have the same effect. As the weathering of limestone (CaCO3) liberates equimolar amounts of Ca

and CO2, it has

no net effect on the CO2 content of the atmosphere and ocean. The weathering of silicate rocks like granite, on the other hand, is a net CO 2 sink because it produces abundant Ca

but very little CO2. History Lime as building material was used since prehistoric times going as far back as 7000 to 14000 BC. The first dated lime kiln dates back to 2500 BC and was found in Khafajah mesopotamia. Calcium (from Latin calx, genitive calcis, meaning "lime") was known as early as the first century when the Ancient Romans prepared lime as calcium oxide. Literature dating back to 975 AD notes that plaster of paris (calcium sulfate), is useful for setting broken bones. It was not isolated until 1808 in England when Sir Humphry Davy electrolyzed a mixture of lime and mercuric oxide. Davy was trying to isolate calcium; when he heard that Swedish chemist Jöns Jakob Berzelius and Pontin prepared calcium amalgam by electrolyzing lime in mercury, he tried it himself. He worked with electrolysis throughout his life and also discovered/isolated sodium, potassium, magnesium, boron and barium. Calcium metal was not available in large scale until the beginning of the 20th century. Occurrence See also category: Calcium minerals Calcium is not naturally found in its elemental state. Calcium occurs most commonly in sedimentary rocks in the minerals calcite, dolomite and gypsum. It also occurs in igneous and metamorphic rocks chiefly in the silicate minerals: plagioclases, amphiboles, pyroxenes and garnets. Applications Calcium is used 

as a reducing agent in the extraction of other metals, such as uranium, zirconium, and thorium.



as a deoxidizer, desulfurizer, or decarbonizer for various ferrous and nonferrous alloys.



as an alloying agent used in the production of aluminium, beryllium, copper, lead,

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and magnesium alloys. 

in the making of cements and mortars to be used in construction.



in the making of cheese, where calcium ions influence the activity of rennin in bringing about the coagulation of milk.

Calcium compounds See also category: Calcium compounds 

Calcium carbonate (CaCO3) is used in manufacturing cement and mortar, lime, limestone (usually used in the steel industry) and aids in production in the glass industry. It also has chemical and optical uses as mineral specimens in toothpastes, for example.



Calcium hydroxide solution (Ca(OH)2) (also known as limewater) is used to detect the presence of carbon dioxide by being bubbled through a solution. It turns cloudy where CO2 is present.



Calcium arsenate (Ca3(AsO4)2) is used in insecticides.



Calcium carbide (CaC2) is used to make acetylene gas (for use in acetylene torches for welding) and in the manufacturing of plastics.



Calcium chloride (CaCl2) is used in ice removal and dust control on dirt roads, in conditioner for concrete, as an additive in canned tomatoes, and to provide body for automobile tires.



Calcium cyclamate (Ca(C6H11NHSO3)2) is used as a sweetening agent in several countries. In the United States it is no longer permitted for use because of suspected cancer-causing properties.



Calcium gluconate (Ca(C6H11O7)2) is used as a food additive and in vitamin pills.



Calcium hypochlorite (Ca(OCl)2) is used as a swimming pool disinfectant, as a bleaching agent, as an ingredient in deodorant, and in algaecide and fungicide.



Calcium permanganate (Ca(MnO4)2) is used in liquid rocket propellant, textile production, as a water sterilizing agent and in dental procedures.



Calcium phosphate (Ca3(PO4)2) is used as a supplement for animal feed, fertilizer, in commercial production for dough and yeast products, in the manufacture of glass,

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and in dental products. 

Calcium phosphide (Ca3P2) is used in fireworks, rodenticide, torpedoes and flares.



Calcium stearate (Ca(C18H35O2)2) is used in the manufacture of wax crayons, cements, certain kinds of plastics and cosmetics, as a food additive, in the production of water resistant materials and in the production of paints.



Calcium sulfate (CaSO4·2H2O) is used as common blackboard chalk, as well as, in its hemihydrate form better known as Plaster of Paris.



Calcium tungstate (CaWO4) is used in luminous paints, fluorescent lights and in Xray studies.



Hydroxylapatite (Ca5(PO4)3(OH), but is usually written Ca10(PO4)6(OH)2) makes up seventy percent of bone. Also carbonated-calcium deficient hydroxylapatite is the main mineral of which dental enamel and dentin are comprised.

Nutrition Main articles: Calcium in biology, Calcium metabolism, and Disorders of calcium metabolism Recommended adequate intake by the IOM for calcium: Age

Calcium (mg/day)

0–6 months

200

7–12 months

260

1–3 years

700

4–8 years

1000

9–18 years

1300

19–50 years

1000

51–70 years (male)

1000

51–70 years (female)

1200

71+ years

1200

Calcium is an important component of a healthy diet and a mineral necessary for life. The

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National Osteoporosis Foundation says, "Calcium plays an important role in building stronger, denser bones early in life and keeping bones strong and healthy later in life." Approximately 99 percent of the body's calcium is stored in the bones and teeth. The rest of the calcium in the body has other important uses, such as some exocytosis, especially neurotransmitter release, and muscle contraction. In the electrical conduction system of the heart, calcium replaces sodium as the mineral that depolarizes the cell, proliferating the action potential. In cardiac muscle, sodium influx commences an action potential, but during potassium efflux, the cardiac myocyte experiences calcium influx, prolonging the action potential and creating a plateau phase of dynamic equilibrium. Long-term calcium deficiency can lead to rickets and poor blood clotting and in case of a menopausal woman, it can lead to osteoporosis, in which the bone deteriorates and there is an increased risk of fractures. While a lifelong deficit can affect bone and tooth formation, over-retention can cause hypercalcemia (elevated levels of calcium in the blood), impaired kidney function and decreased absorption of other minerals. Several sources suggest a correlation between high calcium intake (2000 mg per day, or twice the U.S. recommended daily allowance, equivalent to six or more glasses of milk per day) and prostate cancer. High calcium intakes or high calcium absorption were previously thought to contribute to the development of kidney stones. However, a high calcium intake has been associated with a lower risk for kidney stones in more recent research.Vitamin D is needed to absorb calcium. Dairy products, such as milk and cheese, are a well-known source of calcium. Some individuals are allergic to dairy products and even more people, in particular those of non Indo-European descent, are lactose-intolerant, leaving them unable to consume nonfermented dairy products in quantities larger than about half a liter per serving. Others, such as vegans, avoid dairy products for ethical and health reasons. Many good vegetable sources of calcium exist, including seaweeds such as kelp, wakame and hijiki; nuts and seeds like almonds, hazelnuts, sesame, and pistachio; blackstrap molasses; beans (especially soy beans); figs; quinoa; okra; rutabaga; broccoli; dandelion leaves; and kale. In addition, for some drinks (like soy milk or orange juice[citation needed]) it is typical to be fortified with calcium. Numerous vegetables, notably spinach, chard and rhubarb have a high calcium content, but they may also contain varying amounts of oxalic acid that binds calcium and reduces its absorption. The same problem may to a degree affect the absorption of calcium from amaranth, collard greens, and chicory greens. This process may also be related to the generation of calcium oxalate. An overlooked source of calcium is eggshell, which can be ground into a powder and mixed into food or a glass of water.

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The calcium content of most foods can be found in the USDA National Nutrient Database. Dietary supplements

500 milligram calcium supplements made from calcium carbonate Calcium supplements are used to prevent and to treat calcium deficiencies. Office of Dietary Supplements (National Institutes of Health) recommends that no more than 600 mg of supplement should be taken at a time because the percent of calcium absorbed decreases as the amount of calcium in the supplement increases. It is therefore recommended to spread doses throughout the day. Recommended daily calcium intake for adults ranges from 1000 to 1300 mg. Calcium supplements may have side effects such as bloating and constipation in some people. It is suggested that taking the supplements with food may aid in nullifying these side effects. Vitamin D is added to some calcium supplements. Proper vitamin D status is important because vitamin D is converted to a hormone in the body, which then induces the synthesis of intestinal proteins responsible for calcium absorption. 

The absorption of calcium from most food and commonly used dietary supplements is very similar. This is contrary to what many calcium supplement manufacturers claim in their promotional materials.



Milk is an excellent source of dietary calcium for those whose bodies tolerate it because it has a high concentration of calcium and the calcium in milk is excellently absorbed.



Soymilk and other vegetable milks are usually sold with calcium added so that their calcium concentration is as high as in milk.



Also different kind of juices boosted with calcium are widely available.



Calcium carbonate is the most common and least expensive calcium supplement. It should be taken with food, and depends on low pH levels (acidic) for proper

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absorption in the intestine. Some studies suggests that the absorption of calcium from calcium carbonate is similar to the absorption of calcium from milk. While most people digest calcium carbonate very well, some might develop gastrointestinal discomfort or gas. Taking magnesium with it can help to avoid constipation. Calcium carbonate is 40% elemental calcium. 1000 mg will provide 400 mg of calcium. However, supplement labels will usually indicate how much calcium is present in each serving, not how much calcium carbonate is present. 

Antacids frequently contain calcium carbonate, and are a commonly used, inexpensive calcium supplement.



Coral calcium is a salt of calcium derived from fossilized coral reefs. Coral calcium is composed of calcium carbonate and trace minerals.



Calcium citrate can be taken without food and is the supplement of choice for individuals with achlorhydria or who are taking histamine-2 blockers or proton-pump inhibitors. Calcium citrate is about 21% elemental calcium. 1000 mg will provide 210 mg of calcium. It is more expensive than calcium carbonate and more of it must be taken to get the same amount of calcium.



Calcium phosphate costs more than calcium carbonate, but less than calcium citrate. Microcrystalline Hydroxyapatite (MH) is one of several forms of calcium phosphate used as a dietary supplement. Hydroxyapatite is about 40% calcium.



Calcium lactate has similar absorption as calcium carbonate, but is more expensive. Calcium lactate and calcium gluconate are less concentrated forms of calcium and are not practical oral supplements.



Calcium chelates are synthetic calcium compounds in which calcium is bound to an organic molecule, such as malate, aspartate, or fumarate. These forms of calcium may be better absorbed on an empty stomach. However, in general they are absorbed similarly to calcium carbonate and other common calcium supplements when taken with food. The "chelate" mimics the action that natural food performs by keeping the calcium soluble in the intestine. Thus, on an empty stomach, in some individuals, chelates might, in theory, be absorbed better.

Cardiovascular Impact A study investigating the effects of personal calcium supplement use on cardiovascular risk in the Women’s Health Initiative Calcium/Vitamin D Supplementation Study (WHI CaD Study) found a modestly increased risk of cardiovascular events, particularly myocardial infarction in

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postmenopausal women. A broad recommendation of calcium/vitamin D supplements is therefore not warranted. In contrast, the authors of a 2013 literature review concluded that the benefits of calcium supplementation, such as on bone health, appear to outweigh any risk calcium supplementation may theoretically pose to the cardiovascular health. Osteoporosis Such studies often do not test calcium alone, but rather combinations of calcium and vitamin D. Randomized controlled trials found both positive and negative effects. The different results may be explained by doses of calcium and underlying rates of calcium supplementation in the control groups. Cancer A meta-analysis by the international Cochrane Collaboration of two randomized controlled trials found that calcium "might contribute to a moderate degree to the prevention of adenomatous colonic polyps". More recent studies were conflicting, and one that was positive for effect (Lappe, et al.) did control for a possible anti-carcinogenic effect of vitamin D, which was found to be an independent positive influence from calcium-alone on cancer risk (see second study below). 

A randomized controlled trial found that 1000 mg of elemental calcium and 400 IU of vitamin D3 had no effect on colorectal cancer.



A randomized controlled trial found that 1400–1500 mg supplemental calcium and 1100 IU vitamin D3 reduced aggregated cancers with a relative risk of 0.402.



An observational cohort study found that high calcium and vitamin D intake was associated with "lower risk of developing premenopausal breast cancer."

Hazards and toxicity Compared with other metals, the calcium ion and most calcium compounds have low toxicity. This is not surprising given the very high natural abundance of calcium compounds in the environment and in organisms. Calcium poses few serious environmental problems, with kidney stones the most common side-effect in clinical studies. Acute calcium poisoning is rare, and difficult to achieve unless calcium compounds are administered intravenously. For 50

example, the oral median lethal dose (LD ) for rats for calcium carbonate and calcium chloride are 6.45 and 1.4 g/kg, respectively. Calcium metal is hazardous because of its sometimes-violent reactions with water and acids. Calcium metal is found in some drain cleaners, where it functions to generate heat and

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calcium hydroxide that saponifies the fats and liquefies the proteins (e.g., hair) that block drains. When swallowed calcium metal has the same effect on the mouth, esophagus and stomach, and can be fatal. Excessive consumption of calcium carbonate antacids/dietary supplements (such as Tums) over a period of weeks or months can cause milk-alkali syndrome, with symptoms ranging from hypercalcemia to potentially fatal renal failure. What constitutes ―excessive‖ consumption is not well known and, it is presumed, varies a great deal from person to person. Persons consuming more than 10 grams/day of CaCO3 (=4 g Ca) are at risk of developing milk-alkali syndrome, but the condition has been reported in at least one person consuming only 2.5 grams/day of CaCO3 (=1 g Ca), an amount usually considered moderate and safe. Oral calcium supplements diminish the absorption of thyroxine when taken within four to six hours of each other. Thus, people taking both calcium and thyroxine run the risk of inadequate thyroid hormone replacement and thence hypothyroidism if they take them simultaneously or near-simultaneously. Excessive calcium supplementation can be detrimental to cardiovascular health, especially in men. See also Further reading 

Rebecca J. Donatelle. Health, The Basics. 6th ed. San Francisco: Pearson Education, Inc. 2005.

External links 

WebElements.com – Calcium



Calcium at The Periodic Table of Videos (University of Nottingham)



USDA National Nutrient Database, Calcium content of selected foods



UK Food Standards Agency: Calcium



Nutrition fact sheet from the National Institutes of Health

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Chapter 6: Strontium Strontium is a chemical element with symbol Sr and atomic number 38. An alkaline earth metal, strontium is a soft silver-white or yellowish metallic element that is highly reactive chemically. The metal turns yellow when it is exposed to air. Strontium has physical and chemical properties similar to those of its two neighbors calcium and barium. It occurs naturally in the minerals celestine and strontianite. While natural strontium is stable, the synthetic

Sr isotope is present in radioactive fallout and has a half-life of 28.90 years.

Both strontium and strontianite are named after Strontian, a village in Scotland near which the mineral was first discovered in 1790 by Adair Crawford and William Cruickshank. The production of sugar from sugar beet was in the 19th century its largest application (see strontian process). Strontium compounds are today mostly used for the production of cathode ray tubes for televisions. The displacement of cathode ray tubes by other display methods in television sets is changing strontium's overall consumption. Characteristics

Oxidized dendritic strontium Strontium is a grey, silvery metal that is softer than calcium and even more reactive toward water, with which it reacts on contact to produce strontium hydroxide and hydrogen gas. It burns in air to produce both strontium oxide and strontium nitride, but since it does not react with nitrogen below 380 째C, at room temperature it will only form the oxide spontaneously. Three allotropes of metallic strontium exist, with transition points at 235 and 540 째C. Because of its extreme reactivity with oxygen and water, this element occurs naturally only in compounds with other elements, such as in the minerals strontianite and celestite. It is kept

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under a liquid hydrocarbon such as mineral oil or kerosene to prevent oxidation; freshly exposed strontium metal rapidly turns a yellowish color with the formation of the oxide. Finely powdered strontium metal is pyrophoric meaning it will ignite spontaneously in air at room temperature. Volatile strontium salts impart a bright red color to flames, and these salts are used in pyrotechnics and in the production of flares. Natural strontium is a mixture of four stable isotopes. History

Flame test for Strontium Strontium is named after the Scottish village of Strontian (Gaelic Sron an t-Sithein), having been discovered in the ores taken from the lead mines there. In 1790, Adair Crawford, a physician engaged in the preparation of barium, and his colleague William Cruickshank, recognised that the Strontian ores exhibited properties that differed from those normally seen in other "heavy spars" sources. This allowed Adair to conclude on page 355 "... it is probable indeed, that the scotch mineral is a new species of earth which has not hitherto been sufficiently examined." The physician and mineral collector Friedrich Gabriel Sulzer analysed together with Johann Friedrich Blumenbach the mineral from Strontian and named it strontianite. He also came to the conclusion that it was distinct from the witherite and contained a new earth (neue Grunderde). In 1793 Thomas Charles Hope, a professor of chemistry at the University of Glasgow proposed the name strontites. He confirmed the

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earlier work of Crawford and recounted: " ... Considering it a peculiar earth I thought it necessary to give it an name. I have called it Strontites, from the place it was found; a mode of derivation in my opinion, fully as proper as any quality it may possess, which is the present fashion." The element was eventually isolated by Sir Humphry Davy in 1808 by the electrolysis of a mixture containing strontium chloride and mercuric oxide, and announced by him in a lecture to the Royal Society on 30 June 1808. In keeping with the naming of the other alkaline earths, he changed the name to strontium. The first large scale application of strontium was in the production of sugar from sugar beet. Although a crystallisation process using strontium hydroxide was patented by AugustinPierre Dubrunfaut in 1849 the large scale introduction came with the improvement of the process in the early 1870s. The German sugar industry used the process well into the 20th century. Prior to World War I the beet sugar industry used 100,000 to 150,000 tons of strontium hydroxide for this process per year. The strontium hydroxide was recycled in the process, but the demand to substitute losses during production was high enough to create a significant demand initiating mining of strontianite in the M端nsterland. The mining of strontianite in Germany ended when mining of the celestite deposits in Gloucestershire started. These mines supplied most of the world strontium supply from 1884 to 1941. Although the celestite deposits in the Granada basin were known for some time the large scale mining did not start before the 1950s. During the atmospheric nuclear weapons testing it was observed that strontium-90 is one of the nuclear fission products with a relative high yield. The similarity to calcium and the chance that the strontium-90 might become enriched in bones made research on the metabolism of strontium an important topic. Occurrence

Strontium output in 2005 Strontium commonly occurs in nature, the 15th most abundant element on Earth, estimated to average approximately 360 parts per million in the Earth's crust and is found chiefly as the form of the sulfate mineral celestite (SrSO 4) and the carbonate strontianite (SrCO 3). Of the two, celestite occurs much more frequently in sedimentary deposits of sufficient size to make development of mining facilities attractive. Because strontium is used most often in the

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carbonate form, strontianite would be the more useful of the two common minerals, but few deposits have been discovered that are suitable for development. The mean strontium content of ocean water is 8 mg/l. At a concentration between 82 and 90 µmol/l of strontium the concentration is considerable lower than the calcium concentration which is normally between 9.6 and 11.6 mmol/l. Production According to the British Geological Survey, China was the top producer of strontium in 2007, with over two-thirds world share, followed by Spain, Mexico, Turkey, Argentina, and Iran. Large amounts of the mined celestite (SrSO 4) are converted to the carbonate by two processes. Either the celestite is directly leached with sodium carbonate solution or the celestite is roasted with coal to form the sulfide. The second process results a dark coloured material containing mostly strontium sulfide. This so-called black ash is dissolved in water and filtered. Strontium carbonate is precipitated from the strontium sulfide solution by introduction of carbon dioxide. The sulfate is reduced to the sulfide by the carbothermic reduction: SrSO4 + C → SrS + 2 CO2 About 300,000 tons are processed in this way annually. The metal is produced commercially by reducing strontium oxide with aluminium. The strontium is distilled from the mixture. Strontium metal can in principle be prepared by electrolysis of a solution of strontium chloride in molten potassium chloride: Sr

+ 2 e− → Sr −

2 Cl → Cl2 + 2 e− Isotopes Main article: Isotopes of strontium Strontium has four stable, naturally occurring isotopes: 88

(7.0%) and

Sr (82.58%). Only

radioactive alkali metal sources of 86

Sr, and

Sr (0.56%),

Sr (9.86%),

Sr is radiogenic; it is produced by decay from the

Rb, which has a half-life of 4.88 × 10

years. Thus, there are two

Sr in any material: first the portion formed in stars along with the isotopes

Sr; and second the portion formed by radioactive decay of

Sr,

Rb. The ratio

Sr/ Sr is the parameter typically reported in geologic investigations; ratios in minerals and

rocks have values ranging from about 0.7 to greater than 4.0. Because strontium has an

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atomic radius similar to that of calcium, it readily substitutes for Ca in minerals. Sixteen unstable isotopes are known to exist. Of greatest importance are of 28.78 years and

Sr with a half-life of 50.5 days.

Sr with a half-life

Sr is a by-product of nuclear fission

found in nuclear fallout and presents a health problem since it substitutes for calcium in bone, preventing expulsion from the body. This isotope is one of the best long-lived highenergy beta emitters known, and is used in SNAP (Systems for Nuclear Auxiliary Power) devices. These devices hold promise for use in spacecraft, remote weather stations, navigational buoys, etc., where a lightweight, long-lived, nuclear-electric power source is required. The 1986 Chernobyl nuclear accident contaminated a vast area with

Sr.

confined inside a concave silver plaque is also used for the medical treatment of a resected pterygium. 89

Sr is a short-lived artificial radioisotope that is used in the treatment of bone cancer. In

circumstances where cancer patients have widespread and painful bony metastases (secondaries), the administration of

Sr results in the delivery of radioactive emissions (beta

particles in this case) directly to the area of bony problem (where calcium turnover is greatest). The

Sr is manufactured as the chloride salt (which is soluble), and when

dissolved in normal saline can be injected intravenously. Typically, cancer patients will be treated with a dose of 150 MBq. Patients must take precautions following this because their urine becomes contaminated with radioactivity, so they must sit to urinate and double-flush the toilet. The beta particles travel about 3.5 mm in bone (energy 0.583 MeV) and 6.5 mm in tissue, so there is no requirement to isolate patients having been treated, except to say they should not have any one (especially young children) sitting in their laps for 10â&#x20AC;&#x201C;40 days.[citation needed] The variation in time results from the variable clearing time for

Sr,

which depends on renal function and the number of bony metastases. With a lot of bony metastases, the entire

Sr dose can be taken up into bone and so the radioactivity is

retained to decay over a 50.5-day half-life. It takes about 10 half-lives or about 500 days for 99.9% of the radioactive strontium to decay. However, where there are few bony metastases, the large proportion of

Sr not taken up by the bone will be filtered by the kidney, so that the

effective half-life (a combination of the physical and biological half-life) will be much shorter. Applications

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CRT computer monitor front panel made from strontium and barium oxide-containing glass. This application consumes most of the world's production of strontium. Consuming 75% of production, the primary use for strontium is in glass for colour television cathode ray tubes. It prevents X-ray emission. All parts of the CRT must absorb X-rays. In the neck and the funnel of the tube, lead glass is used for this purpose, but this type of glass shows a browning effect due to the interaction of the X-rays with the glass. Therefore, the front panel has to use a different glass mixture, in which strontium and barium are the X-rayabsorbing materials. The average values for the glass mixture determined for a recycling study in 2005 is 8.5% strontium oxide and 10% barium oxide. The amount of strontium used for the production of cathode ray tube is declining because the CRTs are replaced by other display methods. This decline has a significant influence on the mining and refining of strontium. Because strontium is so similar to calcium, it is incorporated in the bone. All four stable isotopes are incorporated, in roughly similar proportions, as they are found in nature. However, the actual distribution of the isotopes tends to vary greatly from one geographical location to another. Thus, analyzing the bone of an individual can help determine the region it came from. This approach helps to identify the ancient migration patterns as well as the origin of commingled human remains in battlefield burial sites. Strontium, thus, helps forensic scientists too. 87

Sr/ Sr ratios are commonly used to determine the likely provenance areas of sediment in

natural systems, especially in marine and fluvial environments. Dasch (1969) showed that surface sediments of Atlantic displayed averages of the

Sr/ Sr ratios that could be regarded as bulk

Sr/ Sr ratios of geological terranes from adjacent landmasses. A good

example of a fluvial-marine system to which Sr isotope provenance studies have been successfully employed is the River Nile-Mediterranean system, Due to the differing ages of the rocks that constitute the majority of the Blue and White Nile, catchment areas of the

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changing provenance of sediment reaching the River Nile delta and East Mediterranean Sea can be discerned through Sr isotopic studies. Such changes are climatically controlled in the Late Quaternary. More recently,

Sr/ Sr ratios have also been used to determine the source of ancient 87

archaeological materials such as timbers and corn in Chaco Canyon, New Mexico. Sr/ Sr ratios in teeth may also be used to track animal migrations or in criminal forensics. Pyrotechnics Strontium carbonate or other strontium salts are used in the manufacture of fireworks, as they impart a deep red color to the firework. This application consumes about 5% of the world's production. Uses for radioactive strontium 89

Sr is the active ingredient in Metastron (the generic version of Metastron, Generic

Strontium

Chloride

Sr-89

Injection,

its

manufactured

Bio-Nucleonics

Inc.),

radiopharmaceutical used for bone pain secondary to metastatic bone cancer. The strontium acts like calcium and is preferentially incorporated into bone at sites of increased osteogenesis. This localization focuses the radiation exposure on the cancerous lesion.

RTGs from Soviet era lighthouses 90

Sr has been used as a power source for radioisotope thermoelectric generators (RTGs).

Sr produces approximately 0.93 watts of heat per gram (it is lower for the form of

in RTGs, which is strontium fluoride). However, shorter and has a lower density than that it is cheaper than

238

Sr used

Sr has a lifetime approximately 3 times

Pu, another RTG fuel. The main advantage of

Sr is

Pu and is found in nuclear waste. The Soviet Union deployed nearly

1000 of these RTGs on its northern coast as a power source for lighthouses and meteorology stations. 90

Sr is also used in cancer therapy. Its beta emission and long half-life is ideal for superficial

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radiotherapy. Niche applications Strontium chloride is sometimes used in toothpastes for sensitive teeth. One popular brand includes 10% total strontium chloride hexahydrate by weight. Small amounts are used in the refining of zinc, to remove small amounts of lead impurities. Research trends Other possible applications follow: 

Strontium titanate has an extremely high refractive index and an optical dispersion greater than that of diamond, making it useful in a variety of optics applications. This quality has also led to its being cut into gemstones, in particular as a diamond simulant. However, it is very soft and easily scratches so it is rarely used.



Ferrite magnets.



Strontium aluminate is used as a bright phosphor with long persistence of phosphorescence.



Strontium oxide is sometimes used to improve the quality of some pottery glazes.



Strontium ranelate is used in the treatment of osteoporosis. It is a prescription drug in the EU, but not in the USA.



Strontium barium niobate can be used in outdoors holographic 3D displays as a "screen".

Strontium metal is used in strontium 90%-aluminium 10% alloys of an eutectic composition for the modification of aluminium-silicon casting alloys.AJ62, a durable, creep-resistant magnesium alloy used in car and motorcycle engines by BMW, contains 2% strontium by weight. Strontium is used in scientific studies of neurotransmitter release in neurons. Like calcium, strontium facilitates synaptic vesicle fusion with the synaptic membrane. But, unlike calcium, strontium causes asynchronous vesicle fusion. Therefore, replacing calcium in a culture medium with strontium allows scientists to measure the effects of a single-vesicle fusion event, e.g., the size of the postsynaptic response elicited by the neurotransmitter content of a single vesicle.

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The important concept for isotopic tracing is that Sr derived from any mineral through weathering reactions will have the same 87

Sr/ Sr as the mineral. Therefore, differences in

Sr/ Sr among ground waters require either (a) differences in mineralogy along contrasting

flowpaths or (b) differences in the relative amounts of Sr weathered from the same suite of minerals. This latter situation can arise in several ways. First, differences in initial water chemistry within a homogeneous rock unit will affect the relative weathering rates of the minerals. For example, sections of the soil zone affected by evaporative concentration of recharge waters or by differences in pCO2 can be expected to have different

Sr/ Sr.

Secondly, differences in the relative mobilities of water at scales ranging from inter-grain pores to the catchment scale may also profoundly affect example, the chemical composition and the resultant

Sr/ Sr (Bullen et al., 1996). For 86

Sr/ Sr in immobile waters at a

plagioclase-hornblende grain boundary versus a quartz-mica boundary will be different. Third, a difference in the relative "effective" surface areas of minerals in one portion of the rock unit will also cause differences in chemistry and isotopic composition; "poisoning" of reactive surfaces by organic coatings is an example of this kind of process. In a fundamental sense, because the waters in shallow systems are not in chemical equilibrium with the rocks, it is unrealistic to expect that waters along flowpaths within even a constant-mineralogy unit should have a constant

Sr/ Sr. Instead, the waters moving along specific flowpaths slowly

react with the rocks and gradually approach chemical equilibrium over long time-periods. Compounds

The mineral celestine (SrSO4) illustrates the fact that most Sr compounds are colourless or white. See also: Category:Strontium compounds Strontium forms a variety of salts, the properties of which are always intermediate between those of barium and calcium. The salts tend to be colourless. The sulfate and carbonate are poorly soluble, hence their occurrence as minerals. Most compounds are derived from the carbonate or the sulfide, which is obtained from the minerals. Typical for an alkaline earth derivative, the sulfide hydrolyzes readily:

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SrS + 2 H2O → Sr(OH)2 + H2S Similar reactions are used in the production of commercially useful compounds, including the most useful strontium compound, strontium carbonate. SrS + H2O + CO2 → SrCO3 + H2S Strontium nitrate can also be prepared in this way. Biological role Acantharea a relative large group of marine radiolarian protozoa produce intricate mineral skeletons composed of strontium sulfate. In biological systems calcium is substituted in a small extent by strontium. In the human body most of the absorbed strontium is deposited in the bones. The ratio of strontium to calcium in human bones is between 1:1000 and 1:2000 roughly in the same range as in the blood serum. Effect on the human body The human body absorbs strontium as if it were calcium. Due to the chemical similarity of the elements, the stable forms of strontium might not pose a significant health threat — in fact, the levels found naturally may actually be beneficial (see below) – but the radioactive

can lead to various bone disorders and diseases, including bone cancer. The strontium unit is used in measuring radioactivity from absorbed

Sr.

A recent in-vitro study conducted the NY College of Dental Sciences using strontium on osteoblasts showed marked improvement on bone-building osteoblasts. The drug strontium ranelate, made by combining strontium with ranelic acid, was found to aid bone growth, increase bone density, and lessen vertebral, peripheral, and hip fractures. Women receiving the drug showed a 12.7% increase in bone density. Women receiving a placebo had a 1.6% decrease. Half the increase in bone density (measured by X-ray densitometry) is attributed to the higher atomic weight of Sr compared with calcium, whereas the other half a true increase in bone mass. Strontium ranelate is registered as a prescription drug in Europe and many countries worldwide. It must be prescribed by a doctor, must be delivered by a pharmacist, and requires strict medical supervision. There is a long history of medical research regarding strontium's benefits, beginning in the 1950s. Studies indicate a lack of undesirable side-effects. Several other salts of strontium such as strontium citrate and strontium carbonate are available in the United States under the Dietary Supplements Health and Education Act of 1994, providing close to the recommended strontium content, about 680 milligrams per day, of strontium ranelate. Their

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long-term safety and efficacy have not been evaluated on humans in large-scale medical trials.[citation needed] However, some companies do manufacture strontium pills for increasing bone health. See also External links 

WebElements.com – Strontium



Chemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World: Strontium



Strontium at The Periodic Table of Videos (University of Nottingham)

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Chapter 7: Barium Barium is a chemical element with symbol Ba and atomic number 56. It is the fifth element in Group 2, a soft silvery metallic alkaline earth metal. Because of its high chemical reactivity barium is never found in nature as a free element. Its hydroxide was known in pre-modern history as baryta; this substance does not occur as a mineral, but can be prepared by heating barium carbonate. The most common naturally occurring minerals of barium are barite (barium sulfate, BaSO 4) and witherite (barium carbonate, BaCO3), both being insoluble in water. Barium's name originates from the alchemical derivative "baryta", which itself comes from Greek βαρύς (barys), meaning "heavy." Barium was identified as a new element in 1774, but not reduced to a metal until 1808, shortly after electrolytic isolation techniques became available. Barium has only a few industrial applications. The metal has been historically used to scavenge air in vacuum tubes. It is a component of YBCO (high-temperature superconductors) and electroceramics, and is added to steel and cast iron to reduce the size of carbon grains within the microstructure of the metal. Barium compounds are added to fireworks to impart a green color. Barium sulfate is used as an insoluble heavy additive to oil well drilling fluid, and in purer form, as X-ray radiocontrast agents for imaging the human gastrointestinal tract. Soluble barium compounds are poisonous due to release of the soluble barium ion, and therefore have been used as rodenticides. Characteristics Physical properties

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Oxidized barium Barium is a soft, silvery-white metal, with a slight golden shade when ultrapure. The silverywhite color of barium metal rapidly vanishes upon oxidation in air yielding a dark gray oxide layer. Barium has a medium specific weight and good electrical conductivity. Ultrapure barium is very hard to prepare, and therefore many properties of barium have not been accurately measured yet. At room temperature and pressure, barium has a body-centered cubic structure, with a barium–barium distance of 503 picometers, expanding with heating at a rate of −5

approximately 1.8×10 /°C. It is a very soft metal with a Mohs hardness of 1.25. Its melting temperature of 1,000 K (730 °C; 1,340 °F) is intermediate between those of the lighter strontium (1,050 K or 780 °C; 1,430 °F) and heavier radium (973 K or 700 °C; 1,292 °F); however, its boiling point of 2,170 K (1,900 °C; 3,450 °F) exceeds that of strontium (1,655 K −3

or 1,382 °C; 2,519 °F). The density (3.62 g·cm ) is again intermediate between those of −3

−3

strontium (2.36 g·cm ) and radium (~5 g·cm ). Chemical reactivity Barium is chemically similar to magnesium, calcium, and strontium, being even more reactive. It always exhibits the oxidation state of +2. Reactions with chalcogens are highly exothermic (release energy); the reaction with oxygen or air occurs at room temperature, and therefore barium is stored under oil or inert gas atmosphere. Reactions with other nonmetals, such as carbon, nitrogen, phosphorus, silicon, and hydrogen, are generally exothermic and proceed upon heating. Reactions with water and alcohols are also very exothermic and release hydrogen gas:

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Ba + 2 ROH → Ba(OR)2 + H2↑ (R is an alkyl or a hydrogen atom) Additionally, barium reacts with ammonia to form complexes such as Ba(NH 3)6. The metal is readily attacked by most acids. Sulfuric acid is a notable exception, as passivation stops the reaction by forming the insoluble barium sulfate. Barium combines with several metals, including aluminium, zinc, lead, and tin, forming intermetallic phases and alloys. Compounds Selected alkaline earth and zinc salts densities, g·cm O2−

Ca2+

Sr2+

Ba2+

Zn2+

S2−

F−

Cl−

−3

SO2−

CO2−

O2−

H−

3.34

2.59

3.18

2.15

2.96

2.83

2.9

1.7

5.1

3.7

4.24

3.05

3.96

3.5

4.78

3.26

5.72

4.3

4.89

3.89

4.49

4.29

4.96

4.16

5.6

4.09

4.95

2.09

3.54

4.4

1.57

—

Barium salts are typically white when solid and colorless when dissolved, as barium ions provide no specific coloring. They are also denser than their strontium or calcium analogs, except for the halides (see table; zinc is given for comparison). Barium hydroxide ("baryta") was known to alchemists who produced it by heating barium carbonate. Unlike calcium hydroxide, it absorbs very little CO 2 in aqueous solutions and is therefore insensitive to atmospheric fluctuations. This property is used in calibrating pH equipment. Volatile barium compounds burn with a green to pale green flame, which is an efficient test to detect a barium compound. The color results from spectral lines at 455.4, 493.4, 553.6, and 611.1 nm. Organobarium compounds are a growing class of compounds: for example, dialkylbariums are known, as are alkylhalobariums.

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Isotopes Main article: Isotopes of barium Barium occurs naturally on Earth as a mixture of seven primordial nuclides, barium-130, 132, and 134 through 138. The first two are thought to be radioactive: barium-130 should decay to xenon-130 via double beta plus decay, and barium-132 should similarly decay to xenon-132. The corresponding half-lives should exceed the age of the Universe by at least thousand times. Their abundances are ~0.1% relative to that of natural barium. Their radioactivity is so weak that they pose no danger to life. Out of the stable isotopes, barium-138 makes up 71.7% of all barium, and the lighter the isotope, the less it is abundant. In total, barium has about 50 known isotopes, ranging in mass between 114 and 153. The most stable metastable isotope is barium-133, which has a half-life of approximately 10.51 years, and five more isotopes have their half-lives longer than a day. Barium also has 10 meta states, out of which barium-133m1 is the most stable, having a half-live of about 39 hours. History

Sir Humphry Davy, who first isolated barium metal Alchemists in the early Middle Ages knew about some barium minerals. Smooth pebble-like stones of mineral barite found in Bologna, Italy, were known as "Bologna stones." Witches and alchemists were attracted to them because after exposure to light they would glow for years. The phosphorescent properties of barite heated with organics were described by V. Casciorolus in 1602. Carl Scheele identified barite as containing a new element in 1774, but could not isolate barium, only barium oxide. Johan Gottlieb Gahn also isolated barium oxide two years later in similar studies. Oxidized barium was at first called "barote," by Guyton de Morveau, a name that was changed by Antoine Lavoisier to baryta. Also in the 18th century, English

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mineralogist William Withering noted a heavy mineral in the lead mines of Cumberland, now known to be witherite. Barium was first isolated by electrolysis of molten barium salts in 1808, by Sir Humphry Davy in England. Davy, by analogy with calcium named "barium" after baryta, with the "-ium" ending signifying a metallic element.Robert Bunsen and Augustus Matthiessen obtained pure barium by electrolysis of a molten mixture of barium chloride and ammonium chloride. The production of pure oxygen in the Brin process was a large-scale application of barium peroxide in the 1880s, before it was replaced by electrolysis and fractional distillation of liquefied air in the early 1900s. In this process barium oxide reacts at 500–600 °C (932– 1,112 °F) with air to form barium peroxide, which decomposes at above 700 °C (1,292 °F) by releasing oxygen: 2 BaO + O2 ⇌ 2 BaO2 In 1908, barium sulfate was first applied as a radiocontrast agent in X-ray imaging of the digestive system. Occurrence and production The abundance of barium is 0.0425% in the Earth's crust and 13 µg/L in sea water. The main commercial source of barium is barite (also called barytes or heavy spar), which is a barium sulfate mineral. Its deposits are spread all over the world. The only other commercial source is far less important than barite; it is witherite, a barium carbonate mineral. Its main deposits are located in England, Romania, and the former USSR.

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Barite, left to right: appearance, graph showing trends in production over time, and the map showing shares of the most important producer countries in 2010. The barite reserves are estimated between 0.7 and 2 billion tonnes. The maximum production was achieved in 1981, at 8.3 million tonnes, and only 7–8% of it was used to make barium or its compounds. The barite production has again risen since the second half of the 1990s: from 5.6 million tonnes in 1996 to 7.6 in 2005 and 7.8 in 2011. China accounts for more than 50% of this output, followed by India (14% in 2011), Morocco (8.3%), US (8.2%), Turkey (2.5%), Iran and Kazakhstan (2.6% each). The mined ore is washed, crushed, classified, and separated from quartz. If the quartz penetrates too deep into the ore, or the iron, zinc, or lead content is abnormally high, then froth flotation methods are applied. The product is a 98% pure barite (by mass); the purity should be no less than 95%, with a minimal content of iron and silicon dioxide. It is then reduced by carbon to barium sulfide: BaSO4 + 2 C → BaS + 2 CO2↑ The water-soluble barium sulfide is the starting point for other compounds: dissolved BaS upon reaction with oxygen gives the hydroxide, with nitric acid the nitrate, with carbon dioxide the carbonate, and so on. The nitrate can be thermally decomposed to yield the oxide. Barium metal is produced by reduction with aluminium at 1,100 °C (2,010 °F). The

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intermetallic compound BaAl4 is produced first: 3 BaO + 14 Al → 3 BaAl4 + Al2O3 It is an intermediate, which reacts with barium oxide to give the metal. Note that not all barium is reduced. 8 BaO + BaAl4 → Ba↑ + 7 BaAl2O4 The remaining barium oxide reacts with the formed aluminium oxide: BaO + Al2O3 → BaAl2O4 and the overall reaction is 4 BaO + 2 Al → 3 Ba↑ + BaAl2O4 The thus produced barium vapor is collected at the cooler part of the apparatus and then packed into molds under argon atmosphere. This method is used commercially and can yield ultrapure barium. Commonly sold barium is about 99% pure, with main impurities being strontium and calcium (up to 0.8% and 0.25%) and other contaminants contributing less than 0.1%. A similar reaction with silicon at 1,200 °C (2,190 °F) yields barium and barium metasilicate. Electrolysis is not used because barium readily dissolves in molten halides and is rather impure when isolated with this method.

Benitoite crystals on natrolite. The mineral is named for the San Benito River in San Benito County where it was first found. Gemstone A barium-containing mineral benitoite (barium titanium silicate) occurs as a very rare blue fluorescent gemstone, and is the official state gem of California. Applications

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Metal and alloys Barium, as a metal or when alloyed with aluminium, is used to remove unwanted gases (gettering) from vacuum tubes, such as TV picture tubes. Barium is suitable for this purpose because of its low vapor pressure and reactivity towards oxygen, nitrogen, carbon dioxide, and water; it can even partly remove noble gases by dissolving them in the crystal lattice. This application is gradually disappearing due to the rising popularity of the tubeless LCD and plasma sets. Other uses of elemental barium are minor and include an additive to silumin (aluminium– silicon alloys) that refines their structure, as well as 

bearing alloys;



lead–tin soldering alloys – to increase the creep resistance;



alloy with nickel for spark plugs;



additive to steel and cast iron as an inoculant;



alloys with calcium, manganese, silicon, and aluminium as high-grade steel deoxidizers.

Barium sulfate and barite

Amoebiasis as seen in radiograph of barium-filled colon Barium sulfate (the mineral barite, BaSO4) is important to the petroleum industry, for example, as a drilling fluid in oil and gas wells. The precipitate of the compound (called "blanc fixe", from a French expression meaning "permanent white") is used in paints and varnishes, and also as a filler in ringing ink, plastics, and rubbers. It is also a paper coating pigment. In the form of nanoparticles, it can improve physical properties of some polymers, such as epoxies.

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Barium sulfate has a low toxicity and relatively high density of ca. 4.5 g·cm

−3

(and thus

opacity to X-rays). For this reason it is used as a radiocontrast agent in X-ray imaging of the digestive system ("barium meals" and "barium enemas").Lithopone, a pigment that contains barium sulfate and zinc sulfide, is a permanent white that has good covering power, and does not darken when exposed to sulfides. Other barium compounds

Green barium fireworks Aside from the sulfate, other compounds of barium find only niche applications. Applications are limited by the toxicity of Ba

ions (barium carbonate is a rat poison), which is not a

problem for the insoluble BaSO4. 

Barium oxide is used in a coating for the electrodes of fluorescent lamps, which facilitates the release of electrons.



Barium carbonate is used in glassmaking. Being a heavy element, barium increases the refractive index and luster of the glass. The compound is also used to reduce leaks of X-rays from cathode ray tubes (CRT) TV sets.



Barium, typically as barium nitrate, is added to fireworks to impart them a green color. The species responsible for the brilliant green is barium monochloride; in the absence of chlorine a yellow or "apple" green is produced instead.



Barium peroxide can be used as a catalyst to start an aluminothermic reaction when welding rail tracks together. It can also be used in green tracer ammunition and as a bleaching agent.



Barium titanate is a promising electroceramic.

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

Barium fluoride is used for optics in infrared applications because of its wide transparency range of 0.15–12 micrometers.



YBCO was the first high-temperature superconductor that could be cooled by liquid nitrogen, as its transition temperature of 93 K (−180.2 °C; −292.3 °F) exceeded the boiling point of nitrogen (77 K or −196.2 °C; −321.1 °F).

Biological dangers and precautions Because of the high reactivity of the metal, toxicological data are available only for compounds. Water-soluble barium compounds are poisonous. At low doses, barium ions act as a muscle stimulant, whereas higher doses affect the nervous system, causing cardiac irregularities, tremors, weakness, anxiety, dyspnea and paralysis. This may be due to the ability of Ba

to block potassium ion channels, which are critical to the proper function of the

nervous system. Other target organs for water-soluble barium compounds (i.e., barium ions) are eyes, immune system, heart, respiratory system, and skin. They affect the body strongly, causing, for example, blindness and sensitization. Barium is not carcinogenic, and it does not bioaccumulate. However, inhaled dust containing insoluble barium compounds can accumulate in the lungs, causing a benign condition called baritosis. For comparison to the soluble poisons, the insoluble sulfate is nontoxic and is thus not classified as a dangerous good. To avoid a potentially vigorous chemical reaction, barium metal is kept under argon or mineral oils. Contact with air is dangerous, as it may cause ignition. Moisture, friction, heat, sparks, flames, shocks, static electricity, reactions with oxidizers and acids should be avoided. Everything that may make contact with barium should be grounded. Those who work with the metal should wear pre-cleaned non-sparking shoes, flame-resistant rubber clothes, rubber gloves, apron, goggles, and a gas mask; they are not allowed to smoke in the working area and must wash themselves after handling barium. See also 

Han purple and Han blue

External links 

Barium at The Periodic Table of Videos (University of Nottingham)



Elementymology & Elements Multidict



3-D Holographic Display Using Strontium Barium Niobate

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Chapter 8: Radium Radium is a chemical element with symbol Ra and atomic number 88. Radium is an almost pure-white alkaline earth metal, but it readily oxidizes on exposure to air, becoming black in color. All isotopes of radium are highly radioactive, with the most stable isotope being radium-226, which has a half-life of 1601 years and decays into radon gas. Because of such instability, radium is luminescent, glowing a faint blue. Radium, in the form of radium chloride, was discovered by Marie Curie and Pierre Curie in 1898. They extracted the radium compound from uraninite and published the discovery at the French Academy of Sciences five days later. Radium was isolated in its metallic state by Marie Curie and André-Louis Debierne through the electrolysis of radium chloride in 1910. Since its discovery, it has given names like radium A and radium C 2 to several isotopes of other elements that are decay products of radium-226. In nature, radium is found in uranium ores in trace amounts as small as a seventh of a gram per ton of uraninite. Radium is not necessary for living organisms, and adverse health effects are likely when it is incorporated into biochemical processes because of its radioactivity and chemical reactivity. Characteristics Physical characteristics Although radium is not as well studied as its stable lighter homologue barium, the two elements have very similar properties. Their first two ionization energies are very similar: 509.3 and 979.0 kJ·mol

−1

for radium and 502.9 and 965.2 kJ·mol

−1

for barium. Such low

figures yield both elements' high reactivity and the formation of the very stable Ra

ion and

similar Ba . Pure radium is a white, silvery, solid metal, melting at 700 °C (1292 °F) and boiling at 1737 3

°C (3159 °F), similar to barium. Radium has a density of 5.5 g/cm ; the radium-barium density ratio is comparable to the radium-barium atomic mass ratio, as these elements have very similar body-centered cubic structures. Chemical characteristics and compounds See also category: Radium compounds Radium is the heaviest known alkaline earth metal; its chemical properties mostly resemble those of barium. When exposed to air, radium reacts violently with it, forming radium nitride,

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which causes blackening of this white metal. It exhibits only the +2 oxidation state in solution. Radium ions do not form complexes easily, due to the highly basic character of the ions. Most radium compounds coprecipitate with all barium, most strontium, and most lead compounds, and are ionic salts. The radium ion is colorless, making radium salts white when freshly prepared, turning yellow and ultimately dark with age owing to self-decomposition from the alpha radiation.Compounds of radium flame red-purple and give a characteristic spectrum. Like other alkaline earth metals, radium reacts violently with water to form radium hydroxide and is slightly more volatile than barium. Because of its geologically short half-life and intense radioactivity, radium compounds are quite rare, occurring almost exclusively in uranium ores. Radium chloride, radium bromide, radium hydroxide, and radium nitrate are soluble in water, with solubilities slightly lower than those of barium analogs for bromide and chloride, and higher for nitrate. Radium hydroxide is more soluble than hydroxides of other alkaline earth metals, actinium, and thorium, and more basic than barium hydroxide. It can be separated from these elements by their precipitation with ammonia. Insoluble radium compounds include radium sulfate, radium chromate, radium iodate, radium carbonate, and radium tetrafluoroberyllate; the radium sulfate is the most insoluble known sulfate. Radium oxide remains uncharacterized, despite the fact that oxides are common compounds for other alkaline-earth metals. The 6s and 6p electrons participate in the bonding in radium fluoride and radium astatide, making the bonding there more covalent in character. Isotopes Main article: Isotopes of radium Radium has 25 different known isotopes, four of which are found in nature, with the most common.

223

Ra,

224

Ra,

226

either uranium (U) or thorium (Th).

Ra and

228

226

Ra being

Ra are all generated naturally in the decay of

226

Ra is a product of

238

U decay, and is the longest-lived

isotope of radium with a half-life of 1601 years; next longest is

228

Ra, a product of

232

breakdown, with a half-life of 5.75 years. Radium has no stable isotopes; however, four isotopes of radium are present in decay chains,all of which are present in trace amounts. The most abundant and the longest-living one is radium-226, with a half-life of 1601 years. To date, 34 isotopes of radium have been synthesized, ranging in mass number from 202 to 234. At least 12 nuclear isomers have been reported; the most stable of them is radium-205m, with a half-life of between 130 and 230 milliseconds. All ground states of isotopes from radium-205 to radium-214, and from radium-221 to radium-234, have longer ones.

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Three other natural radioisotopes had received historical names in the early 20th century: radium-223 was known as actinium X, radium-224 as thorium X and radium-228 as mesothorium I. Radium-226 has given historical names to its decay products after the whole element, such as radium A for polonium-218. Radioactivity Radium is three million times as radioactive as the same mass of uranium. Its decay occurs in at least seven stages; the successive main products have been studied and were called radium emanation or exradio (now identified as radon), radium A (polonium), radium B (lead), radium C (bismuth), etc. Radon is a heavy gas, and the later products are solids. These products are themselves radioactive elements, each with an atomic weight a little lower than its predecessor. Radium loses about 1% of its activity in 25 years, being transformed into elements of lower atomic weight, with lead being the final product of disintegration. The SI unit of radioactivity is the becquerel (Bq), equal to one disintegration per second. The curie is a non-SI unit defined as that amount of radioactive material that has the same disintegration rate as 1 gram of radium-226 (3.7Ă&#x2014;10

disintegrations per second, or 37

GBq). Radium metal maintains itself at a higher temperature than its surroundings because of the radiation it emits â&#x20AC;&#x201C; alpha particles, beta particles, and gamma rays. More specifically, the alpha particles are produced by the radium decay, whereas the beta particles and gamma rays are produced by relatively short-half-life elements further down the decay chain. Occurrence All radium occurring today is produced by the decay of heavier elements, being present in decay chains. Owing to such short half-lives of its isotopes, radium is not primordial but trace. It cannot occur in large quantities due both to the fact that isotopes of radium have short half-lives and that parent nuclides have very long ones. Radium is found in tiny quantities in the uranium ore uraninite and various other uranium minerals, and in even tinier quantities in thorium minerals. Radium-226 is a decay product of uranium and is therefore found in all uranium-bearing ores. (One ton of pitchblende typically yields about one seventh of a gram of radium). Production Uranium had no large scale application in the late 18th century and therefore no large

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uranium mines existed. In the beginning the only larger source for uranium ore was the silver mines at Joachimsthal (now Jáchymov) in the Austrian Empire. The uranium ore was only a by-product of the mining activities. After the isolation of radium by Marie and Pierre Curie from uranium ore from Joachimsthal several scientists started to isolate radium in small quantities. Later small companies purchased mine tailings from Joachimsthal mines and started isolating radium. In 1904 the Austrian government took over the ownership of the mines and stopped exporting raw ore. For some time the radium availability was low. The formation of an Austrian monopoly and the strong urge of other countries to have access to radium led to a world wide search for uranium ores. The United States took over as leading producer in the early 1910s. The Carnotite sands in Colorado provide some of the element, but richer ores are found in the Congo and the area of the Great Bear Lake and the Great Slave Lake of northwestern Canada. Radium can also be extracted from the waste from nuclear reactors. Large radium-containing uranium deposits are located in Russia, Canada (the Northwest Territories), the United States (New Mexico, Utah and Colorado, for example) and Australia. Neither of the deposits is mined for radium but the uranium content makes mining profitable. The amounts produced were always relatively small; for example, in 1918 13.6 g of radium were produced in the United States. As of 1954, the total worldwide supply of purified radium amounted to about 5 pounds (2.3 kg). History For more details on this topic, see Marie Curie#New elements.

Marie and Pierre Curie experimenting with radium, a drawing by André Castaigne Radium (Latin radius, ray) was discovered by Marie Skłodowska-Curie and her husband Pierre on December 21, 1898, in a uraninite sample. While studying the mineral, the Curies removed uranium from it and found that the remaining material was still radioactive. They then separated out a radioactive mixture consisting mostly of compounds of barium which gave a brilliant green flame color and crimson carmine spectral lines that had never been documented before. The Curies announced their discovery to the French Academy of Sciences on 26 December 1898. The naming of radium dates to about 1899, from the French word radium, formed in Modern Latin from radius (ray), called for its power of emitting energy in the form of rays. In 1910, radium was isolated as a pure metal by Curie and André-Louis Debierne through the

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electrolysis of a pure radium chloride solution by using a mercury cathode and distilling in an atmosphere of hydrogen gas. The same year, E. Eoler produced radium by heating its azide, Ra(N3)2. The Curies' new element was first industrially produced in the beginning of the 20th century by Biraco, a subsidiary company of Union MiniĂ¨re du Haut Katanga (UMHK) in its Olen plant in Belgium. UMHK offered to Marie Curie her first gram of radium. It gave historical names to the decay products of radium, such as radium A, B, C, etc., now known to be isotopes of other elements. On 4 February 1936, radium E (bismuth-210) became the first radioactive element to be made synthetically in the United States. Dr. John Jacob Livingood, at the radiation lab at University of California, Berkeley, was bombarding several elements with 5-MeV deuterons. He noted that irradiated bismuth emits fast electrons with a 5-day half-life, which matched the behavior of radium E. The common historical unit for radioactivity, the curie, is based on the radioactivity of

226

Ra.

Historical applications Some of the few practical uses of radium are derived from its radioactive properties. More recently discovered radioisotopes, such as 60Co and 137Cs, are replacing radium in even these limited uses because several of these isotopes are more powerful emitters, safer to handle, and available in more concentrated form. Luminescent paint Radium watch hands under ultraviolet light Radium was formerly used in self-luminous paints for watches, nuclear panels, aircraft switches, clocks, and instrument dials. A typical self-luminous watch that uses radium paint contains around 1 microgram of radium. In the mid-1920s, a lawsuit was filed against the United States Radium Corporation by five dying "Radium Girl" dial painters who had painted radium-based luminous paint on the dials of watches and clocks. The dial painters routinely licked their brushes to give them a fine point, thereby ingesting radium. Their exposure to radium caused serious health effects which included sores, anemia, and bone cancer. This is because radium is treated as calcium by the body, and deposited in the bones, where radioactivity degrades marrow and can mutate bone cells. During the litigation, it was determined that the company's scientists and management had taken considerable precautions to protect themselves from the effects of radiation, yet had not seen fit to protect their employees. Worse, for several years the companies had attempted to cover up the effects and avoid liability by insisting that the Radium Girls were

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instead suffering from syphilis. This complete disregard for employee welfare had a significant impact on the formulation of occupational disease labor law. As a result of the lawsuit, the adverse effects of radioactivity became widely known, and radium-dial painters were instructed in proper safety precautions and provided with protective gear. In particular, dial painters no longer licked paint brushes to shape them (which caused some ingestion of radium salts). Radium was still used in dials as late as the 1960s, but there were no further injuries to dial painters. This highlighted that the harm to the Radium Girls could easily have been avoided. From the 1960s the use of radium paint was discontinued. In many cases luminous dials were implemented with non-radioactive fluorescent materials excited by light; such devices glow in the dark after exposure to light, but the glow fades. Where indefinite self-luminosity in darkness was required, safer radioactive promethium paint was initially used, later replaced by tritium which continues to be used today. Tritium emits beta radiation which cannot penetrate the skin, rather than the penetrating gamma radiation of radium and is regarded as safer. It has a half-life of 12 years. Clocks, watches, and instruments dating from the first half of the 20th century, often in military applications, may have been painted with radioactive luminous paint. They are usually no longer luminous; however, this is not due to radioactive decay of the radium (which has a half-life of 1600 years) but to the fluorescence of the zinc sulfide fluorescent medium being worn out by the radiation from the radium. The appearance of an often thick layer of green or yellowish brown paint in devices from this period suggests a radioactive hazard. The radiation dose from an intact device is relatively low and usually not an acute risk; but the paint is dangerous if released and inhaled or ingested.

Recreational use Radium was once an additive in products such as toothpaste, hair creams, and even food items due to its supposed curative powers. Such products soon fell out of vogue and were prohibited by authorities in many countries after it was discovered they could have serious adverse health effects. (See, for instance, Radithor or Revigator types of "Radium water" or "Standard Radium Solution for Drinking".)Spas featuring radium-rich water are still occasionally touted as beneficial, such as those in Misasa, Tottori, Japan. In the U.S., nasal radium irradiation was also administered to children to prevent middle-ear problems or enlarged tonsils from the late 1940s through the early 1970s.

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Medical use Radium (usually in the form of radium chloride) was used in medicine to produce radon gas which in turn was used as a cancer treatment; for example, several of these radon sources were used in Canada in the 1920s and 1930s. The isotope 223Ra (under the trade name Xofigo) was approved by the FDA in 2013 for use in medicine as a cancer treatment of bone metastasis. Howard Atwood Kelly, one of the founding physicians of Johns Hopkins Hospital, was a major pioneer in the medical use of radium to treat cancer. His first patient was his own aunt in 1904, who died shortly after surgery. Kelly was known to use excessive amounts of radium to treat various cancers and tumors. As a result, some of his patients died from high amounts of radium exposure. His method of radium application was inserting a radium capsule near the affected area then sewing the radium "points" directly to the tumor. This was the same method used to treat Henrietta Lacks, the host of the original HeLa cells, for cervical cancer.

Research In 1909, the famous Rutherford experiment used radium as an alpha source to probe the atomic structure of gold. This experiment led to the Rutherford model of the atom and revolutionized the field of nuclear physics. When mixed with beryllium, it is a neutron source. This type of neutron source were for a long time the main source for neutrons in research. Precautions Radium is highly radioactive and its decay product, radon gas, is also radioactive. Since radium is chemically similar to calcium, it has the potential to cause great harm by replacing calcium in bones. Exposure to radium can cause cancer and other disorders, because radium and its decay product radon emit alpha particles upon their decay, which kill and mutate cells. At the time of the Manhattan Project in 1944, the "tolerance dose" for workers was set at 0.1 microgram of ingested radium. Some of the biological effects of radium were apparent from the start. The first case of socalled "radium-dermatitis" was reported in 1900, only 2 years after the element's discovery. The French physicist Antoine Becquerel carried a small ampoule of radium in his waistcoat pocket for 6 hours and reported that his skin became ulcerated. Marie Curie experimented with a tiny sample that she kept in contact with her skin for 10 hours, and noted that an ulcer appeared several days later. Handling of radium has been blamed for Curie's death due to aplastic anemia. Stored radium should be ventilated to prevent accumulation of radon. Emitted energy from the decay of radium also ionizes gases, fogs photographic emulsions, and produces many other detrimental effects.

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Further reading 

Albert Stwertka (1998). Guide to the Elements – Revised Edition. Oxford University Press. ISBN 0-19-508083-1.



Denise Grady (October 6, 1998). "A Glow in the Dark, and a Lesson in Scientific Peril". The New York Times. Retrieved 2007-12-25.



Nanny Fröman (1 December 1996). "Marie and Pierre Curie and the Discovery of Polonium and Radium". Nobel Foundation. Retrieved 2007-12-25.



Macklis, R. M. (1993). "The great radium scandal". Scientific American 269 (2): 94– 99. doi:10.1038/scientificamerican0893-94. PMID 8351514.



Clark, Claudia (1987). Radium Girls: Women and Industrial Health Reform, 1910– 1935. University of North Carolina Press. ISBN 0-8078-4640-6.

External links 

Lateral Science – Radium Discovery



Photos of Radium Water Bath in Oklahoma



NLM Hazardous Substances Databank – Radium, Radioactive



Reproduction of a 1942 comic book ad selling a "Radiumscope" to children



Annotated bibliography for radium from the Alsos Digital Library for Nuclear Issues



The Poisoner Next Door – Japan Today, 10/20/2001



Radium at The Periodic Table of Videos (University of Nottingham)

Text here. Article Sources and Contributors

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

Alkaline earth metal Source: https://en.wikipedia.org/w/index.php?oldid=592568139 Contributors:

1297,

Ahoerstemeier,

64.26.98.xxx,

Akjar13,

Alansohn,

777sms, Ale

jrb,

Aadal,

Aashaa,

Abshirdheere,

Andre

Engels,

AndyVolykhov,

Antony1103, Apollo, BD2412, Babychanel, Badocter, Bennybp, Bggoldie, Bkell, Bomac, Braincricket, Brion VIBBER, ChemNerd, Chris Dybala, Chris the speller, Chrislk02, Chrumps, Chukonu xbow, Chuunen Baka, Cimex, Cnilep, Conversion script, DMacks, DO11.10, Darekun, Darnir redhat, Darth Panda, DePiep, Dittaeva, Double

sharp,

Eddideigel,

Edgar181,

Element16,

Ellywa,

Eric119,

Espi,

EtymAesthete, Eugene van der Pijll, Eumolpo, Falkflyer, Fastily, Feline1, Femto, Fieldday-sunday, Firas@user, Firsfron, FocalPoint, Fonzy, Fred88freddy, GDonato, Gamma, Gentgeen, Giftlite, Glendoremus, GoingBatty, Graham87, Hadal, Hardy1Heron, Hatkill, Hellbus, Heron, Hugo-cs, Hydrogen Iodide, ITshnik, Icairns, IlGreven,

Izalithium,

J.delanoy,

JForget,

Jasper

Deng,

Jilkmarine,

Jimp,

Jjbernardiscool, Jojhutton, Jóna Þórunn, Karuna13, Kesac, Khazar, Khukri, Kingturtle, Kvnchris, Kyle1278, Lanthanum-138, Laurinavicius, LizardJr8, Longhair, Lop242438, MER-C, Mariagomez39, Mark grindell, Materialscientist, Mav, Mentifisto, Michael Hardy, MickiWiki, MikeLynch, Mikeo, Miquonranger03, Molestash, Monterey Bay, Mr Stephen, MusikAnimal, Myasuda, NaNi3320, Nergaal, Noctibus, Nonagonal Spider, Northumbrian, Omicron777, Paul-L, Pete Hobbs, Piast93, PlanetStar, Raymondwinn, Regibox, Remember, Res2216firestar, Rex the first, Rfc1394, Rjd0060, Rjwilmsi, Roadrunner, Robert Fraser, Robo37, Romanm, Rrv1988, Scerri, Scs, Sfan00 IMG, Shellreef, Shinglor, Skatebiker, Smack, Soliloquial, Some jerk on the Internet, Souravmohanty2005, Squids and Chips, Stone, StringTheory11, Suisui, Syd Henderson, Synergism, TCO, Tarquin, The Thing That Should Not Be, The User 567, Tide rolls, Tim1357, Timmeh, Titus III, Tom harrison, Tom.k, Trappist the monk, Versus22, Vsmith, Walkerma, Weasel the spriggan, Wiki alf, Wikipelli, Wyklety, Xerxes314, Xtv, Xxcandyy, आशीष भटनागर 

Electron

, 388 anonymous edits

configuration

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