alkaline earth metals

Alkaline Earth Metals

Lulu Press, Raleigh, N.C. USA

Dr. Pramod Kothari Assistant Professor, Department Of Chemistry

Government Post Graduate College, Berinag, District Pithoragarh

Uttarakhand (India)

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

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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

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While every attempt has been made to provide information that is both accurate and

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this information.

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

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 Pithoragarh

Uttarakhand (India)

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

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 Page 1

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

4 beryllium 2, 2 [He] 2s2

12 magnesium 2, 8, 2 [Ne] 3s2

20 calcium 2, 8, 8, 2 [Ar] 4s2

38 strontium 2, 8, 18, 8, 2 [Kr] 5s2

56 barium 2, 8, 18, 18, 8, 2 [Xe] 6s2

88 radium 2, 8, 18, 32, 18, 8, 2 [Rn] 7s2

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


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 M2+

HE

F-

HE

"MF2"

unit

HE

MF2

lattice

energies

Solubility

Be 2,455 458 3,371 3,526 soluble

Mg 1,922 458 2,838 2,978 0.0012

Ca 1,577 458 2,493 2,651 0.0002

Sr 1,415 458 2,331 2,513 0.0008

Ba 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.


Alkaline

earth

metal

Standard

atomic

weight

(u)

Melting

point

(K)

Melting

point

(C)

Boiling

point

(K)

Boiling

point

(C)

Density

(g/cm3)

Electronegativity

(Pauling)

First

ionization

energy

(kJmol1

)

Covalent

radius

(pm)

Flame test

color

Beryllium 9.012182(3) 1560 1287 2742 2469 1.85 1.57 899.5 105 White

Magnesium 24.3050(6) 923 650 1363 1090 1.738 1.31 737.7 150 Brilliant-

white

Calcium 40.078(4) 1115 842 1757 1484 1.54 1.00 589.8 180 Brick-

red

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 Apple-

green

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.


History

Etymology

The alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned 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 heatingproperties shared by these oxides. The realization that these earths were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Trait lmentaire 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 Whler and Antoine Bussy in 1828 by reacting beryllium compounds with potassium. In 1910, radium was isolated as a pure metal by Curie and Andr-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 Whler 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


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. Davys 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 Jns 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.


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-most-abundant 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.


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


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.


Bibliography

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


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. For example, the electron configuration of the neon atom is 1s

2 2s

2 2p

6.

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


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 (=0) p (=1)

m=0 m=0 m=1

s pz px py

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 electron shell can accommodate 2n

2 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 spineach 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, , within a shell. The values = 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(2 + 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.


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 written 1s

1. Lithium has two electrons in the 1s-subshell and one in the (higher-energy) 2s-

subshell, so its configuration is written 1s2 2s

1 (pronounced "one-s-two, two-s-one").

Phosphorus (atomic number 15) is as follows: 1s2 2s

2 2p

6 3s

2 3p

3.

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 the noble gases. Phosphorus, for instance, differs from neon (1s

2 2s

2 2p

6) only by the

presence of a third shell. Thus, the electron configuration of neon is pulled out, and phosphorus is written as follows: [Ne] 3s

2 3p

3. 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 Madelung's rule. Hence the electron configuration of iron can be written as [Ar] 3d

6 4s

2

(keeping the 3d-electrons with the 3s- and 3p-electrons which are implied by the configuration of argon) or as [Ar] 4s

2 3d

6 (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 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 electron-electron interactions. The configuration that corresponds to the lowest electronic energy is called the ground state. Any other configuration is an excited state.

As an example, the ground state configuration of the sodium atom is 1s22s

22p

63s, as

deduced from the Aufbau principle (see below). The first excited state is obtained by promoting a 3s electron to the 3p orbital, to obtain the 1s

22s

22p

63p 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 1s

22s

22p

53s

2 configuration.


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 present-day chemist: sulfur was given as 2.4.4.6 instead of 1s

2 2s

2 2p

6 3s

2 3p

4 (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 Schrdinger 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 higher-

energy 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.)


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 [E] ns

2 (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.


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 HartreeFock 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 nave 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 periodic table before the transition metals, and have electron configurations [Ar] 4s

1 and

[Ar] 4s2 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 (n = 3, l = 2). However, chromium and copper have electron configurations [Ar] 3d

5 4s

1 and

[Ar] 3d10

4s1 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 3d

6, 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 doesn't. There is no special reason why the Fe

2+ 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 HartreeFock calculations,


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 Period 5 Period 6 Period 7

Element Z Electron

Configuration

Element Z Electron

Configuration

Element Z Electron

Configuration

Element Z Electron

Configuration

Lanthanum 57 [Xe] 6s25d

1 Actinium 89 [Rn] 7s

26d

1

Cerium 58 [Xe] 6s24f

1 5d

1 Thorium 90 [Rn] 7s

26d

2

Praseodymium 59 [Xe] 6s2 4f

3 Protactinium 91 [Rn] 7s

25f

2 6d

1

Neodymium 60 [Xe] 6s2 4f

4 Uranium 92 [Rn] 7s

25f

3 6d

1

Promethium 61 [Xe] 6s2 4f

5 Neptunium 93 [Rn] 7s

25f

4 6d

1

Samarium 62 [Xe] 6s2 4f

6 Plutonium 94 [Rn] 7s

2 5f

6

Europium 63 [Xe] 6s2 4f

7 Americium 95 [Rn] 7s

2 5f

7

Gadolinium 64 [Xe] 6s24f

7 5d

1 Curium 96 [Rn] 7s

25f

7 6d

1

Terbium 65 [Xe] 6s2 4f

9 Berkelium 97 [Rn] 7s

2 5f

9

Scandium 21 [Ar] 4s2 3d

1 Yttrium 39 [Kr] 5s

2 4d

1 Lutetium 71 [Xe] 6s

2 4f

14

5d1

Lawrencium 103 [Rn] 7s2

5f14

7p1

Titanium 22 [Ar] 4s2 3d

2 Zirconium 40 [Kr] 5s

2 4d

2 Hafnium 72 [Xe] 6s

2 4f

14

5d2

Rutherfordium 104 [Rn] 7s2 5f

14

6d2

Vanadium 23 [Ar] 4s2 3d

3 Niobium 41 [Kr] 5s

1 4d

4 Tantalum 73 [Xe] 6s

2 4f

14

5d3

Chromium 24 [Ar] 4s1 3d

5 Molybdenum 42 [Kr] 5s

1 4d

5 Tungsten 74 [Xe] 6s

2 4f

14

5d4

Manganese 25 [Ar] 4s2 3d

5 Technetium 43 [Kr] 5s

2 4d

5 Rhenium 75

[Xe] 6s2 4f

14

5d5


Iron 26 [Ar] 4s2 3d

6 Ruthenium 44 [Kr] 5s

1 4d

7 Osmium 76

[Xe] 6s2 4f

14

5d6

Cobalt 27 [Ar] 4s2 3d

7 Rhodium 45 [Kr] 5s

1 4d

8 Iridium 77

[Xe] 6s2 4f

14

5d7

Nickel 28

[Ar] 4s2 3d

8 or

[Ar] 4s1 3d

9

(disputed)

Palladium 46 [Kr] 4d10

Platinum 78 [Xe] 6s

1

4f14

5d9

Copper 29 [Ar] 4s1 3d

10 Silver 47 [Kr] 5s

1 4d

10 Gold 79

[Xe] 6s1

4f14

5d10

Zinc 30 [Ar] 4s2 3d

10 Cadmium 48 [Kr] 5s

2 4d

10 Mercury 80

[Xe] 6s2 4f

14

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 the dioxygen molecule, O2, is 1g

2 1u

2 2g

2 2u

2 1u

4 3g

2 1g

2. The term 1g

2 represents

the two electrons in the two degenerate *-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 rationalization of chemical properties, in both inorganic and organic chemistry. In effect, electron configurations, along with some simplified form of molecular orbital theory, have become the modern equivalent of the valence concept, describing the number and type of chemical bonds that an atom can be expected to form.

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


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: McGraw-Hill. pp. 123. ISBN 0-07-112651-1.


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 close-

packed 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

(1925 Jkg1

K1

) and thermal conductivity (216 Wm1

K1

), 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 (11.4106

K1

), these characteristics

result in a unique stability under conditions of thermal loading.


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 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

interstellar gas and dust. Primordial beryllium contains only one stable isotope, 9Be, and

therefore beryllium is a monoisotopic element.

Plot showing variations in solar activity, including variation in sunspot number (red) and 10

Be

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

indicate lower 10

Be levels


Radioactive cosmogenic 10

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

spallation of oxygen.10

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, 10

Be

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 10

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 10

Be by the

reaction of fast neutrons with 13

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

past activity at nuclear weapon test sites. The isotope 7Be (half-life 53 days) is also

cosmogenic, and shows an atmospheric abundance linked to sunspots, much like 10

Be.

8Be has a very short half-life of about 710

17 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

phase to produce carbon by the fusion of 4He nuclei and the very low concentrations of

available beryllium-8. The British astronomer Sir Fred Hoyle first showed that the energy

levels of 8Be and

12C 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

7Be 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 13

Be which decays through neutron

emission. It has a half-life of 2.7 1021

s. 6Be is also very short-lived with a half-life of

5.0 1021

s. The exotic isotopes 11

Be and 14

Be are known to exhibit a nuclear halo. This

phenomenon can be understood as the nuclei of 11

Be and 14

Be have, respectively, 1 and 4

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


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

the soils, 6 ppm, and is found in 0.2 parts per trillion (ppt) of sea water. Trace amounts of 9Be

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.


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 is added to beryllium hydroxide to yield a precipitate of ammonium

tetrafluoroberyllate, which is heated to 1,000C 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.


Beryllium hydrolysis as a function of pH

Water molecules attached to Be are omitted

The beryllium atom has the electronic configuration [He] 2s2. 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 thus absorbing other cations such as Al3+

which

might interfere for example, in the solvent extraction of a complex formed between Be2+

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+

ion.

[Be(H2O)4]2+

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

+

Other products of hydrolysis include the trimeric ion [Be3(OH)3(H2O)6]3+

. 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


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). BeF2 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

structure adopted by calcium

difluoride and many other difluorides

Disordered structure of beryllium

glass (sketch, two dimensions)

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


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 Ren

Just Hay 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 Whler in 1828.


Friedrich Whler was one of the men who independently isolated beryllium

Friedrich Whler 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, Whler 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 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


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 Bry; Greek , brullos, beryl; Prakrit veruliya (); Pli veuriya (),

veiru () or viar () "to become pale", in reference to the pale semiprecious

gemstone beryl. The original source is probably the Sanskrit word 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: , 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.


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


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 liquid-

fuel 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.


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

and remains more uniform 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-


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, radium-

226, 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


shape due to its brittleness, and toxic if mishandled, beryllium tweeters are limited to high-

end 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). Cross-

rolled beryllium sheet is an excellent structural support for printed circuit boards in surface-

mount 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 high-

power 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


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

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 AZ Guide to the Elements.

Oxford, England, UK: Oxford University Press. ISBN 0-19-850340-7.


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-most-common 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 or 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.

alkaline earth metals

Documents

reactive metals

group of chemical elements

periodic table

similar properties

outermost electrons

standard temperature

oxidation state

calcium ca