# Alkaline earth metal

The alkaline earth metals are six chemical elements in group 2 of the periodic table. They are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra).[1] The elements have very similar properties: they are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure.[2]

Structurally, they have in common an outer s- electron shell which is full;[2][3][4] that is, this orbital contains its full complement of two electrons, which these elements readily lose to form cations with charge +2, and an oxidation state of +2.[5]

All the discovered alkaline earth metals occur in nature, although radium occurs only through the decay chain of uranium and thorium and not as a primordial element.[6] Experiments have been conducted to attempt the synthesis of element 120, the next potential member of the group, but they have all met with failure.

Alkaline earth metals
 Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
IUPAC group number 2
Name by element beryllium group
Trivial name alkaline earth metals
CAS group number
(US, pattern A-B-A)
IIA
old IUPAC number
(Europe, pattern A-B)
IIA

↓ Period
2
Beryllium (Be)
4
3
Magnesium (Mg)
12
4
Calcium (Ca)
20
5
Strontium (Sr)
38
6
Barium (Ba)
56
7
88

Legend
 primordial element element by radioactive decay Atomic number color:black=solid

## 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[n 1]
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;[2] 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 earth metals react with the halogens to form the alkaline earth metal halides, all of which are 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.[2] The alkaline earth metals have the second-lowest first ionization energies in their respective periods of the periodic table[4] 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.[2][4]

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.[7] Even the compound beryllium fluoride, which is the most ionic beryllium compound, has a low melting point and a low electrical conductivity when melted.[8][9][10]

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[n 2]
Metal
M2+
HE
[11]
F
HE
[12]
"MF2"
unit
HE
MF2
lattice
energies
[13]
Solubility
[14]
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

#### Carbonyl complexes

Calcium, strontium, and barium all form octacarbonyl complexes and an 18-electron valence shell. The compounds were characterized in frozen neon matrices by vibrational spectroscopy and in gas phase by mass spectrometry. Analysis of the electronic structure reveals that not only barium but also strontium and calcium may effectively use their (n − 1)d atomic orbitals in chemical bonding, suggesting they behave like transition metals.[15]

### 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)[n 3][17][18]
Melting point
(K)
Melting point
(°C)
Boiling point
(K)[4]
Boiling point
(°C)[4]
Density
(g/cm3)
Electronegativity
(Pauling)
First ionization energy
(kJ·mol−1)
(pm)[19]
Flame test color
Beryllium 9.012182(3) 1560 1287 2742 2469 1.85 1.57 899.5 105 White[20]
Magnesium 24.3050(6) 923 650 1363 1090 1.738 1.31 737.7 150 Brilliant-white[2]
Calcium 40.078(4) 1115 842 1757 1484 1.54 1.00 589.8 180 Brick-red[2]
Strontium 87.62(1) 1050 777 1655 1382 2.64 0.95 549.5 200 Crimson[2]
Barium 137.327(7) 1000 727 2170 1897 3.594 0.89 502.9 215 Apple-green[2]
Radium [226][n 4] 973 700 2010 1737 5.5 0.9 509.3 221 Crimson red[n 5]

## 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 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,[25] 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.[26] The same is true for the beryllium compounds beryl and emerald.[27] 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.[28][29][30]

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.[28][29][30]

#### Beryllium

Emerald is a form of beryl, the principal mineral of beryllium.

Beryl, a mineral that contains beryllium, has been known since the time of the Ptolemaic Kingdom in Egypt.[27] Although it was originally thought that beryl was an aluminium silicate,[31] beryl was later found to contain a then-unknown element when, in 1797, Louis-Nicolas Vauquelin dissolved aluminium hydroxide from beryl in an alkali.[32] In 1828, Friedrich Wöhler[33] and Antoine Bussy[34] 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.[35] 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.[35]

#### Magnesium

Magnesium was first produced by Sir Humphry Davy in England in 1808 using electrolysis of a mixture of magnesia and mercuric oxide.[36] Antoine Bussy prepared it in coherent form in 1831. Davy's first suggestion for a name was magnium,[36] but the name magnesium is now used.

#### Calcium

Lime has been used as a material for building since 7000 to 14,000 BCE,[26] and kilns used for lime have been dated to 2,500 BCE in Khafaja, Mesopotamia.[37][38] 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,[39] after hearing that Jöns Jakob Berzelius had prepared a calcium amalgam from the electrolysis of lime in mercury.

#### Strontium

In 1790, physician Adair Crawford discovered ores with distinctive properties, which were named strontites in 1793 by Thomas Charles Hope, a chemistry professor at the University of Glasgow,[40] 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.[41]

#### 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.[42][43]

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.[44] Radium was named in 1899 from the word radius, meaning ray, as radium emitted power in the form of rays.[45]

## Occurrence

Series of alkaline earth metals.

Beryllium occurs in the earth's crust at a concentration of two to six parts per million (ppm),[46] 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.[47][48] However, in freshwater, beryllium is somewhat more common, with a concentration of 0.1 parts per billion.[49]

Magnesium and calcium are very common in the earth's crust, being respectively the fifth- eighth-most-abundant elements. None of the alkaline earth metals are found in their elemental state. Common magnesium—containing minerals are carnallite, magnesite, and dolomite. Common calcium-containing minerals are chalk, limestone, gypsum, and anhydrite.[2]

Strontium is the fifteenth-most-abundant element in the Earth's crust. The principal minerals are celestite and strontianite.[50] Barium is slightly less common, much of it in the mineral barite.[51]

Radium, being a decay product of uranium, is found in all uranium-bearing ores.[52] Due to its relatively short half-life,[53] radium from the Earth's early history has decayed, and present-day samples have all come from the much slower decay of uranium.[52]

## 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.[7]

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.[54]

To produce barium, barite is subjected to barium sulfide by carbothermic reduction. 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.[51] The most important supplier of barium is China, which produces more than 50% of world supply.[55]

## Applications

Beryllium is used mostly for military applications,[56] but there are other uses of beryllium, as well. In electronics, beryllium is used as a p-type dopant in some semiconductors,[57] and beryllium oxide is used as a high-strength electrical insulator and heat conductor.[58] 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.[59][60]

Magnesium has many uses. It offers advantages over other materials such as aluminium, although this usage has fallen out of favor due to magnesium's flammability.[61] Magnesium is also often alloyed with aluminium or zinc to form materials with more desirable properties than any pure metal.[62] 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.[63]

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.[64]

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,[65] and pure strontium is used in the study of neurotransmitter release in neurons.[66][67] Radioactive strontium-90 finds some use in RTGs,[68][69] which utilize its decay heat. Barium has some use in vacuum tubes to remove gases,[51] and barium sulfate has many uses in the petroleum industry,[4] as well as other industries.[4][51][70]

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,[71] although this use was stopped after workers got sick.[72] 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.[61] Radium is no longer even used for its radioactive properties, as there are more powerful and safer emitters than radium.[73][74]

## Representative reactions of alkaline earth metals

Reaction with halogens

Ca + Cl2 → CaCl2

Anhydrous calcium chloride is a hygroscopic substance that is used as a desiccant. Exposed to air, it will absorb water vapour from the air, forming a solution. This property is known as deliquescence.

Reaction with oxygen

Ca + 1/2O2 → CaO
Mg + 1/2O2 → MgO

Reaction with sulphur

Ca + 1/8S8 → CaS

Reaction with carbon

With carbon, they form acetylides indirectly. Beryllium forms carbide.

2Be + C → Be2C
CaO + 3C → CaC2 + CO (at 25000C in furnace)
CaC2 + 2H2O → Ca(OH)2 + C2H2
Mg2C3 + 4H2O → 2Mg(OH)2 + C3H4

Reaction with nitrogen

Only Be and Mg form nitrides directly.

3Be + N2 → Be3N2
3Mg + N2 → Mg3N2

Reaction with hydrogen

Alkaline earth metals react with hydrogen to generate saline hydride that are unstable in water.

Ca + H2 → CaH2

Reaction with water

Ca, Sr and Ba readily react with water to form hydroxide and hydrogen gas. Be and Mg are passivated by an impervious layer of oxide. However, amalgamated magnesium will react with water vapour.

Mg + H2O → MgO + H2

Reaction with acidic oxides

Alkaline earth metals reduce the nonmetal from its oxide.

2Mg + SiO2 → 2MgO + Si
2Mg + CO2 → 2MgO + C (in solid carbon dioxide)

Reaction with acids

Mg + 2HCl → MgCl2 + H2
Be + 2HCl → BeCl2 + H2

Reaction with bases

Be exhibits amphoteric properties. It dissolves in concentrated sodium hydroxide.

Be + NaOH + 2H2O → Na[Be(OH)3] + H2

Reaction with alkyl halides

Magnesium reacts with alkyl halides via an insertion reaction to generate Grignard reagents.

RX + Mg → RMgX (in anhydrous ether)

## Identification of alkaline earth cations

The flame test

The table below[75] presents the colours observed when the flame of a Bunsen burner is exposed to salts of alkaline earth metals. Be and Mg do not impart colour to the flame due to their small size.[76]

Metal Colour
Ca Brick-red
Sr Crimson red
Ba Green/Yellow
Ra Carmine red

In solution

Mg2+

Disodium phosphate is a very selective reagent for magnesium ions and, in the presence of ammonium salts and ammonia, forms a white precipitate of ammonium magnesium phosphate.

Mg2+ + NH3 + Na2HPO4 → (NH4)MgPO4 + 2Na+

Ca2+

Ca2+ forms a white precipitate with ammonium oxalate. Calcium oxalate is insoluble in water, but is soluble in mineral acids.

Ca2+ + (COO)2(NH4)2 → (COO)2Ca + NH4+

Sr2+

Strontium ions precipitate with soluble sulphate salts.

Sr2+ + Na2SO4 → SrSO4 + 2Na+

All ions of alkaline earth metals form white precipitate with ammonium carbonate in the presence of ammonium chloride and ammonia.

## Compounds of alkaline earth metals

Oxides

The alkaline earth metal oxides are formed from the thermal decomposition of the corresponding carbonates.

CaCO3 → CaO + CO2 (at approx. 9000C)

In laboratory, they are obtained from hydroxides:

Mg(OH)2 → MgO + H2O

or nitrates:

Ca(NO3)2 → CaO + 2NO + 1/2O2

The oxides exhibit basic character: they turn phenolphthalein red and litmus, blue. They react with water to form hydroxides in an exothermic reaction.

CaO + H2O → Ca(OH)2 + Q

Calcium oxide reacts with carbon to form acetylide.

CaO + 3C → CaC2 + CO (at 25000)
CaC2 + N2 → CaCN2 + C
CaCN2 + H2SO4 → CaSO4 + H2N—CN
H2N—CN + H2O → (H2N)CO (urea)
CaCN2 + 2H2O → CaCO3 + NH3

Hydroxides

They are generated from the corresponding oxides on reaction with water. They exhibit basic character: they turn phenolphthalein pink and litmus, blue. Beryllium hydroxide is an exception as it exhibits amphoteric character.

Be(OH)2 + 2HCl → BeCl2 + H2O
Be(OH)2 + NaOH → Na[Be(OH)3]

Salts

Ca and Mg are found in nature in many compounds such as dolomite, aragonite, magnesite (carbonate rocks). Calcium and magnesium ions are found in hard water. Hard water represents a multifold issue. It is of great interest to remove these ions, thus softening the water. This procedure can be done using reagents such as calcium hydroxide, sodium carbonate or sodium phosphate. A more common method is to use ion-exchange aluminosilicates or ion-exchange resins that trap Ca2+ and Mg2+ and liberate Na+ instead:

Na2O·Al2O3·6SiO2 + Ca2+ → CaO·Al2O3·6SiO2 + 2Na+

## 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.[7] 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.[77] 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.[78] 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.[79]

The chemistry of element 120 is predicted to be closer to that of calcium or strontium[80] 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.[80]

## Notes

1. ^ Noble gas notation is used for conciseness; the nearest noble gas that precedes the element in question is written first, and then the electron configuration is continued from that point forward.
2. ^ Energies are given in −kJ/mol, solubilities in mol/L; HE means "hydration energy".
3. ^ The number given in parentheses refers to the measurement uncertainty. This uncertainty applies to the least significant figure(s) of the number prior to the parenthesized value (i.e., counting from rightmost digit to left). For instance, 1.00794(7) stands for 1.00794±0.00007, whereas 1.00794(72) stands for 1.00794±0.00072.[16]
4. ^ The element does not have any stable nuclides, and a value in brackets indicates the mass number of the longest-lived isotope of the element.[17][18]
5. ^ The color of the flame test of pure radium has never been observed; the crimson-red color is an extrapolation from the flame test color of its compounds.[21]

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## Bibliography

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

• Group 2 – Alkaline Earth Metals, Royal Chemistry Society.
• Hogan, C.Michael. 2010. Calcium. eds. A.Jorgensen, C. Cleveland. Encyclopedia of Earth. National Council for Science and the Environment.
• Maguire, Michael E. "Alkaline Earth Metals." Chemistry: Foundations and Applications. Ed. J. J. Lagowski. Vol. 1. New York: Macmillan Reference USA, 2004. 33–34. 4 vols. Gale Virtual Reference Library. Thomson Gale.
• Silberberg, M.S., Chemistry: The molecular nature of Matter and Change (3e édition, McGraw-Hill 2009)
• Petrucci R.H., Harwood W.S. et Herring F.G., General Chemistry (8e édition, Prentice-Hall 2002)
122 iron arsenide

The 122 iron arsenide unconventional superconductors are part of a new class of iron-based superconductors. They form in the tetragonal I4/mmm, ThCr2Si2 type, crystal structure. The shorthand name "122" comes from their stoichiometry; the 122s have the chemical formula AEFe2Pn2, where AE stands for alkaline earth metal (Ca, Ba, Sr or Eu) and Pn is pnictide (As, P, etc.). These materials become superconducting under pressure and also upon doping. The maximum superconducting transition temperature found to date is 38 K in the Ba0.6K0.4Fe2As2. The microscopic description of superconductivity in the 122s is yet unclear.

Alkali

In chemistry, an alkali (; from Arabic: al-qaly "ashes of the saltwort") is a basic, ionic salt of an alkali metal or alkaline earth metal chemical element. An alkali also can be defined as a base that dissolves in water. A solution of a soluble base has a pH greater than 7.0. The adjective alkaline is commonly, and alkalescent less often, used in English as a synonym for basic, especially for bases soluble in water. This broad use of the term is likely to have come about because alkalis were the first bases known to obey the Arrhenius definition of a base, and they are still among the most common bases.

Alkali salt

Alkali salts or basic salts are salts that are the product of the neutralization of a strong base and a weak acid.

Rather than being neutral (as some other salts), alkali salts are bases as their name suggests. What makes these compounds basic is that the conjugate base from the weak acid hydrolyzes to form a basic solution. In sodium carbonate, for example, the carbonate from the carbonic acid hydrolyzes to form a basic solution. The chloride from the hydrochloric acid in sodium chloride does not hydrolyze, though, so sodium chloride is not basic.

The difference between a basic salt and an alkali is that an alkali is the soluble hydroxide compound of an alkali metal or an alkaline earth metal. A basic salt is any salt that hydrolyzes to form a basic solution.

Another definition of a basic salt would be a salt that contains amounts of both hydroxide and other anions. White lead is an example. It is basic lead carbonate, or lead carbonate hydroxide.

These materials are known for their high levels of dissolution in polar solvents.

These salts are insoluble and are obtained through precipitation reactions.

Barium

Barium is a chemical element with symbol Ba and atomic number 56. It is the fifth element in group 2 and is a soft, silvery alkaline earth metal. Because of its high chemical reactivity, barium is never found in nature as a free element. Its hydroxide, known in pre-modern times as baryta, does not occur as a mineral, but can be prepared by heating barium carbonate.

The most common naturally occurring minerals of barium are barite (now called baryte) (barium sulfate, BaSO4) and witherite (barium carbonate, BaCO3), both insoluble in water. The name barium originates from the alchemical derivative "baryta", from Greek βαρύς (barys), meaning "heavy." Baric is the adjectival form of barium. Barium was identified as a new element in 1774, but not reduced to a metal until 1808 with the advent of electrolysis.

Barium has few industrial applications. Historically, it was used as a getter for vacuum tubes and in oxide form as the emissive coating on indirectly heated cathodes. 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. Barium compounds are added to fireworks to impart a green color. Barium sulfate is used as an insoluble additive to oil well drilling fluid, as well as in a purer form, as X-ray radiocontrast agents for imaging the human gastrointestinal tract. The soluble barium ion and soluble compounds are poisonous, and have been used as rodenticides.

Beauvericin

Beauvericin is a depsipeptide with antibiotic and insecticidal effects belonging to the enniatin family. It was isolated from the fungus Beauveria bassiana, but is also produced by several other fungi, including several Fusarium species; it may therefore occur in grain (such as corn, wheat and barley) contaminated with these fungi. Beauvericin is active against Gram-positive bacteria and mycobacteria, and is also capable of inducing programmed cell death in mammals.Chemically, beauvericin is a cyclic hexadepsipeptide with alternating N-methyl-phenylalanyl and D-hydroxy-iso-valeryl residues. Its ion-complexing capability allows beauvericin to transport alkaline earth metal and alkali metal ions across cell membranes.

Beauvericin has in vitro fungicidal effects on Candida parapsilosis when used in combination with the antifungal drug ketoconazole at dosages of 0.1 μg/ml. Increased survivability rates and low cytotoxicity were also observed in mouse models.

Calcium

Calcium is a chemical element with symbol Ca and atomic number 20. As an alkaline earth metal, calcium is a reactive metal that forms a dark oxide-nitride layer when exposed to air. Its physical and chemical properties are most similar to its heavier homologues strontium and barium. It is the fifth most abundant element in Earth's crust and the third most abundant metal, after iron and aluminium. The most common calcium compound on Earth is calcium carbonate, found in limestone and the fossilised remnants of early sea life; gypsum, anhydrite, fluorite, and apatite are also sources of calcium. The name derives from Latin calx "lime", which was obtained from heating limestone.

Some calcium compounds were known to the ancients, though their chemistry was unknown until the seventeenth century. Pure calcium was isolated in 1808 via electrolysis of its oxide by Humphry Davy, who named the element. Calcium compounds are widely used in many industries: in foods and pharmaceuticals for calcium supplementation, in the paper industry as bleaches, as components in cement and electrical insulators, and in the manufacture of soaps. On the other hand, the metal in pure form has few applications due to its high reactivity; still, in small quantities it is often used as an alloying component in steelmaking, and sometimes, as a calcium–lead alloy, in making automotive batteries.

Calcium is the most abundant metal and the fifth-most abundant element in the human body. As electrolytes, calcium ions play a vital role in the physiological and biochemical processes of organisms and cells: in signal transduction pathways where they act as a second messenger; in neurotransmitter release from neurons; in contraction of all muscle cell types; as cofactors in many enzymes; and in fertilization. Calcium ions outside cells are important for maintaining the potential difference across excitable cell membranes as well as proper bone formation.

Difluoride

Difluorides are chemical compounds with two fluorine atoms per molecule (or per formula unit).

Metal difluorides are all ionic. Despite being highly ionic, the alkali earth metal difluorides generally have extremely high lattice stability and are thus insoluble in water. One exception is beryllium difluoride. In addition, many transition metal difluorides are water-soluble.

Calcium difluoride is a notable compound. In the form of the mineral fluorite it is the major source of commercial fluorine. It also has an epynomic crystal structure, which is an end member of the spectrum starting from bixbyite and progressing through pyrochlore.

Isotopes of strontium

The alkaline earth metal strontium (38Sr) has four stable, naturally occurring isotopes: 84Sr (0.56%), 86Sr (9.86%), 87Sr (7.0%) and 88Sr (82.58%). Its standard atomic weight is 87.62(1).

Only 87Sr is radiogenic; it is produced by decay from the radioactive alkali metal 87Rb, which has a half-life of 4.88 × 1010 years (i.e. more than three times longer than the current age of the universe). Thus, there are two sources of 87Sr in any material: primordial, formed during nucleosynthesis along with 84Sr, 86Sr and 88Sr; and that formed by radioactive decay of 87Rb. The ratio 87Sr/86Sr 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 electron configuration similar to that of calcium, it readily substitutes for Ca in minerals.

In addition to the four stable isotopes, thirty-one unstable isotopes of strontium are known to exist (see Table, below): the longest-lived of these are 90Sr with a half-life of 28.9 years and 85Sr with a half-life of 64.853 days. Of importance are strontium-89 (89Sr) with a half-life of 50.57 days, and strontium-90 (90Sr). They decay by emitting an electron and an antineutrino (${\displaystyle {\ce {{\bar {\nu }}_{e}}}}$) in beta decay (β decay) to become yttrium:

${\displaystyle {\ce {^{89}_{38}Sr->{^{89}_{39}Y}+{e^{-}}+{\bar {\nu }}_{e}}}}$
${\displaystyle {\ce {^{90}_{38}Sr->{^{90}_{39}Y}+{e^{-}}+{\bar {\nu }}_{e}}}}$

89Sr is an artificial radioisotope used in treatment of bone cancer. In circumstances where cancer patients have widespread and painful bony metastases, the administration of 89Sr results in the delivery of beta particles directly to the area of bony problem,[further explanation needed] where calcium turnover is greatest.

90Sr is a by-product of nuclear fission, present in nuclear fallout. The 1986 Chernobyl nuclear accident contaminated a vast area with 90Sr. It causes health problems, as it substitutes for calcium in bone, preventing expulsion from the body. Because it is a long-lived high-energy beta emitter, it 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 lightest isotope is 73Sr and the heaviest is 107Sr.

All other strontium isotopes have half-lives shorter than 55 days, most under 100 minutes.

List of UN numbers 1301 to 1400

The UN numbers from UN1301 to UN1400 as assigned by the United Nations Committee of Experts on the Transport of Dangerous Goods.

Magnesite in Greece

Magnesium (Mg) is a chemical element, an alkaline earth metal, the eighth-most abundant element in the Earth's crust and the fourth-most common element on Earth. It is contained in magnesite, dolomite, brucite, carnallite, talc, and magnesium minerals. China is now the biggest producer of crude magnesite in the world for refractory and agricultural uses.

Crude magnesite was produced in Greece in 1910. It was first found in Atalanti and in the Province of Lokris, central Greece.

Other localities were Perachori, near Corinth; Ermioni (or Kastri) and on Spetses Island in southern Argolis in the Peloponnese; on Paros Island (Cyclades); around Thebes in Boeotia and in Papades and Troupi in northern Euboea. Euboea was mostly worked until the 1980s.

Galataki (near Limni) and Afrati (near Chalcis) were exploited by the English company Petrified Ltd. (founded 1897 and based in London), which finally sold its assets to the Anglo-Greek Magnesite Company. Ltd (AGM) in 1902. Besides, northern Greece was found to have magnesite mining interest, in the concessions of Aghia Paraskevi (east of Thessaloniki) in small production, and in Chalkidiki’s concessions of Vavdos,

Patelidas

and Yerakini with the largest deposits.

N-Methyltaurine

N-Methyltaurine (2-methylaminoethanesulfonic acid) is an aminosulfonic acid which is present as a zwitterion in the crystalline state and in polar solvents (just like amino acids). In contrast to the widespread taurine, N-methyltaurine has been found in nature only in red algae, where it is formed by methylation of taurine. It is suitable for esterification (actually amide formation) with long-chain carboxylic acids to taurides (acylaminoethansulfonaten) because of its high polarity and the relatively good solubility of its alkaline earth metal salts, which are also used as mild anionic surfactants.

Period (periodic table)

A period in the periodic table is a row of chemical elements. All elements in a row have the same number of electron shells. Each next element in a period has one more proton and is less metallic than its predecessor. Arranged this way, groups of elements in the same column have similar chemical and physical properties, reflecting the periodic law. For example, the alkali metals lie in the first column (group 1) and share similar properties, such as high reactivity and the tendency to lose one electron to arrive at a noble-gas electronic configuration. As of 2016, a total of 118 elements have been discovered and confirmed.

Modern quantum mechanics explains these periodic trends in properties in terms of electron shells. As atomic number increases, shells fill with electrons in approximately the order shown at right. The filling of each shell corresponds to a row in the table.

In the s-block and p-block of the periodic table, elements within the same period generally do not exhibit trends and similarities in properties (vertical trends down groups are more significant). However, in the d-block, trends across periods become significant, and in the f-block elements show a high degree of similarity across periods.

Period 3 element

A period 3 element is one of the chemical elements in the third row (or period) of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when the periodic table skips a row and a chemical behaviour begins to repeat, meaning that elements with similar behavior fall into the same vertical columns. The third period contains eight elements: sodium, magnesium, aluminium, silicon, phosphorus, sulfur, chlorine, and argon. The first two, sodium and magnesium, are members of the s-block of the periodic table, while the others are members of the p-block. Note that there is a 3d subshell, but it is not filled until period 4, such giving the period table its characteristic shape of "two rows at a time". All of the period 3 elements occur in nature and have at least one stable isotope.

Period 4 element

A period 4 element is one of the chemical elements in the fourth row (or period) of the periodic table of the elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The fourth period contains 18 elements, beginning with potassium and ending with krypton. As a rule, period 4 elements fill their 4s shells first, then their 3d and 4p shells, in that order; however, there are exceptions, such as chromium.

Period 7 element

A period 7 element is one of the chemical elements in the seventh row (or period) of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The seventh period contains 32 elements, tied for the most with period 6, beginning with francium and ending with oganesson, the heaviest element currently discovered. As a rule, period 7 elements fill their 7s shells first, then their 5f, 6d, and 7p shells, in that order; however, there are exceptions, such as plutonium.

Periodic table (detailed cells)

The periodic table is a tabular method of displaying the chemical elements. It can show much information, after name, symbol and atomic number. Also, for each element mean atomic mass value for the natural isotopic composition of each element can be noted.

The two layout forms originate from two graphic forms of presentation of the same periodic table. Historically, when the f-block was identified it was drawn below the existing table, with markings for its in-table location (this page uses dots or asterisks). Also, a common presentation is to put all 15 lanthanide and actinide columns below, while the f-block only has 14 columns. One lanthanide and actinide each are d-block elements, belonging to group 3 with scandium and yttrium, though whether these are the first of each series (lanthanum and actinium) or the last (lutetium and lawrencium) has been disputed. The tables below show lanthanum and actinium as group 3 elements, as this is the more common form in the literature.

Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869. Mendeleev invented the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table has been refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain chemical behavior.

Strontium

Strontium is the chemical element with symbol Sr and atomic number 38. An alkaline earth metal, strontium is a soft silver-white yellowish metallic element that is highly chemically reactive. The metal forms a dark oxide layer when it is exposed to air. Strontium has physical and chemical properties similar to those of its two vertical neighbors in the periodic table, calcium and barium. It occurs naturally mainly in the minerals celestine and strontianite, and is mostly mined from these. While natural strontium is stable, the synthetic 90Sr isotope is radioactive and is one of the most dangerous components of nuclear fallout, as strontium is absorbed by the body in a similar manner to calcium. Natural stable strontium, on the other hand, is not hazardous to health.

Both strontium and strontianite are named after Strontian, a village in Scotland near which the mineral was discovered in 1790 by Adair Crawford and William Cruickshank; it was identified as a new element the next year from its crimson-red flame test color. Strontium was first isolated as a metal in 1808 by Humphry Davy using the then-newly discovered process of electrolysis. During the 19th century, strontium was mostly used in the production of sugar from sugar beet (see strontian process). At the peak of production of television cathode ray tubes, as much as 75 percent of strontium consumption in the United States was used for the faceplate glass. With the replacement of cathode ray tubes with other display methods, consumption of strontium has dramatically declined.

Transmetalation

Transmetalation (alt. spelling: transmetallation) is a type of organometallic reaction that involves the transfer of ligands from one metal to another. It has the general form:

M1–R + M2–R′ → M1–R′ + M2–Rwhere R and R′ can be, but are not limited to, an alkyl, aryl, alkynyl, allyl, halogen, or pseudo-halogen group. The reaction is usually an irreversible process due to thermodynamic and kinetic reasons. Thermodynamics will favor the reaction based on the electronegativities of the metals and kinetics will favor the reaction if there are empty orbitals on both metals. There are different types of transmetalation including redox-transmetalation and redox-transmetalation/ligand exchange. During transmetalation the metal-carbon bond is activated, leading to the formation of new metal-carbon bonds. Transmetalation is commonly used in catalysis, synthesis of main group complexes, and synthesis of transition metal complexes.

Unbinilium

Unbinilium, also known as eka-radium or simply element 120, is the hypothetical chemical element in the periodic table with symbol Ubn and atomic number 120. Unbinilium and Ubn are the temporary systematic IUPAC name and symbol, until a permanent name is decided upon. In the periodic table of the elements, it is expected to be an s-block element, an alkaline earth metal, and the second element in the eighth period. It has attracted attention because of some predictions that it may be in the island of stability, although newer calculations expect the island to actually occur at a slightly lower atomic number, closer to copernicium and flerovium.

Unbinilium has not yet been synthesized, despite multiple attempts from German and Russian teams. One 2011 attempt from the German team at the GSI Helmholtz Centre for Heavy Ion Research had a suggestive but not conclusive result suggesting the possible production of 299Ubn, but the data was incomplete and did not match theoretical expectations. Planned attempts from Russian, Japanese, and French teams are scheduled for 2017–2020. Experimental evidence from these attempts shows that the period 8 elements would likely be far more difficult to synthesise than the previous known elements, and that unbinilium may even be the last element that can be synthesized with current technology.

Unbinilium's position as the seventh alkaline earth metal suggests that it would have similar properties to its lighter congeners, beryllium, magnesium, calcium, strontium, barium, and radium; however, relativistic effects may cause some of its properties to differ from those expected from a straight application of periodic trends. For example, unbinilium is expected to be less reactive than barium and radium and be closer in behavior to strontium, and while it should show the characteristic +2 oxidation state of the alkaline earth metals, it is also predicted to show the +4 oxidation state, which is unknown in any other alkaline earth metals.

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